chemistry 120 handouts/notes

Chemistry 120 Chapter Two Notes

Chapter Two: Measurements Types of Measurements

Suppose I ask the question: “How far is it from here to San Diego?”

Your answer must contain two pieces information A numerical value A unit Examples About 70 miles About 1.5 hours Length Distance from one location to another Distance between the “center” of one atom and one it is attached to Mass The quantity of matter in whatever it is we are measuring The more matter an object contains, the greater its mass

Mass vs. Weight

Weight measures the gravitational attraction between an object and the Earth (or moon, Jupiter, etc.)

Weight is not the same as mass! Example: Consider a 1-ton (2,000 pound) rock on the surface of the Earth. How will its weight change if we move it to the moon? To Jupiter? How will its mass change?

Certain and Uncertain Digits

In every measurement we take, digits we are sure of are said to be certain. A digit which must be estimated is said to be uncertain.In taking a measurement, first determine which digits you know without question. Then estimate exactly one uncertain digit.

When & Where To Estimate

Suppose you are taking a measurement with a ruler that has tenths (0.1) of a centimeter as its smallest markings.

– You will estimate between the tenths mark, giving you an uncertain digit in the hundredths (0.01) place.

Suppose you are using a thermometer which has whole degrees as its smallest markings. – You will estimate between each degree to the tenth of a degree. – Be careful not to drop the estimated digit!

If the thermometer is exactly at the 15 degree mark, report 15.0 degrees.

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Certain and Uncertain Digits

In summary, in a given measurement, digits which are unambiguous are considered as certain.

In any measurement there is always exactly one uncertain digit which results from estimation. It is always the right-most digit.

In general, when taking a measurement on a device which measures to a given decimal place, estimate one digit to the right of this decimal place.

– This would not apply for most electronic devices (like digital balances), which already estimate for you.

Significant Figures Clearly, when we take measurements, there is a limit as to how accurate our data can be. For example, compare A bathroom scale A truck scale Significant figures are those values in a measurement which we can rely on for accuracy.

The Rules for Determining Which Digits are Significant

Non-zero digits – All non-zero digits (1-9) are always significant. – Example: How many significant digits are there in each value below?

48.9 7231.228 Leading zeros are zeros to the left of the first non-zero digit.

– Leading digits are not significant. – So, the zeros in 0.000825 and 0.0134 are not significant. – Why?

Enclosed zeros are zeros between other non-zero digits. – Enclosed zeros are always significant. – Examples of enclosed zeros:

5,002 901.0802 101 Trailing zeros are those zeros which come at the end (to the right) of a value.

– Trailing zeros are significant if there is a decimal point in the value. If there is no decimal, the zeros are ambiguous and generally not considered significant.

– Why?

Practice With “Sig Digs” How many significant digits do each of these values have?

462.0 1230 1230. 0.030900 0.0005

Calculations: Multiplication & Division When performing multiplication and division, count the number of significant digits in each

part of the calculation. Determine which value has the least number of significant digits. Your answer will be rounded to that many significant digits.

Examples

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Calculations: Addition and Subtraction

When adding or subtracting, consider the value with the least precision. Your answer must extend to that level of precision.

– For example, suppose you are adding 60.82 and 7.831 60.82 goes to the hundredths place 7.831 extends to the thousandths place Therefore, your answer will round to the hundredths place.

VERY IMPORTANT: Note that in adding and subtracting, the number of significant figures in the values does not matter!

Example More Examples

23.649 – 17.2 =

8.75 + 9.414 + 105.32 =

(6.834 × 8.32) – 3.45 =

Scientific Notation Many measurements in chemistry involve numbers that are extremely large or small. Examples:

Number of atoms in a glass of water. Distance between two atoms in a sugar molecule.

Scientific notation is a convenient mathematical notation which simplifies working with numbers of this magnitude.

How it Works

Take the value under consideration, and move the decimal point after the first non-zero digit.

Example: 582 5.82 Next, consider how many places the decimal was moved, and in what direction it was

moved. The decimal moved 2 places to the left. Our new value is 102, or 100 times, smaller than it was to begin with.

Undo this by multiplying by 10 to the power of how many places the decimal moved. This power is positive if the decimal moved to the left, negative if it moved to the right Our value is 5.82 × 102

Do not omit any significant digits! Examples

Express these numbers in scientific notation. 23,894. 0.004289

20.000

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Calculations Involving Significant Digits

In calculations involving the multiplication and division of numbers expressed in scientific notation, it is helpful to remember the algebraic expressions:

Examples Units of Measurement

Units are used in measurements to tell us – What type of measurement we are making

Distance Time Energy

– The general magnitude of the measurement Use feet to measure a person’s height Use miles to measure distance between distant cities Use light years to measure distance between distant galaxies.

The English System The English System of units includes many familiar units

Inches, Feet, and Miles Liquid Ounces Pounds and Tons Despite its popularity, the English System has a serious flaw

SI Units SI units (also called metric units) are based on powers of 10

– There are 10 millimeters in a centimeter – There are 10 centimeters in a decimeter – There are 10 decimeters in a meter.

You only need to know what each prefix means to be able to easily convert from one unit to another.

These prefixes are used in all types of measurements, including distance, volume, time, and energy among many others.

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Converting Into Different Units

Consider the following well known equality: 1 foot = 12 inches

Suppose we divide both sides by 12 inches:

Consider the following: “Convert 18.5 feet to inches.”

We need to get rid of feet, and end with inches. – Strategy: Divide by feet (in denominator), multiply by inches (in numerator)

Consider this example, using SI units: “How many centimeters are in 9.86 m?”

Another example using the SI system. “How many nanoseconds (ns) are there in 2.83 milliseconds (ms)?”

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Multiple Unit Conversions

In many measurements, such as in this problem, there are units in both the numerator and the denominator: “A car is traveling at 60. miles per hour. How fast is this in centimeters per second?”

A Side Note… By now you have noticed that units can be factored (or divided out) just like numbers You should also note that

– Units also can be multiplied Ex. 2 ft × 2 ft = 4 (ft × ft) = 4 ft2

Ex. 3 ft. × 2 lbs. = 6 ft.·lbs. (read “foot pounds”) You can only add and subtract measurements with the same units!

– Before adding and subtracting, perform any unit conversions to make the measurements have the same units (of course, they must be the same type of measurement!).

Area & Volume

Translating length to area involves squaring the units as well as the values which accompany them.

For example, what is the area of a square which is 5.0 cm on each side? 5.0 cm × 5.0 cm = 25 cm2

Let’s convert that to square meters:

Common units of volume you should know Liter(L) – the SI unit of volume milliliter(mL) – another common unit Cubic centimeter (cm3 or cc) – a unit equivalent to the milliliter

1 mL = 1 cm3 = 1 cc Dealing with volumes involves a similar procedure to that of areas.

A Volume Problem

A rectangular block is 14.6 in. by 7.2 in. by 6.8 in. What is its volume in cm3?

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Density Density (d) is the ratio between mass (m) and volume(V):

volumemassdensity =

Units of density for liquids are usually g/mL. For solids, we use the equivalent g/cm3. If something has high density, then a small volume of it will have a large mass.

– Ex. Which has greater mass, a liter of rocks or a liter of feathers?

Density Problems

The mass of an empty graduated cylinder is found to be 22.57 g. 7.25 mL of a liquid is added to it. The graduated cylinder and liquid have a combined mass of 30.79 g. What is the density of the liquid?

A 73.43 g cube of gold is dropped into a graduated cylinder whose volume reads 27.8 mL. After the cube sinks to the bottom, what volume reading will the graduated cylinder have? Note that gold has density 19.3 g/cm3.

Temperature Temperature is a measurement of the average kinetic energy of a body.

Kinetic Energy is the energy of motion, so the molecules in hot water are, on average, moving more rapidly than those in colder water.

Note, temperature is not the same as heat!

Units of Temperature Fahrenheit (°F)

– Commonly used to measure temperatures in the U.S., but not very practical for scientific use.

Celsius (°C) – Unit most commonly used in determining temperature in the laboratory.

Kelvin (K)

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– The SI unit of temperature. – 0 K is called “absolute zero.” It is impossible to for a system to have a temperature of

0 K (or lower).

The Temperature of the Freezing and Boiling Points of Water

TemperatureScale Freezing Point Boiling Point

°F 32 212

°C 0 100

K 273.15 373.15

Converting Temperatures

To convert between Celsius and Kelvin, use the following relationship: T(in K) = T(in °C) + 273.15

To convert between °C and °F, use the equation: °F = (1.8 × °C) + 32

A Temperature Problem

The temperature in Paris is 23 °C. What is the equivalent temperature in Kelvin? In Fahrenheit?

Accuracy & Precision in Measurements If data is accurate, this means that the results obtained are close to the “true and correct”

value(s). If a group of measurements give data which are close in value, these measurements are

said to be precise.

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Chemistry 120 Chapter Two Notes Chemistry 120 Chapter Two Notes

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Chemistry 120 Chapter Three Notes

Basic Concepts About Matter

Chemistry is the study of the properties and changes of matter What exactly is matter?

Matter is anything which has mass and takes up space (volume) Examples of matter:

Sand (a solid) Water (a liquid) Air (a mixture of gases)

How do we learn chemistry?

Chemistry is an empirical science, meaning that it is based on the results of experiments. In the lecture we will study theories and laws based on many years of observations and experiments. In the laboratory we will verify many of these principles.

A Conceptual Approach

During this course, we will focus on the fundamental concepts which define chemistry. Specifically, we will often look at matter at the smallest level (the submicroscopic level) to try to

reason why matter behaves the way it does on a larger, or macroscopic, scale. This is an essential skill for the aspiring chemist; a good imagination is all you need to develop it!

Atoms: A Brief Overview

Atoms are the most fundamental units of matter we consider in this course All matter which we encounter in our daily lives contains an extremely large number of these tiny

particles There are many different types of atoms

Some common examples are hydrogen, oxygen, gold, and sodium We will look closely at the structures of atoms at a later point in this course

Molecules

Two or more atoms may join together to form a molecule Molecules are held together by bonds

Specifically, these are called covalent bonds; more on these later! A diatomic molecule is made up of exactly two atoms (which may be the same or different)

Molecules Physical States of Matter

There are three common physical states of matter which we will consider in this class Gases Liquids Solids

We compare the three states by asking two questions Does the substance have a definite shape, or does it take the shape of its container? Does the substance have a definite volume?

Solids

Solids have a definite shape They do not assume the shape of their container

Solids have a definite volume The particles making up a solid

are close together do not move about, but vibrate in place tend to form organized patterns

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Liquids

Liquids have an indefinite shape They take the shape of their container

Liquids have a definite volume The particles making up a liquid

are fairly close together, but not to the same extent as solids move freely throughout the liquid are not organized in any particular pattern

Gases

Gases have an indefinite shape Like liquids, they take the shape of their container

Gases have indefinite volume The particles making up a gas

are generally far apart from one another move freely throughout their container in a random fashion

Classification of Matter

We classify matter as either a pure substance or a mixture There are two types of pure substances

Elements Compounds

There are two types of mixtures Homogeneous mixtures Heterogeneous mixtures

Elements

A pure sample of an element contains only one type of atom A sample of gold—an element—contains only gold atoms Helium contains only helium atoms

There are over a hundred known elements Each element is assigned a name and a symbol

Each symbol consists of one to three letters The first letter of the symbol is always capitalized; any other letters are always written in lower-

case The symbols (and occasionally the names) are catalogued on the Periodic Table of the Elements

Some Atomic Symbols H: Hydrogen O: Oxygen C: Carbon N: Nitrogen Na: Sodium Cu: Copper Cl: Chlorine K: Potassium

Note that some of these symbols are unusual! Elements

Some elements are generally only found as diatomic molecules We will call these seven elements the diatomic elements

H2 N2 O2 F2 Cl2 Br2 I2 Be sure you know these!

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Compounds Atoms come together in whole number ratios to form compounds. The chemical formula lists the ratio of these elements in the compound.

Each unit of water contains two hydrogen atoms and one oxygen atom H2O Other examples: NaCl, SiO2, C6H12O6 This formula/ratio is always the same for a chemical compound Different compounds may have the same formula. The compound is said to have a definite composition

Mixtures

The composition of a mixture is not fixed. Consider salt water, a mixture of H2O and NaCl.

Is salt water always found in the same proportion? The same atom-to-atom ratio?

Mixtures can be classified into two types: Homogeneous mixtures have all parts in the same state (gas, liquid or solid) and all parts must be

mixed together. If the parts of the mixture are visually inseparable, we will call the mixture homogeneous.

Heterogeneous mixtures are simply those which are not homogeneous. Separations

Mixtures can be separated by physical methods A heterogeneous mixture of coffee grounds and water can be separated by filter paper The water can be removed from a salt water solution (a homogeneous mixture) by boiling the

water off. The salt will remain in the container. Compounds can only be separated into their individual elements by chemical means (i.e. through the

result of a chemical reaction)

Mixtures Ex. Are each of these mixtures homogeneous or heterogeneous? Why?

Vodka (a mixture of water and ethyl alcohol)

Cheerios in milk

A mixture of oil and water

A salt water solution Note that the term solution is often used to refer to homogeneous mixtures, especially for

compounds dissolved in water.

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Air

Dirty, dust-filled air

Classification Problems Classify each of these as an element, a compound, a homogeneous mixture, or a heterogeneous

mixture. Tap water

Steel (an alloy of several metals) Note that alloys can be mixed in different proportions.

Helium

Mud

Carbon dioxide

Properties of Matter We can describe matter in two ways:

By its chemical properties, which describe how a type of matter interacts (or “reacts”) with another type of matter.

For example, hydrogen is able to react with oxygen to form water. Helium reacts with virtually nothing.

By its physical properties, which include all non-chemical properties.

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For example, water is a liquid at room temperature, freezes at 0 °C, has a density of 1.0 g/mL, and is both clear (we can see through it) and colorless.

Changes of Matter

Common changes of matter are described in essentially the same way as properties are A chemical change is a change which involves a chemical reaction

Bonds are formed and/or broken in a chemical change The chemical substances you end with are fundamentally different than what you began with

A physical change is a change which does not involve a chemical reaction All changes of state of a given substance are physical changes

Examples include ice melting and liquid water boiling

Classify each change as either a physical change or a chemical change Gasoline evaporates off of the ground

A glass vase is shattered

Sodium reacts with chlorine, forming sodium chloride

Dry grass burns in a large fire

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Chemistry 120 Elements and Symbols You Must Know

Group1A

8A

HHydrogen

2A 3A 4A 5A 6A 7A HeHelium

LiLithium

BeBeryllium

BBoron

CCarbon

NNitrogen

OOxygen

FFluorine

NeNeon

NaSodium

MgMagnesium

3B 4B 5B 6B 7B 8B 1B 2B AlAluminum

SiSilicon

PPhosphorus

SSulfur

ClChlorine

ArArgon

KPotassium

CaCalcium

ScScandium

TiTitanium

VVanadium

CrChromium

MnManganese

FeIron

CoCobalt

NiNickel

CuCopper

ZnZinc

GaGallium

GeGermanium

AsArsenic

SeSelenium

BrBromine

KrKrypton

RbRubidium

SrStrontium

PdPalladium

AgSilver

CdCadmium

SnTin

SbAntimony

IIodine

XeXenon

CsCesium

BaBarium

WTungsten

PtPlatinum

AuGold

HgMercury

PbLead

BiBismuth

UUranium

You must memorize the symbols and names (with the correct spelling) for each elementlisted on this periodic table. I recommend making flashcards to help you in this.

Chemistry 120 Standard States of the Elements

Group1A 8A

H2Hydrogen


LiLithium

BeBeryllium

Key: solid liquid gas BBoron

CCarbon

N2Nitrogen

O2Oxygen

F2Fluorine

NeNeon

NaSodium

MgMagnesium

3B 4B 5B 6B 7B 8B 1B 2B AlAluminum

SiSilicon

PPhosphorus

SSulfur

Cl2Chlorine

ArArgon

KPotassium

CaCalcium

ScScandium

TiTitanium

VVanadium

CrChromium

MnManganese

FeIron

CoCobalt

NiNickel

CuCopper

ZnZinc

GaGallium

GeGermanium

AsArsenic

SeSelenium

Br2Bromine

KrKrypton

RbRubidium

SrStrontium

PdPalladium

AgSilver

CdCadmium

SnTin

SbAntimony

I2Iodine

XeXenon

CsCesium

BaBarium

WTungsten

PtPlatinum

AuGold

HgMercury

PbLead

BiBismuth

UUranium

This periodic table shows you the state (gas, liquid, or solid) of many important elements at room temperature. Notice that all metals and metalloids except mercury are solids at room temperature. The only two elements which are found as liquids are bromine and mercury. All the elements in the last group (8A) are gases. The rest are fairly easy to memorize. On this periodic table I have indicated the diatomic elements with their formulas; you will not see this on a "normal" periodic table. I have left the boxes of the elements which are not important in this course blank.

Chemistry 120 Chapter Four Notes

Chapter Four: Matter and Energy Matter Recall that matter Has mass Takes up space (i.e. has volume) Matter also has one of four states Gas Liquid Solid Plasma (We will not be concerned with the plasma state in this course!) States of Matter The states of matter are classified by two parameters Does it take the shape of its container? Does it completely fill its container? Gases Take the shape of their container Completely fill their container Liquids Take the shape of their container Do not fill their container completely Solids Do not take the shape of their container Do not fill their container completely Which state a substance is in depends on two factors: Temperature and Pressure Gases are comprised of molecules which generally are far apart from one another, and travel in random paths. The particles in liquids and solids are much closer together than those of gases The particles of solids are generally “stuck in place”, but are able to vibrate Liquid particles freely move across one another Changes of State Special terms are associated when matter changes from one state to another. You may be familiar with the term “freezing,” which describes a change from liquid to solid. Solid Liquid Gas

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Properties of Matter We can describe matter in two ways: By its chemical properties, which describe how a type of matter interacts (or “reacts”) with another type of matter.

For example, hydrogen is able to react with oxygen to form water. Helium reacts with virtually nothing.

By its physical properties, which include all non-chemical properties. For example, water is a liquid at room temperature, freezes at 0 °C, has a density of 1.0 g/mL, and is both clear (we can see through it) and colorless.

Changes of Matter Similarly, changes can be classified in the same way as properties: Chemical changes involve a rearrangement of atoms, producing chemical compounds that were not there before.

When iron (Fe) is exposed to oxygen on a wet day, rust (Fe2O3) is formed. Physical changes are those which do not involve a chemical change.

Boiling water (liquid H2O changes to gaseous H2O), glass shattering, a chemical evaporating.

Metals The metals include many elements found on the left side of the periodic table. Many metals are known for having the following properties

They are malleable (meaning they are soft and easily shaped) They are ductile (they can be twisted and drawn into a wire) They can conduct both electricity and heat. They tend to be lustrous (shiny). All metals are solids at room temperatures except for mercury, which is a liquid.

Nonmetals Nonmetals are generally found on the right side of the periodic table (except hydrogen, which is placed on the left). Their properties are generally the opposite of the metals.

Those which are solids tend to be brittle. Most are poor conductors of electricity at room temperature (insulators) and do not conduct heat well. Some are gases at room temperature; others are solids. Bromine is the only other element which is a liquid.

Metalloids Metalloids are found between the metals and the nonmetals on the periodic table. They include B, Si, Ge, As, Sb, and Te. The properties of the metalloids are often a cross between those of the metals and the nonmetals. All the metalloids listed are solids at room temperature.

(I do not include Po and At with the metalloids. They are rather unstable and unimportant to our studies at this point.)

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The Law of Conservation of Mass One of the most fundamental statements about matter is described by this law, which states: “Matter can neither be created nor destroyed in any chemical process” Another consequence of this law is that one type of atom cannot be changed into another through a chemical reaction.

If you start a chemical process with 10 million hydrogen atoms and 10 million oxygen atoms, then you will end the process with 10 million hydrogen atoms and 10 million oxygen atoms. ALWAYS!

Energy Energy is an important subject in chemistry, as chemical reactions either give off or take in energy. The Law of Conservation of Energy states that “Energy can neither be created nor destroyed in a process.” Energy can be transferred between two systems in two ways

As work (w), which we will consider as energy put to some specific use, like making a motor run. As heat (q), which will be random energy(not directed to some useful purpose), like that given off by your car’s engine.

Any change in energy is just the sum of the work and heat changes. The SI unit of energy is the Joule (J). Other units include the kJ, and the calorie (cal)

1 calorie = 4.184 Joules For example, suppose that burning a certain amount of gasoline produces 50 kJ of energy. If only 20 kJ of it goes into doing work in a car engine, the remaining 30 kJ must be lost as heat. Remember, energy cannot be destroyed! An exothermic process is one which gives off more heat than it takes in. For an exothermic reaction, q < 0. For example, consider burning gasoline. An endothermic process is one which brings in more heat than it gives off. For an endothermic reaction, q > 0. Specific Heat Some substances require more energy to raise their temperatures than others.

For example, it requires much less energy to raise the temperature of 50. g of aluminum by 10 °C than it would to raise 50. g of water by the same amount.

This difference is represented by a constant called the specific heat, c. Its units are J/(g ·°C) or cal/(g ·°C) . Transferring Heat The amount of heat (q) absorbed or given off by a substance when it changes temperature can be found using the following equation: q = m × ΔT × c m is the mass of the substance ΔT is the change in the temperature; it equals (final T – starting T) c is the specific heat of the substance Note that heat always flows from an area of high temperature to one of lower temperature!

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Example How much heat is required to raised the temperature of 50.0 g of aluminum from 32 °C to 47 °C? The specific heat of aluminum is 0.215 cal/(g ·°C). A More Difficult Example A 50.0 g aluminum block is heated to 75 °C, then dropped into a sealed container containing 350. g of water at 15 °C. What will be the temperature of the block when it is finished cooling? More On Conservation Note that in the last problem, heat was transferred from the aluminum to the water, but never destroyed. Einstein is credited with the famous equation: E = mc2

Where E is energy, m is mass, and c is a constant (the speed of light, not the specific heat!) This shows that matter can be converted to energy. More on Conservation Note however, this will not apply to chemical reactions. We see this with nuclear reactions. In this class, mass is always conserved.

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We can now state the Law of Conservation of Mass-Energy: “Mass and energy can neither be created nor destroyed in a process, though they may be interchanged.” Kinetic and Potential Energy Kinetic Energy (K.E.) is the energy of motion. The amount of K.E. an object has is described by the formula K.E. = ½mv2

where m is the object’s mass, and v is its velocity. Potential Energy is the energy of position. Objects may be in a “high energy” position, and can convert this potential energy to kinetic energy while moving to a “lower energy” position. Examples:

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Chemistry 120 Chapter Five Notes

Chapter Five: Atomic Theory and Structure Evolution of Atomic Theory

The ancient Greek scientist Democritus is often credited with developing the idea of the atom Democritus proposed that matter was, on the smallest scale, composed of particles described

as atomos, meaning “indivisible” While atoms can in fact be broken down into smaller particles, the atoms of each element are

distinct from each other, making them the fundamental unit of matter.

Dalton’s Atomic Theory Compiling experimental information available during his lifetime, John Dalton described an

accurate picture of atoms in 1803 with his atomic theory. Dalton’s Atomic Theory maintains that

All elements are made up of tiny, indivisible particles called atoms, which can be neither created nor destroyed in reactions.

Atoms of the same type of element are the same; those of different elements are different.

Atoms of different elements form compounds by combining in fixed, whole number ratios. This is also called “The Law of Definite Proportions”

If the same elements combine to make more than one compound, there can be a different, but definite, atom ratio for each compound.

This is also called “The Law of Multiple Proportions” In actuality, atoms combining in the same ratio can make different compounds.

A chemical reaction does not involve a fundamental change of the atoms themselves, but of the way they are combined together.

Probing Further into Atomic Structure: The Electron

One of the three particles which make up all atoms is the electron Benjamin Franklin observed that, when a cloth was rubbed across a glass rod, a charge was

developed on each. The charges, called positive (+) and negative (-), show an attractive force between opposites

and repulsion between identical charges.

The Atom: A Complete Picture Each atom contains at its core a nucleus, a region of positive charge. Positively charged particles, called protons, are contained in the nucleus. Electrons (negatively charged particles) “orbit” around the nucleus throughout the atom. Later experiments also confirmed that all atoms except hydrogen must contain one or more

neutral (non-charged) particles called neutrons. Note that the protons and neutrons are each almost 2,000 times more massive than an

electron; therefore, most of the mass of an atom is in the nucleus. The protons and neutrons are called the nucleons.

Characteristics of an Atom

The number of protons in an atom is called the elements atomic number, which is symbolized by Z.

The number Z indicates which type of element that atom represents. These values are found in order on the periodic table. Example: Which element contains 5 protons? 10? 34?

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Recall that the “Law of Conservation of Mass” indicated that atoms cannot be changed in different elements by chemical means.

So, protons are never “changed around” in a chemical reaction!

Characteristics of an Atom Atoms that have no charge must have the same amount of positive charge as negative

charge. Therefore, the number of electrons in an atom is equal to the number of protons in an atom

(Z) for any neutral atom. Atoms which contain more or less electrons than protons therefore must have a charge.

These are called ions.

Ions: A Lesson in Thinking Backwards Suppose an ion has exactly one more electron than it does protons.

Will the ion be positively or negatively charged?

What if an atom lost two electrons? Will the ion be positively or negatively charged?

For each electron an ion has more than it does protons, we indicate it with a – as a superscript.

For each electron an ion has less than it does protons, we indicate it with a + as a superscript.

Examples of Ions

A bromine atom normally has_______protons and________electrons.

Suppose we add exactly one electron; now we have________protons and_______electrons. We symbolize this as Br- (or Br1-).

An aluminum atom normally has________protons and_________electrons. Suppose the atom loses three electrons. We symbolize this ion as Al3+. Note that losing electrons is indicated with +, and gaining electrons is indicated with -.

Neutrons and Isotopes

The number of neutrons in an atom cannot be easily predicted or found on the periodic table. Atoms of the same element (i.e. have the same number of protons) can have different

numbers of neutrons. They are called isotopes of one another. Examples:

A carbon atom must contain 6 protons, but it may contain 6, 7, or 8 neutrons. Almost all hydrogen atoms contain zero neutrons. The isotopes of hydrogen which

contain one and two neutrons are called deuterium and tritium, respectively; they are symbolized as D and T, but do not appear on the periodic table.

More on Isotopes The sum of the number of protons and neutrons in an atom is called the mass number of that

atom, and is symbolized by A. Isotopes are named by stating the name of the element, followed by A.

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Carbon (Z=6) with 6 neutrons is “carbon-12” Carbon with 7 neutrons is “carbon-13” Carbon with 8 neutrons is “carbon-14”

More on Isotopes

Isotopes can be written in shorthand notation in two different ways,

X or X AAZ

Examples For each of the following, indicate how many protons, electrons, and neutrons each atom or

ion possesses. 35Cl

81Br-

27Al3+

The Mass of an Atom Recall that virtually all of the mass of an atom comes from its nucleus. Knowing the mass of protons and neutrons allows us to calculate the mass of one atom of a

particular isotope. Since most elements have more than one isotope, and these isotopes are mixed in nature, it

is not possible to provide an exact mass for the element. Instead, we consider a weighted average mass, based on the weight of each individual

isotope and its abundance in nature. The periodic table provides the atomic weight of each element, which corresponds to this

weighted average mass. These values are in atomic mass units (amu), a very small unit of mass.

1 amu = 1.66 × 10-24 g For example, the mass of a single helium atom is 4.003 amu.

Calculating Atomic Weights

The atomic weight of an element can be calculated if we know: The atomic weight of each isotope of that element, and The percent abundance of that element in nature.

For example, suppose that 2 isotopes exist for an element X, and that the isotopes have atomic weights of 64. amu and 66. amu. If there are 50.% of each in nature, what is the atomic weight of X?

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Examples There are three isotopes of potassium, as the data in the table below indicates. Calculate the

atomic weight of potassium from this data.

Isotope Mass(amu) % Abundance 39K 38.963707 93.2581 40K 39.963999 0.001171 41K 40.961826 6.7302

Chlorine has only two naturally occurring isotopes. The atomic weight of 35Cl is 34.96885 amu, and its % abundance is 75.53%. What is the atomic weight of the other isotope, 37Cl? Percent Composition

The percent composition of a compound indicates the percentage of each element in the compound by mass

For example, suppose that a compound containing only nitrogen and oxygen is 36.85% nitrogen

From this, we see that, for every 100. grams of the compound, 36.85 grams of it comes from the nitrogen in the compound

What about oxygen?

5-4


Suppose a compound is analyzed and is found to contain 6.00 grams of carbon, 2.00 grams of hydrogen, and 8.00 grams of oxygen. What is the percent composition of the compound? What mass of hydrogen is in a 1.450 kg sample of this compound?

5-5

Chemistry 120 Chapter Ten Notes

Chapter Ten: Modern Atomic Theory

Light and Energy

Electromagnetic radiation(EM) is an especially important form of energy for scientific study. Many types of “radiant” energy are included under this description, including visible light,

X-rays, and radio waves. EM can be described as waves, which can be characterized by their wavelengths (λ), the

distance between the peaks of each wave. We also may describe light as a particle, called a photon

Wavelength and Frequency The frequency (ν) of waves is measured as the number of waves which occur in a period of

time, usually a second. For example, there are 445 peaks in a second, we would indicate this as 445 s-1. Alternatively, the Hertz(Hz) may be used

1 Hz = 1 s-1 The frequency and the wavelength are inversely proportional

That is, as the wavelength increases, the frequency decreases, and vice versa.

The Mathematical Relationship Between Wavelength and Frequency

Where c is the speed of light c = 3.00 ×108 m/s λ

=νc

Energy of EM Radiation

The total energy of a single photon depends on the frequency (or wavelength) which that photon is associated with.

This is described by the equation E = hν where E is the energy (in J), ν is the frequency (in s-1), and h is Planck’s Constant (6.626 ×10-34 J·s).

Greater frequency (and shorter wavelengths) indicate higher energy photons. Provide a formula relating energy to wavelength rather than frequency:

The Electromagnetic Spectrum EM radiation can be classified into different regions based on its wavelength or its frequency.

The longest wavelenths (smallest frequencies) correspond to radio waves. The shortest wavelengths (greatest frequencies) correspond to gamma rays.

10-1


Electrons & Energy Levels

Evidence suggested to scientists studying atomic structure that electrons “orbit” around the nucleus in specific energy levels (also called shells).

The farther the energy level is from the nucleus, the higher the energy of its electrons. Normally, electrons “fill up” the lower energy levels first, and as new electrons are added they

go into higher and higher energy levels. An atom is said to be in the ground state when this is true.

An energy source, such as light or heat, can give electrons the energy necessary to “jump” from its energy level to a higher one.

In this situation, the atom is said to be in an excited state.

Emission of Light An atom in the excited state is unstable and must release the energy it gained in the first

place to return to the ground state. According to the Law of Conservation of Energy, this extra energy cannot simply

disappear One way in which this is accomplished is for the atom to give off this energy as light. This

process is called emission. The released light is called a quanta (energy packet) or a photon.

Excited electrons return to lower energy levels. Depending on the atom itself and which energy levels are involved, EM radiation of

different wavelengths are given off. Example Suppose that a photon has 2.45 × 10-19

J of energy. (a) What is its wavelength in meters?

(b) What is its frequency in s-1? (c) What type of EM radiation is this?

So What is Light, Really?

The answer to this is not as simple as you might think Light (EM in general) is correctly described as both

Particles (photons) Waves

Light is not one or the other; it is both simultaneously!

10-2


Quantization Notice that, for a given atom, electrons are only allowed to jump to discrete energy levels.

In other words, an electron in the third level may fall to the first or second level, but not “in between” the levels.

As a result, only specific energy changes are allowed and can be observed for each atom. We say that these energy changes are “quantized” since only specific quantities of

energy can be emitted, As an analogy, compare a ball moving down a set of stairs (step-by-step) to one moving

down a ramp. Which corresponds to “quantized” energy values?

Orbitals & Subshells In truth, electrons do not simply rotate about the nucleus in simple patterns. Instead, electrons are mostly contained in orbitals, which are shapes which describe the

regions where an electron is likely to be found. We will call a complete “set” of orbitals in a given energy level a subshell. Four orbital types are commonly encountered in modern chemistry.

These are designated s, p, d, and f, and are listed here from lowest energy orbital type to highest.

Each individual orbital can hold, at most, two electrons in its ground state. The s Orbital

The s orbital is the least energetic of the orbitals, and, as a sphere, is also the most simple-looking.

All energy levels have exactly one s orbital. Since any orbital can only hold two electrons, each s orbital is limited to this.

The p Orbital

Starting with the 2nd energy level, each level possesses three “sets” of p orbitals. There are three 2p orbitals, three 3p orbitals, etc.

Recall that each orbital can hold up to two electrons, so the three orbitals combined will hold up to six electrons.

The p orbitals are more complex in shape than the s orbital, and electrons residing in them typically have greater energy than those in an s orbital.

10-3


d and f Orbitals

Energy levels three and higher have a total of 5 sets of d orbitals.

A full set of d orbitals can therefore hold up to_____electrons. Energy levels four and higher have a total of 7 sets of f orbitals.

A full set of f orbitals can therefore hold up to_____electrons.

Electron Configurations

Each element has its own unique electron configuration which describes which orbitals have electrons in them and how many.

For example, the electron configuration of boron is 1s22s22p1, which means that there are 2 electrons in an s orbital on the first energy level (called the 1s subshell) 2 electrons in the 2s orbital (the 2s subshell) 1 electron in a 2p orbital (the three p orbitals collectively make up the 2p subshell)

Atoms always fill their lowest energy orbitals first, and successively higher ones as more

electrons are added. This is known as the Aufbau principle; aufbau is German for “building up”

The order in which the electrons fill can be found on the periodic table. Notice that the periodic table is broken up into 4 distinct “blocks,” each of which indicates

where the highest energy electrons are. Examples Determine the electron configuration for each of the following atoms/ions: Be

N

Na

V

10-4


Shorthand Notation for Electron Configurations You can abbreviate the electron configuration with a special notation

Consider V for example. The last noble gas before V was Ar (element 18) So, we can write [Ar]4s23d3 as vanadium’s electron configuration. What is the electron configuration of antimony in shorthand notation?

Unusual Electron Configurations Unfortunately, the electron configurations we predict are not always correct For example, predict the electron configuration of copper

The actual electron configuration is:

There are several other elements which do not “follow the rules”; however, at this point in your study of chemistry, you do not need to be concerned with these exceptions Spin

Each electron is able to rotate about its central axis, a process we call spin Electrons may spin in one of two possible directions We refer to these directions as “spin up” and “spin down”

The Pauli Exclusion Principle & Hund’s Rule

In a ground state atom, two electrons in the same orbital must have opposite spins; this is a simplified version of a rule called the Pauli Exclusion Principle

Another important principle, Hund’s Rule, states that, for ground state atoms, electrons will fill unoccupied orbitals within a subshell before filling singularly-occupied ones

For example, in a nitrogen atom, with electron configuration 1s22s22p3, the three electrons in the 2p subshell reside in different orbitals (i.e. they do not pair up)

Energy Level Diagrams

Energy level diagrams are a graphical form of the electron configuration, representing electrons as arrows (up arrow for spin-up, down for spin-down) and orbitals as lines

The lines are listed from lowest energy at the bottom (i.e. the 1s orbital) to highest

10-5


Example

Draw the energy level diagrams for a. titanium b. S2-

10-6

Chemistry 120 Orbital Blocks

Group1A 8A


3B 4B 5B 6B 7B 8B 1B 2B

Here you see the periodic table separated into blocks, which should assist you in your studyof electron configurations. Note that helium belongs in the s block.

s block

d block

p block

fblock

Chemistry 120 Chapter Eleven Notes

Chapter Eleven Chemical Bonding Valence Electrons

The electrons occupying the outermost energy level of an atom are called the valence electrons; all other electrons are called the core electrons.

The valence electrons, as we will see, are responsible for chemical bonding. Knowing the number of valence electrons an atom has is the single most important

information you can have in predicting how an atom will react chemically. Note that the valence electrons are not always the electrons with the highest energy, as we

will see with the transition metals For elements in group IA through VIIIA, the number of valence electrons an atom has is

simply its group number. So phosphorus, in group VA, has five valence electrons.

Helium is the only significant exception; it has 2 valence electrons, even though it is in group VIIIA

Lewis Dot Diagrams

Lewis dot diagrams are simple symbols which show the symbol of an element surrounded by as many dots as that atom has valence electrons.

The first four electrons are drawn (in any order) on each of the four “sides” of the symbol. The next four electrons are “paired up” with the other four atoms (again, in any order). These symbols will be used to model chemical bonding.

Examples

Draw the Lewis dot diagrams for selenium, rubidium, and manganese atoms.

The Octet Rule The Octet Rule states that atoms react in such a way as to give them eight electrons in their

outermost energy level (their valence shell). The Noble Gases, those elements in the period on the far right of the Periodic Table,

generally do not react with other atoms, as their valence shell is filled with eight electrons (or two in the case of helium).

By gaining or losing enough electrons to give an atom eight electrons in its valence shell, the atom gains the stable features of a noble gas.

1


Getting an Octet

Atoms usually achieve an octet in the following way: Metals lose their valence electrons, giving them an empty outer shell.

Note that the ion now has the electron configuration of a noble gas. Nonmetals gain or share enough electrons to give themselves eight electrons in their

outer shell. Nonmetals generally share with other nonmetal atoms, or take electrons away from

metal atoms.

Example How many electrons will usually be gained or lost by each of the following atoms?

Bromine

Barium

Potassium

Aluminum

Oxygen

Neon

The Duet Rule Hydrogen atoms attempt to pick up another electron, giving them the same electron

configuration as the noble gas helium; this is called The Duet Rule. Likewise, when lithium loses an electron to become Li+, it has the same electron configuration

as helium.

A Cautionary Note Gaining and losing electrons does not occur unless there are other atoms around to

encourage this. A sodium atom cannot just toss its valence electron away; it must give it to another atom

which wants to take it, like chlorine. Considering this, be sure that you do not confuse an atom of an element (which has its usual

number of electrons) with an ion (which has gained or lost electrons).

What is Bonding?

Bonding describes how atoms interact with each other in an attractive sense. There are three types of bonding:

Ionic bonding Covalent bonding Metallic bonding

2


Ionic Bonding

Ionic bonding occurs between ions of different charges. Often, but not always, ionic bonding occurs between the ion of metal with the ion of a

nonmetal. Sometimes, one of these ions is a polyatomic ion, such as sulfate (SO4

2-). Ionic bonding is generally the strongest of the bonding types.

It requires a great amount of energy to separate the ions from each other, resulting in high melting points.

Examples of Ionic Compounds NaCl CaO CuBr2 LiF Formulas of Ionic Compounds

An ionic compound should have no total charge in its formula. Ions of opposite charge are paired in such a way as to cancel out charges.

Positively charged ions are called cations. Negatively charged ions are called anions.

Example: One Na+ combines with one Cl- to give the neutral ionic compound NaCl. Two Na+ combine with one O2- to give the neutral ionic compound Na2O.

Notice that the cation is always listed first in the formula.

Formulas of Ionic Compounds

Further examples Note that the formula of an ionic compound shows the smallest whole number ratio

between the ions. So, when Ca2+ and O2- combine, the formula of the product is CaO, not Ca2O2. Sometimes, the ratio between ions is more complex, like with Al3+ and O2-.

Example What is the formula of the ionic compound which contains ions of each of the two elements

below? sodium and sulfur barium and nitrogen selenium and calcium

3


Covalent Bonding

Covalent bonds result from the sharing of electrons between two atoms. Recall that covalent bonding is usually found between nonmetal atoms.

This is a bit oversimplified, but will work well for our purposes in this class. In following the octet rule, atoms share enough electrons with each other to give each eight

valence electrons (or two in the case of hydrogen). Covalent bonds are generally not as strong as ionic bonds; still, many are relatively strong.

Example: Compare heating salt (ionic) and sugar (covalent). Molecules

Compounds which are held together by covalent bonds are called molecules This term is not appropriate for ionic compounds, which are made up of repeating unit cells For example an H2O molecule contains exactly three elements, but there is no “NaCl

molecule” consisting of only two atoms Examples of Covalent Compounds H2O C6H12O6 Cl2 (an element) C6H6CO2H2S SO3 Diatomic Elements

Seven elements are essentially always found in nature as diatomic molecules, which are molecules which contain exactly two atoms

The formulas of these seven elements are: H2, N2, O2, F2, Cl2, Br2, and I2

When we refer to oxygen gas for example, we mean O2, not individual oxygen atoms. These seven formulas will appear frequently throughout the course and must be learned!

Phosphorus, Sulfur, and Ozone

The most stable form of phosphorus is a molecule containing four P atoms, which is given the formula P4.

Similarly, sulfur forms molecules which contain eight sulfur atoms each, and has the formula S8.

Unlike for the diatomic elements, chemists do not consistently use these formulas when writing chemical reactions

Oxygen also forms the molecule known as ozone, O3. Although both O2 and O3 contain only oxygen atoms, they have entirely different chemical

properties

4


Electronegativity The electronegativity of an atom indicates how strong its attraction for electrons is.

The greater an atom’s electronegativity, the more strongly it pulls electrons toward it. Like the atomic radius and the ionization energy, electronegativity has a periodic trend:

Electronegativity increases up a group and from left to right across a period. Fluorine is the most electronegative atom; in general, the closer a main group element is

to fluorine on the periodic table, the more electronegative it is. The noble gases have virtually no electronegativity.

Lewis Structures

Lewis Structures are simple drawings that show the covalent bonds between atoms. As in Lewis dot diagrams, electrons are represented by dots.

Since two electrons are required to make a bond, a pair of electrons between two atoms indicates a bond.

Multiple Bonds Remember, atoms bond in order to gain an octet (or duet for hydrogen). In order to do this, atoms can also share more than one pair of electrons, up to a maximum of three pairs.

Sharing one pair of electrons forms a single bond. Sharing two pairs forms a double bond. Sharing three pairs forms a triple bond. There are not quadruple bonds!

Examples of Multiple Bonds More on Lewis Structures

It is much simpler when drawing Lewis structures to represent electron pairs as lines rather than dots.

However, always keep in mind that the lines stand for electron pairs!

5


Drawing Lewis Structures

There are several steps involved in taking a formula and converting it to a Lewis structure; just keep in mind the goals set out by the octet and duet rules.

1. Add up the total number of valence electrons in the formula.

Ex. NH3 has 5 + (3×1)=8 v.e. Note that if the molecule is an ion, each negative charge adds one electron to this count, and each positive charge takes one away.

2. Draw a simple “skeleton” of the molecule, showing which atoms are connected to which. If there is one “central” atom (the one surrounded by the others) it is usually the least electronegative. Note: Hydrogen is never a central atom!

3. Figure out how many electrons you have left. Remember that each bond stands for two electrons, so subtract two from the total number of valence electrons for every bond you drew in the previous step. 8 – 6 = 2 electrons left

4. Give enough electron pairs to each atom in the structure to give each atom (except hydrogen) an octet, using the following guidelines: -Electrons go to the most electronegative atoms first. -You must stop adding electron pairs when you run out of electrons (from the original number you started with in step 1).

5. This step is not always necessary: If you have run out of electrons, and some atoms still need octets, change a lone pair on the least electronegative adjacent atom (except hydrogen) to a double bond between the two atoms. If necessary, you may need to make a triple bond by taking away another lone pair. Note that fluorine will not share its lone pairs.

6


Examples Draw Lewis Structures for each of the following molecules:

CH4

H2O

HCN NO3

- N2H2 C2H2

7


Exceptions to the Octet Rule

Boron and beryllium often will not be able to gain a complete octet. For example, consider BH3 and BF3

Atoms in period 3 and below only may exceed the octet rule for reasons we will not consider

here. You will often see this in cases where the central atom is P, S, Cl, Br, or I; others exist as well.

For example, consider PCl5, SF6,and SO3.

8

Additional Chapter 11 Topics Periodic Trends

By comparing the position of one element to another on the periodic table, it is often possible to make comparisons between the properties of those elements; these are a result of periodic trends.

In this class we consider three of these trends: Electronegativity Atomic radius Ionization Energy

Atomic Radius

The radius of an atom tends to increase as you move down a group in the periodic table. Explanation: Elements in lower groups of the periodic table have electrons in outer

energy levels. These electrons are not held as tightly by the nucleus, so they can travel further away from it.

The radius of an atom tends to decrease from left to right across a period. Explanation: The valence electrons of atoms in the same period fill the same energy

levels. However, the positive charge in the nucleus is increasing as we move from left to right on the periodic table as the number of protons is likewise increasing. This pulls the electrons in closer, making the radius smaller.

Atomic Radius: Ions

Recall that when an ion is formed, an atom gains or loses electrons, while the number of protons remains the same

When an atom loses an electron to form a cation, the positive charge of the nucleus pulls the remaining electrons closer towards the nucleus

A monatomic cation is smaller than the corresponding neutral atom Na+ has a smaller atomic radius than Na Mg2+ has a smaller atomic radius that Mg

When an atom gains an electron to form an anion, the increased repulsion of the additional electrons results in electrons moving farther from the nucleus

Cl- has a larger atomic radius than Cl O2- has a larger atomic radius than O

When comparing atoms and ions with the same number of electrons, the largest ions are those with the most negative charge

Atomic Radius: Ions Rank each of the following atoms/ions in order from largest to smallest

I. Ne Al3+ O2- F- Na+ Mg2+

II. Li K N Ne Cs

Ionization Energy

The ionization energy is the amount of energy required to remove an electron from a mole of atoms in their gas state of a particular element. X(g) X+

(g) + e-

The first ionization energy is the amount of energy required to remove the first electron. The second ionization energy is the amount of energy required to remove the second

electron. Etc. Ionization energy tends to decrease down a group.

Explanation: Electrons in the outer shells are further from the nucleus and feel its pull less than the core electrons do. Less energy is required to remove them.

Ionization energy tends to increase from left to right across a period. Explanation: Protons are added to the nucleus from left to right across the table, and

these protons exert a greater pull on the electrons, requiring more energy to remove them.

Why is the first ionization energy of oxygen less than that of nitrogen?

Comparing Ionization Energies

Generally, the first ionization energy is less than the second, which is less than the third, etc. If an atom has lost enough electrons to give it an octet in its outer shell, it requires a

tremendous amount of energy to remove any electrons beyond this. Covalent Bonds

Recall that Covalent bonds result from the sharing of electrons between atoms. Nonmetals bond with other nonmetal atoms through covalent bonding. The strength of a covalent bond is generally less than that of an ionic bond.

A question: Do atoms involved in covalent bonding share electrons equally?

Polarity and Dipole Moment

If two atoms which are covalently bonded have substantially different electronegativity values, then the bond is said to be polar covalent.

The electrons are not shared equally in the bond. The electrons will tend to stay closer to the more electronegative atom.

The molecule might have a dipole moment, which is shown by drawing a special arrow pointing towards the more electronegative atom. More on this later…

Shapes of Molecules

From Lewis structures we can usually figure out the three-dimensional structure of simple molecules.

VSEPR theory (stands for Valence Shell Electron Pair Repulsion) helps us to understand the shapes that molecules take.

According to VSEPR, electron pairs move as far away from each other (in three dimensions) as possible.

Remember that electrons all have negative charge, so they repel one another.

The Basic Geometries of Molecules Depending on the number of attached atoms and lone pairs, a basic geometry can be

described for an atom in a molecule. Linear: 2 attached groups Trigonal Planar: 3 attached groups Tetrahedral: 4 attached groups Trigonal Bipyramidal: 5 attached groups Octahedral: 6 attached groups

Assigning Shapes to Molecules

A shape is assigned to each central atom in a molecule. To be a central atom, an atom must have two or more atoms bonded to it.

The number of atoms bonded to the central atom is added up. Call this m. Careful! Add the number of atoms, not the number of bonds!

The number of lone pairs on the central atoms is added up. Call this n.

Assigning Shapes to Molecules Take these numbers and put them in the following formula:

AXmEn The A in this formula stands for the central atom, the X for the attached atoms, and the E for

the lone pairs. From this code we can name the shape of the molecule.

Linear AX2 is called the linear shape. The bond angle (atoms between the three bonds) in this shape is 180°.

Trigonal Planar and Bent/V-Shaped

The shape associated with AX3 is trigonal planar because it resembles a flat triangle. The shape associated with AX2E1 is called bent/V-shaped.

When naming a shape we pretend to only “see” the atoms, not the lone pairs. That is why we have two different shape names for the same geometry.

Both of these shapes have bond angles of approximately 120°, which may vary slightly in some molecules.

Tetrahedral Geometries

The shape associated with AX4 is called tetrahedral, with all atoms at the four points of a pyramid.

Replacing an atom with an electron pair gives AX3E1, which is called trigonal pyramidal. Changing another atom to a lone pair gives the the AX2E2 code, which is called the bent/V-

shaped shape. This is the same shape name as we saw earlier, but it has a different geometry and

different bond angles. The bond angles associated with these three are approximately 109.5°.

Trigonal Bipyramidal and Octahedral Shapes

The AX5 shape is called trigonal bipyramidal, and resembles two pyramids sharing a common face.

The bond angles associated with this shape are both 90° and 120°, depending on where you look.

The AX6 shape is called octahedral. The bond angles for this shape are 90°.

Other shapes within these geometries are not included in this course, but they do exist.

Dipole Moments

Molecules which have fairly electronegative atoms may have a dipole moment. Since electrons are drawn towards more electronegative atoms, a “partial negative charge”

(δ-)develops in this region of the molecule, and a partial positive charge (δ+)develops on the other side.

Consider CH3F..

Dipole Moments

In thinking about dipole moments, we have to consider the three-dimensional structure of the molecule.

Dipoles can cancel each other out, if the same type of bonds are oriented in opposite directions.

Example

Example: Which of these molecules have dipole moments? Which, if any, have polar bonds but no dipole moment?

CCl4 NH3 H2O CO2

Chemistry 120 Molecular Geometry

Molecular Geometry AXmEn Code

Approximate Bond Angles (degrees)

Hybridization of Central Atom*

2-dimensional Representation


Linear AX2 180 sp AX X AX X

Trigonal Planar AX3 120 sp2

AX X

X

AX X

X

Bent/V-shaped AX2E1 <120 sp2 AX X A

X X

Tetrahedral AX4 109.5 sp3 AX X

X

X

A

X

X X

X

Trigonal Pyramid AX3E1 <109.5 sp3A

XX

X

A

X X

X

Bent/V-shaped AX2E2 <109.5 sp3 A

X X A

X

X

In the AXE codes, m is the number of atoms bonded to the central atom, and n is the number of lone pairs on the central atom (do not include lone pairs on other atoms). Hybridization is an important topic which will be introduced in CHEM 130. Although you do not need this information for CHEM 120, it will be useful for you when you take General Chemistry.

Page 1




Hybridization of Central Atom



Trigonal Bipyramid AX5 90 & 120 dsp3A

X

X X

X X

X AX

X

X

X

See-saw AX4E1 90 &120 dsp3 AX X

X X A

X

X

X

X

T-shaped AX3E2 90 dsp3AX X

X

X A

X

X

Linear AX2E3 180 dsp3 AX X

A

X

X

Note: None of the shapes on pages 2 or 3 of this handout will be tested in CHEM 120, as they all have central atoms which exceed an octet. These are included to assist you in your later studies in General Chemistry.

Page 2




Hybridization of Central Atom



Octahedral AX6 90 d2sp3 A

X

XX

X

X

X

AX

X X

X

X

X

Square Pyramid AX5E1 90 d2sp3

A

X

X

X

X

X

AX

X X

X

X

Square Planar AX4E2 90 d2sp3 AX

X

X

X

A

X

X X

X

Note: None of the shapes on pages 2 or 3 of this handout will be tested in CHEM 120, as they all have central atoms which exceed an octet. These are included to assist you in your later studies in General Chemistry.

Page 3

Chemistry 120 Chapter Six Notes

1

Chapter Six: Nomenclature—The Names of Chemical Compounds Types of Ions

Ions are classified according to how many atoms they contain If an ion is derived from a single atom, it is called a monatomic ion.

Examples include Na+, O2-, and Pb4+. Ions which are derived from two or more atoms are called polyatomic ions.

Polyatomic Ions

Polyatomic ions can be described as molecules which have collectively gained or lost electrons to become an ion.

Being molecules, the ions themselves are held together by covalent bonds. From this, you would expect the atoms in a polyatomic ion to all be nonmetals.

Examples: NO3-, ClO4

-, NH4+

However, some polyatomic ions contain metals which are able to covalently bond. Examples: Cr2O7

2-, MnO4-

All but one of the more common polyatomic ions are anions. The ammonium ion, NH4

+, is the only common polyatomic cation. The common polyatomic ions must be memorized, and you must learn to recognize them in a

formula on sight.

1- anions OH- hydroxide NO3

- nitrate CN- cyanide ClO3

- chlorate OCN- cyanate BrO3

- bromate SCN- thiocyanate IO3

- iodate MnO4

- permanganate C2H3O2- acetate

2- anions CO3

2- carbonate SO42- sulfate

CrO42- chromate Cr2O7

2- dichromate C2O4

2- oxalate SiO32- silicate

S2O32- thiosulfate O2

2- peroxide 3- anions

PO43- phosphate BO3

3- borate


2

Classifications of Compounds

When naming inorganic compounds, two different naming systems are used, depending on the type of compound you are considering.

Binary nonmetals (sometimes called molecular compounds) contain atoms from exactly two different nonmetals.

Examples: NO, CO2, CO, SO2, H2S. Ionic compounds are composed of a cation and an anion.

For the purposes of this course, if a compound is not a binary nonmetal, it is an ionic compound.

The ions may be monatomic and/or polyatomic. Examples: NaCl, CaCl2, NH4Cl, CaCO3, Na3PO4.

Examples

Classify each of these compounds as binary nonmetals or as ionic compounds. KBr

Na2CO3

NaCN

N2O4

NaH2PO4

P2O5

NH4NO3 Naming Binary Nonmetals

Nonmetals can bond with one another in many possible combinations. For example, nitrogen and oxygen make compounds such as NO, NO2, and N2O.

Although these compounds contain the same elements (and may even contain them in the same ratio), these chemicals have considerably different chemical and physical properties.

Therefore, each must be given its own distinct name to distinguish it. The compound is named by putting prefixes before the name of the element to indicate how

many atoms of each type are in the molecule. The ending of the name of the second element is changed to “-ide”.

prefix[1st element] prefix[2nd element]


3

Naming Binary Nonmetals The first ten prefixes must be memorized:

Number of Atoms in Molecule Prefix

1 mono- 2 di- 3 tri- 4 tetra- 5 penta- 6 hexa- 7 hepta- 8 octa- 9 nona- 10 deca-

For example, a compound with formula S2O3 would be named disulfur trioxide. Exceptions:

Do not use the prefix “mono-” with the first element in the formula; just stating the name of the element implies that there is only one of it in the formula.

For example, CO2 is carbon dioxide, not monocarbon dioxide. What is the name of SO3?

If the last letter of the prefix is an “a” or an “o”, and the first letter of the element after the prefix is an “o”, drop the last letter of the prefix; it looks and sounds awkward.

Example: CO is carbon monoxide, not carbon monooxide. P2O5 should be named diphosphorus pentoxide, which is preferable to diphosphorus

pentaoxide (which is technically acceptable).

Common Names In three cases, the naming rules are totally ignored, and the common name is used. These three are

H2O, which is always called water. NH3, which is always called ammonia. PH3, commonly known as phosphine.

Other common names do exist, but are omitted here (they are less common). Hydrogen-Containing Formulas

When hydrogen appears first in the formula of a binary nonmetal, do not use prefixes at all. Simply name the compound without them.

There will be as many hydrogens bonded to the other nonmetal as are necessary to give it a complete octet.

So, H2S(g) is named hydrogen sulfide, and HCl(g) is named hydrogen chloride. Note that the (g) means “gas”, an important fact which must be specified in the formula

for these types of compounds.

If the hydrogen-containing binary nonmetal is dissolved in water, a different name is used.


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In the formula, the notation (aq) is used, meaning aqueous, to indicate that it is dissolved in water.

These compounds are named as acids. These are not to be confused with oxyacids, which will have a different naming system.

Naming Acids of Binary Nonmetals

The prefix “hydro-” is put before the name of the element. The ending of the name of the element is changed to “-ic”, followed by the word “acid”.

Examples: HCl(aq) is hydrochloric acid HBr(aq) is hydrobromic acid HI(aq) is hydroiodic acid H2S(aq) is hydrosulfuric acid CAUTION: Do not confuse this with sulfuric acid, which is H2SO4!

Examples Name each of the following: N2O5

SO2

HF(aq)

HF(g)

H2Se(aq)

NH3

XeF4

H2S(aq)

Naming Ionic Compounds The naming system for ionic compounds is different than that for binary nonmetals. Most importantly: The prefixes used for binary nonmetals (mono, di, tri, etc.) are never used

to name ionic compounds. This is, by far, the most common mistake made by students in naming compounds.

Naming Individual Ions: Cations

Some cations always have the same charge in virtually all compounds. These include

Group I cations, which only form 1+ ions Ex.: Li+, Na+, K+, Rb+, Cs+

Group II cations, which only form 2+ ions Ex.: Mg2+, Ca2+, Sr2+, Ba2+

Group III cations, which only form 3+ ions Al3+, (Ga3+ if it is forming an ionic compound) Note that this is not true for elements below Ga


5

Three additional common metals also only form one ion. These must be memorized. They are

Zinc: Zn2+ Cadmium: Cd2+ Silver: Ag+

For all other metals, the charge must be specified in naming the cation.

The name of the individual cation is stated by simply stating the name of the metal, followed by the word “ion”.

Examples: Na+ is the sodium ion. Ba2+ is the barium ion. Zn2+ is the zinc ion.

The new method for naming all other cations places the charge of ion in Roman numerals in parenthesis after the name of the metal.

Cu2+ is the copper (II) ion. Cu+ is the copper (I) ion. Pb4+ is the lead (IV) ion. Pb2+ is the lead (II) ion.

Remember, you only use Roman numerals for those ions that form multiple charges. The old names for many ions are still in use and must be learned. The names of the ions must be memorized, noting the following:

The ion possessing the higher charge ends in “-ic” That possessing the lower charge ends in

“-ous” Ex.: Fe3+ is the ferric ion, Fe2+ is the ferrous ion.

Ion Common Name Ion Common Name Cu2+ cupric ion Cu+ cuprous ion Fe3+ ferric ion Fe2+ ferrous ion Pb4+ plumbic ion Pb2+ plumbous ion Sn4+ stannic ion Sn2+ stannous ion Hg2+ mercuric ion Hg2

2+ mercurous ion Cr3+ chromic ion Cr2+ chromous ion Co3+ cobaltic ion Co2+ cobaltous ion Mn3+ manganic ion Mn2+ manganous ion

Examples Name the following cations: Rb+

Fe3+

Hg2

+

Hg22+

Cd2+

Ca2+


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Naming Individual Ions: Anions The charge on a monatomic anion is simply 8 minus the group number of the element.

For example, nitrogen is in group 5, so it will only make a N3- ion. (8-5=3) Monatomic anions are named by simply changing the ending of the element’s name to “ide”. Examples:

S2- is the sulfide ion O2- is oxide ion P3- is the phosphide ion

CAUTION: Do not confuse these monatomic ions (nitride, sulfide, phosphide, etc.) with the similar-sounding oxoanions (nitrate, sulfate, phosphate, etc.)

Naming Ionic Compounds To name the ionic compound, simply state the cation, followed by the anion. Do not use the

word “ion” in the name of a neutral compound. Examples:

NaCl is sodium chloride CaBr2 is calcium bromide CuO is copper (II) oxide, or cupric oxide Cu2O is copper (I) oxide, or cuprous oxide

Again, notice that we never use the prefixes used for binary nonmetals (mono-, di-, tri-, etc.) when naming ionic compounds!

This is because ionic compounds always form predictable ratios, as we have already seen. Na+ and Cl- can only come together as NaCl. Cu2+ and O2- come together as CuO. Cu+ and O2- come together as Cu2O.

Recall that the total charges of the cations and anions must neutralize.

Examples Name each of the following ionic compounds: AgCl

KNO3

PbCl2

ZnBr2

NH4NO3

Hg(ClO3)2

CdSO4

SnCl4

Oxoanions

The oxoanions include many polyatomic anions which contain oxygen. OH- and O2

2- are not considered oxoanions. So far, we have only encountered those oxoanions whose names end in “-ate.” Related oxoanions have different numbers of oxygen atoms but the same charge. Examples:

SO42- is the sulfate anion, SO3

2- is the sulfite anion. ClO3

- is the chlorate anion, ClO4- is the perchlorate anion.


7

Let us consider the chlorate ion, which we know has formula ClO3-.

It is very important that we know how many oxygen atoms are in the “-ate” ion. If the ion has one more oxygen in its formula than that of the “-ate” ion, add the prefix “per-” to

its name. So ClO4

- is perchlorate, SO52- would be persulfate, etc.

Again, consider the chlorate ion, ClO3-.

Taking away one oxygen from the “-ate” anion changes its ending from “-ate” to “-ite.”

Therefore, ClO2- is the chlorite ion, and SO3

2- is the sulfite ion. Taking away two oxygens from the “-ate” anion changes its ending from “-ate” to

“-ite,” and you must add the prefix “hypo-” So, ClO- is the hypochlorite ion, and SO2

2- is the hyposulfite ion. Oxoanions Containing Hydrogen

Some common polyatomic anions have an H+ “within” their formula. In these ions, the H is always listed first, and the normal charge of the ion is changed by

+1. The name of the ion is changed by adding one of the following:

The prefix “bi-”, or The word “hydrogen” before the name of the anion (sometimes you might see

“monohydrogen” instead.) For example, HCO3

- is commonly called the “bicarbonate ion” or the “hydrogen carbonate ion.”

What would HSO3- be called?

Occasionally you may also see two hydrogen atoms added to the ion. This is rare, except in the case of phosphate.

This would increase the charge of the original anion by +2. The word “dihydrogen” is added to the name of the anion. So H2PO4

- is the “dihydrogen phosphate” ion. KH2PO4

is called “potassium dihydrogen phosphate.” NOTE: In all these examples, H+ is not the cation; we should consider it an “attachment” of

the anion for naming purposes. Examples Name each of the following compounds: KIO4

NaHCO3

Al2(SO3)3

CuHPO3

LiH2PO4

NaClO

Co(NO2)3


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Oxoacids

Oxoacids are compounds which contain exactly as many H+ ions as are necessary to cancel out the negative charge of an oxoanion.

For example, SO42- would need two H+ ions to balance out the 2- charge of sulfate; the

oxoacid has the formula H2SO4. Note that the H+ ions are

the only cations in the formula, and always listed first in the formula.

Naming Oxoacids

The naming rules for oxoacids are similar to those for oxoanions. The name of the oxyacid depends on how many oxygen atoms are in the corresponding

oxoanion. For those oxoanions which end in “-ate”, change the ending to “-ic acid.”

So, CO32- is the carbonate ion

H2CO3 is carbonic acid

HNO3 is named _________________________________ Slightly “weird” cases: H2SO4 is sulfuric acid, and H3PO4 is phosphoric acid.

An oxoanion that begins with “per-” follows the same rules. Simply change the ending to “-ic acid.”

HClO4 is called perchloric acid. HIO4 is called periodic acid.

Any oxoanion that ends in “-ite” has its ending changed to “-ous acid” H2SO3 is called sulfurous acid. HClO2 is called chlorous acid.

The acids of oxoanions which begin with “hypo-” likewise have their ending changed to “-ous acid”.

HBrO would be called hypobromous acid. HNO would be called hyponitrous acid.

Examples Name the following oxoacids: H3PO4

HNO3

HNO2

HIO2

H2CO3

HIO


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The Cyanide Ion

One last random bit of information… The cyanide ion (CN-) often behaves like a halogen, so naming rules for it are similar to those

of F-, Cl-, etc. HCN(g) is called “hydrogen cyanide.” HCN(aq) is called hydrocyanic acid.

Hydrates

A hydrate is an ionic compound which has “trapped” a fixed number of water molecules within its structure

For example, the following compound is called copper (II) sulfate pentahydrate:

CuSO4 · 5 H2O The formula tells us that there are five water molecules in the structure for every unit of

copper (II) sulfate The naming of other hydrates is similar:

State the name of the salt End with a prefix indicating how many water molecules are present, followed by the word

“hydrate” Name each of the following hydrates:

BaCl2 · 2 H2O

MgSO4 · 7 H2O

Final Examples Name each of the following compounds. BrO4

Cu(ClO4)2

H2SO3

HF(aq)

LiHCO3

NH3

N2O

F2

HIO3

Chemistry 120 Monatomic Ions I

Group1A 8A

H+/H- 2A 3A 4A 5A 6A 7A

Li+C4-

carbide(rare)

N3-

nitrideO2-

oxideF-

fluoride

Na+ Mg2+ 3B 4B 5B 6B 7B 8B 1B 2B Al3+ P3-

phosphideS2-

sulfideCl-

chloride

K+ Ca2+ Zn2+ Ga3+ Se2-

selenideBr-

bromide

Rb+ Sr2+ Ag+ Cd2+ I-iodide

Cs+ Ba2+

This periodic table contains the formulas of what I consider to be the most important ionswhich only form one charge. You'll notice a few important details: (1) The charge on almost all of the metal ions on this list is the same as the group number. (2) The charge on a nonmetal is simply its group number minus eight (i.e. oxide is derived from oxygen, which is in group 6; 6 - 8= -2 ) (3) Zinc, cadmium, and silver ions are generally only encountered with the charge indicated. (4) As a cation, hydrogen makes a 1+ ion; when hydrogen makes an ionic bond with a metal the result is a "hydride" ion, which has charge 1-. (5) Boron and beryllium are excluded, as they generally form only covalent bonds. (6) Many other elements form ions with two or more different charges; they have also been excluded.

Chemistry 120 Monatomic Ions II

Group1A 8A

2A 3A 4A 5A 6A 7A

3B 4B 5B 6B 7B 8B 1B 2B

Cr3+

Cr2+Fe3+

Fe2+Co3+

Co2+Cu2+

Cu+

Sn4+

Sn2+

Hg2+

Hg22+

Pb4+

Pb2+

This sheet lists the possible charges of many common elements which have more than one stable charge staYou should know both the new name and old name for each (i.e. copper (I) ion and cuprous ion). Note that the mercury (I) ion is a strange case; it is always found as a pair of atoms with a net charge of 2+ shared between each atom in the pair.

Chemistry 120 Chapter Seven Notes

1

Chapter Seven Formula Calculations Introducing the Mole

The dozen is a unit of quanity If I have a dozen atoms, I have 12 atoms by definition.

The mole(mol) is a very important unit of quantity in chemistry. It is used to count large numbers of atoms, molecules, and other submicroscopic pieces of matter.

If you have 1 mole of something, you have 6.02 × 1023 of it.

Examples How many eggs are in 5.5 dozen eggs?

How many atoms are in 5.5 mole of atoms?

How many hydrogen atoms are in one molecule of water? How many oxygen atoms?

How many hydrogen atoms are in one mole of water? How many oxygen atoms? How many atoms total ?

Practice If you have 18.45 mol of oxygen gas, how many individual molecules do you have? How

many individual atoms?

If a sample of methane contains 7.35 × 1025 atoms, how many moles of methane do you have?


2

Atomic Weights and The Mole

The atomic weights provided on the periodic table can be used in much more convenient units than amu.

The values on the periodic table can be interpreted as grams per mole of the atom. For example, 1 mol of calcium atoms has a mass of 40.08 g; 1 mol of neon atoms has a

mass of 20.18 g. Note that atomic weights are often also called molar masses.

Molecular Weights

The mass of molecules can be calculated by adding up the atomic weights of the individual atoms making up the molecule.

For example, suppose we wanted to know the atomic weight of CO2.

Every 1 mol of CO2 contains_____mole of C atoms and_____moles of O atoms. Calculate the total mass of all of these atoms and add them up to get the molecular

weight of CO2.

Aside: what is the mass of a single molecule of CO2 in amu?

Note that the term formula weight is applied for those compounds which do not form true molecules (ionic compounds like NaCl and CaO).

Examples Determine the molecular weights of glucose (C6H12O6) and acetic acid (CH3CO2H).


3

Percent Composition

One common technique used in specialized chemical laboratories is elemental analysis, which can be used to determine the percent each element in a compound contributes to its mass.

Since this type of data is very common, it is important that we know how to interpret it and put it to practical use.

The percent composition is the percent of the total mass percent of each element in a compound.

To determine its value, we use the following formula for each element in the compound:

Examples What is the percent composition of each element in acetic acid? A 19.82 g chunk of an ore is analyzed and found to have a percent composition of 5.75% silver, 64.33% iron, and the remainder is silicon. What mass of each element is contained in a 4.87 kg sample of this ore?

%100mass molar

element of mass total ncompositio % ×=


4

Percent compositions can also be determined for mixtures, such as alloys.

Suppose that 5.50 mols each of copper and zinc are blended with 2.43 mols of tin to make an

alloy. What is the percent composition of the alloy?

A student determines the mass of a hydrate sample to be 6.873 g. After severely heating the sample for 20 minutes to remove water, he reweighs the sample, finding the mass to be 5.276 g. What percent of the mass of the original hydrate is water?

Molecular and Empirical Formulas The molecular formula of a compound tells us exactly how many atoms of each element are

contained in a compound. In contrast, the empirical formula only tells us the lowest whole-number ratio between each

element For example, glucose has molecular formula C6H12O6 (each molecule of glucose contains 6

carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms) Glucose then has empirical formula CH2O (1 C to 2 H to 1 O is the lowest whole-number

ratio).

Empirical Formulas While the molecular formula ultimately is more useful for most applications, it is often more

difficult to determine than the empirical formula. When a sample is analyzed we can easily determine the percent composition, and from here

find the ratio of the number of atoms of each element. Information on the exact number of atoms in each molecule cannot be found in this way.


5

Determining the Emprical Formula

A simple series of steps can be used to determine the empirical formula of a compound: Find the mass of each element in the compound Convert the masses into moles of each element Express the moles of atoms as the smallest possible whole-number ratio Use the numbers from these ratios in the empirical formula for each element.

Examples A sample of a compound is analyzed, and found to contain 1.61 g of phosphorus and 2.98 g of

fluorine. What is the empirical formula of this compound?

The mass of a piece of iron is 1.62 g. The iron is exposed to oxygen and reacts to form a pure oxide of iron, now with mass of 2.31 g. What is the empirical formula of this oxide?

A compound is found to contain 20.0% carbon, 2.2% hydrogen, with the remainder of the mass derived from chlorine. What is the empirical formula of this compound?


6

Molecular Formula

Given the empirical formula and the molecular weight of a compound, it is possible to determine the molecular formula.

For example, suppose we have a compound with empirical formula CH. Its molecular formula could be CH, C2H2, C3H3, etc.

Now, let’s see what the formula weight would be for each: CH = 13.02 g/mol C2H2 = 26.04 g/mol C3H3 = 39.06 g/mol Notice that each is a multiple of 13.02!

To determine the molecular formula from this information Find the formula weight of the empirical formula Divide the molecular weight by this value (you should get a whole number). Multiply the subscripts in your empirical formula by this whole number.

You now have your molecular formula.

Example An unknown compound is determined to be 91.8% silicon and 8.2% hydrogen. The molar mass

of this compound is found to be 122 g/mol. Determine the molecular formula of the compound.

Chemistry 120 Chapter Eight Notes

1

Chapter Eight Chemical Reactions What is a Chemical Reaction?

A chemical reaction involves the conversion of one or more substances into one or more different substances.

The substance(s) which we begin with are called the reactant(s) The substance(s) which we end with are called the product(s)

Writing Chemical Reactions

Chemical reactions are written with the reactants to the left, the products to the right, and an arrow between them to indicate the change.

Occasionally symbols or values may be written over or under the arrow to indicate the reaction conditions.

An Example: H2 + O2 H2O

Balancing Chemical Equations Consider the last reaction:

H2 + O2 H2O There is a problem with this equation! It indicates that we started with two oxygen atoms, but ended with one. What does this contradict?

We must balance chemical equations, which is to say that there must be equal numbers of each type of atom on either side of a chemical reaction.

To accomplish this, we put coefficients in front of the chemical formulas whose atom numbers we wish to increase.

Note that you may never change the subscripts already in place in a chemical formula! Why?

To balance chemical equations first count the number of each type of atom you have on both sides of the reaction.

Identify any lone elements (as opposed to compounds) in the formulas; you will balance these last.

From here, each equation requires its own logic; by trial and error, you should be able to balance the equation.

The only other real “tip” I can give you on this subject is that practice makes perfect!


2

Examples Balance each of the following chemical reactions:

Mg + O2 MgO HgO Hg + O2

Na + MgCl2 Mg + NaCl

Na3PO4 + BaCl2 Ba3(PO4)2 + NaCl

C6H12O6 + O2 CO2 + H2O

Write the balanced equation that corresponds to each statement

When methane reacts with oxygen gas, carbon dioxide and water vapor are produced.

Calcium metal reacts with ferric oxide, yielding calcium oxide and iron metal.

Treatment of carbon monoxide gas with oxygen gas produces carbon dioxide.

Types of Chemical Reactions There are five main classification types of chemical reactions. Knowing how each of these reaction types “works” gives you the ability to predict what

chemical products will be formed in a given reaction in most cases. Types of Chemical Reactions

Combustion Single Substitution (Displacement) Double Substitution (Displacement) Combination (or Synthesis) Decomposition

Within these five classification types, even more specific reaction types can be identified; these will be described as we come to them.


3

Combustion Reactions

In a combustion reaction, a chemical reacts with oxygen gas, forming various products. In this class, we will only consider the combustion of organic compounds containing C, H, and

sometimes O. In these reactions, the compound reacts with oxygen gas, producing carbon dioxide and

water vapor.

[organic compound] + O2(g) CO2(g) + H2O(g)

Examples of Combustion Combustion of benzene

2C6H6(l) + 15O2(g) 12CO2(g) + 6H2O(g)

Combustion of formaldehyde CH2O(l) + O2(g) CO2(g) + H2O(g)

What is the balanced equation for the combustion of glycerol (C3H8O3(l))?

Single Substitution Reactions In a single substitution reaction, an element replaces another element which is present as an

ion in a compound. There are two types of substitutions we should be aware of:

substitution of one metal with another substitution of one halogen with another

Format of the first type of single substitution reactions:

A + BX AX + B Where A and B are metals/metal ions, and X is an anion.

Here are some examples: Mg + ZnCl2 MgCl2 + Zn Zn + 2AgNO3 Zn(NO3)2 + 2Ag 2Na + 2H2O 2NaOH + H2

Activities of Metals We can use the activity of a metal to describe how readily it loses electron(s) in a reaction.

The more active the metal, the more readily it loses its electrons. A more active metal will displace a less active metal in a single substitution reaction; the

reverse will not occur. In other words, an active metal will force the cation of a less active metal to take its

electrons away from it. We look to the activity series to see the relative activities of the metals.

The activity series should not be memorized, but you should become familiar with trends within it.

Displacing Hydrogen

Notice that many metals can displace the H+ from acids, changing it into H2. Notice that the charged hydrogen ion is transformed into the neutral hydrogen molecule.

Some very active metals can even displace H+ from water, leaving OH- behind. In these cases, think of water as HOH.


4

When Does the Reaction “Go”?

So let’s consider one single substitution that works well… Zn + 2AgNO3 Zn(NO3)2 + 2Ag

Since zinc is more active than silver (higher on the activity series), zinc will displace silver. Consider the reverse reaction…

Ag + Zn(NO3)2 no reaction Since silver is less active than zinc, it cannot displace it; therefore, no reaction can occur.

Oxidation & Reduction: An Introduction

Single substitution reactions involve an oxidation and a reduction. When a metal atom or ion is oxidized, it loses one or more electrons, resulting in an

increase in its charge. Example: Zn Zn2+ + 2e-

When a metal ion is reduced, it gains one or more electrons, causing its charge to decrease.

Example: Ag+ + e- Ag An oxidation reaction is always accompanied by a reduction reaction. Redox is the common term for these reactions.

A Look at Some Redox Reactions

Consider the following single substitution reaction: 3Mg + 2FeCl3 3MgCl2 + 2Fe

What is being oxidized?

What is being reduced?

Try this reaction: 2Na + 2H2O 2NaOH + H2

What is being oxidized?

What is being reduced?

Single Substitution Reactions: Substitution of Halogens

Anions derived from halogens (Group VII) can be displaced by a more active halogen. Activity of the halogens decreases down the group:

F2 > Cl2 > Br2 > I2

most active least active An example:

Cl2 + 2NaI 2NaCl + I2 Will the reverse reaction proceed?


5

Examples Predict the products of the following reactions. Write “no rxn.” if none is expected to occur.

Ba + CoBr3

Ni + NaCl

I2 + KF

Li + H2O

Ni + HNO3

Double Substitution Reactions The general format for a double substitution reaction involves two ionic compounds trading

their anions: AX + BY AY + BX

AX and BY are both aqueous solutions, meaning both ionic compounds are dissolved in water.

These reactions will proceed only if at least one of the following is true: A solid precipitate is produced A covalent (molecular) compound is produced, such as H2O, H2, CO2, SO3, etc.

Examples of Double Substitution Reactions AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)

3Ba(OH)2(aq) + 2FeCl3(aq) 3BaCl2(aq) + 2Fe(OH)3(s)

HCl(aq) + NaOH(aq) H2O(l) + NaCl(aq)

Solubility of Compounds

Solubility is best described as the degree to which a compound will dissolve in a solvent (usually water).

A compound that is soluble will dissolve to a significant extent. Compounds that are insoluble will not dissolve, remaining solid in solution.

A reaction which produces an insoluble product can be described as a precipitation reaction; the product “falls out” of the solution like rain precipitates.


6

Solubility

It is important to know whether or not some common chemicals are soluble or not in order to predict how a reaction will proceed.

You should know the following solubility rules (in water) by memory: All nitrates and acetates are soluble. All salts of Group I cations (Li+, Na+, etc.) and ammonium are soluble. All chlorides, bromides, and iodides are soluble, except those of Ag+, Pb2+, and Hg2

2+. All hydroxides are insoluble except those of Group I, NH4

+, Ba2+, Sr2+, and Ca2+. Ca(OH)2 is only slightly soluble.

Unstable Products of Reactions

When produced in a reaction, some compounds will immediately decompose into other products.

Carbonic acid (H2CO3) will decompose into CO2(g) and H2O(l). Ammonium hydroxide (NH4OH) will decompose into NH3(aq) and H2O(l). Sulfurous acid (H2SO3) will decompose into SO2(g) and H2O(l).

Notice that each produces water and a compound formed by the atoms left after water has

been removed from the starting formula. If you produce any of these three compounds in a reaction, cancel it out and replace it with

the decomposition products.

Electrolytes When an ionic compound dissolves in water, it dissociates (separates) into ions. A solution containing ions is capable of conducting an electric current. The term electrolytes is applied to these ions. A solution formed from a soluble ionic compound conducts the current very strongly; the

compound is called a strong electrolyte. A solution formed from an only slightly soluble ionic compound conducts the current weakly;

the compound is called a weak electrolyte. A solution derived from a non-ionic (i.e. covalent) compound will not conduct an electric

current; the compound is called a nonelectrolyte.

Acids and Bases You have heard the term acid applied to several compounds.

All of these compounds mentioned so far contain the H+, called a proton. We will look more in depth at the definitions of acids and bases later, but for now we will use

the following: Acids are compounds which give up H+ ions. Bases are compounds which accept H+ ions.

For now, we will limit our study of bases to some hydroxide-containing compounds and NH3.

Strong Acids Strong acids are strong electrolytes, meaning they completely break up into ions in solution.

So, HCl, which is a strong acid, breaks up into H+ ions and Cl- ions in solution. There are six strong acids you must know:

HCl HBr HI HNO3 H2SO4 HClO4

Note that H2SO4 dissociates into H+ and HSO4-.


7

Example: If you were to “look” into a beaker of the following solutions, what ions, solids, or molecules would you expect to see (besides water)? a. HNO3

b. HBr c. HClO4

Weak Acids Weak acids are weak electrolytes, meaning that relatively few acid molecules will break up

into separate ions. Some samples of weak acids are

H3PO4 H2CO3 HCH3CO2

Strong and Weak Bases Similar definitions apply to strong bases. The strong bases are the Group I hydroxides (LiOH, NaOH, KOH, etc.) and some Group II

hydroxides: Ba(OH)2, Sr(OH)2, and Ca(OH)2. The weak bases include Mg(OH)2 and NH3(aq) (which can be treated as

NH4OH(aq) for now).

Neutralization Reactions Neutralization reactions are a subclass of double substitution reactions. The general form of this reaction is

acid + base water + salt where the salt is any ionic compound.

This can be further simplified as follows: H+ + OH- H2O

Strong acids and strong bases completely neutralize each other. A strong acid will neutralize a weak base, and a strong base will neutralize a weak acid. The reactions of weak acids with weak bases is somewhat more complicated and will not be

considered in this class.

Examples HBr(aq) + KOH(aq) H2O(l) + KBr(aq)

2HNO3(aq)+ Ba(OH)2(aq) 2H2O(l) + Ba(NO3)2(aq)


8

Write the balanced equations showing:

a) the neutralization of lithium hydroxide by perchloric acid.

b) the reaction of sulfuric acid with strontium hydroxide.

c) an ammonia solution is treated with hydrobromic acid.

Combination Reactions

In a combination (or synthesis) reaction, two chemicals combine into one new chemical. These reactions have the general form

A + B C It will not always be possible to predict the products of combination reactions at this level of

preparation, so we will study a few general cases.

Oxide Formation Metals often react with oxygen to form a metal oxide.

Ex. 4Na + O2 2Na2O Nonmetals often react with oxygen to form a nonmetal oxide.

Ex. C + O2 CO2 It is often difficult to predict the products in this case, as CO was another possible oxide

you might have considered.

Reactions of Oxides Metal oxides often react with water to form metal hydroxides.

Ex. Na2O + H2O 2NaOH

Nonmetal oxides often react with water to form oxyacids. Ex. CO2 + H2O H2CO3

P4O10 + 6H2O 4H3PO4

SO3 + H2O H2SO4 N2O5 + H2O 2HNO3

Metal oxides and nonmetal oxides often combine to form a salt. Na2O + CO2 Na2CO3 CaO + SO3 CaSO4

Decomposition Reactions Decomposition reactions are simply the reverse of combination reactions. Their general form is

Z X + Y Like combination reactions, predicting the products of these reactions is often difficult. For now, simply decompose a given compound into two products which could have produced

it from a method we discussed earlier. This is very oversimplified, and not always correct, but it is the best we can do at this

level.


9

Examples

2HgO 2Hg + O2

Na2CO3 Na2O + CO2 Li2SO4 Li2O + SO3

Heat in Chemical Reactions

Virtually every chemical reaction results in either a gain or loss of heat to the reaction system A reaction which produces more heat than it consumes is called an exothermic reaction A reaction which consumes more heat than it produces is called an endothermic reaction The chemical reaction written below describes an exothermic reaction:

CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) + 890 Kj

For an endothermic reaction, the heat would appear with the reactants in the chemical equation: PCl5 + 92.9 kJ PCl3 + Cl2

Although this method is convenient for indicating heat changes, it often leads to confusion in balancing equations

A better method for indicating energy changes will be presented in Chem 130

Heat and Chemical Bonds Energy is always consumed in the breaking of chemical bonds. The opposite is also true: energy is always released when new chemical bonds are formed. Ultimately, whether a reaction is exothermic or endothermic therefore depends on the number and type of

bonds being broken and formed in a chemical reaction Activation Energy and Energy Diagrams

In order for a chemical reaction to begin, some amount of energy specific to that reaction must be supplied to start the reaction process

The amount of this initial investment of energy is called the activation energy The activation energy, as well as the energy changes in chemical reactions, may be illustrated by energy

diagrams

Activity Series of Metals

Solubilities of Selected Ionic Compounds in Water In general, all ionic compounds containing ammonium (NH4

+) and Group IA cations (Li+, Na+, K+, etc.) are soluble.

Most Li Displaces hydrogen from Active Rb cold water, steam, and acids

K Cs Ba Sr Ca Na Mg Displaces hydrogen from Al steam and acids Mn Zn Cr Fe Cd Displaces hydrogen from Co acids Ni Sn Pb H2 Cu Cannot displace hydrogen from Hg cold water, steam, or acids Ag

Least Pt Active Au

nitrates acetates chlorates

all are soluble

chlorides bromides iodides

all are soluble except Ag+, Pb2+, & Hg22+

sulfates

chromates

all are soluble except Ag+, Pb2+, & Hg22+, Hg2+,

Ba2+, Sr2+, & Ca2+ note: CaCrO4 is soluble, CuCrO4 is not

oxides hydroxides

none soluble except Group IA, NH4+,

Ba2+, Sr2+, & Ca2+

sulfides none soluble except Group IA, NH4+, and

Group IIA carbonates phosphates sulfites

none soluble except Group IA & NH4+

Activity Series of Metals—Test Version

Solubilities of Selected Ionic Compounds in Water Note that the solubility rules for those compounds which you were to have memorized have been deleted from the table.

Most Li Displaces hydrogen from Active Rb cold water, steam, and acids

K Cs Ba Sr Ca Na Mg Displaces hydrogen from Al steam and acids Mn Zn Cr Fe Cd Displaces hydrogen from Co acids Ni Sn Pb H2 Cu Cannot displace hydrogen from Hg cold water, steam, or acids Ag

Least Pt Active Au

chlorates all are soluble

sulfates

chromates

all are soluble except Ag+, Pb2+, & Hg22+, Hg2+,

Ba2+, Sr2+, & Ca2+ note: CaCrO4 is soluble, CuCrO4 is not

sulfides none soluble except Group IA, NH4+, and

Group IIA carbonates phosphates sulfites

none soluble except Group IA & NH4+

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Chemistry 120 Net Ionic Equation Notes

1

Net Ionic Equations Revisiting “aqueous”

The term aqueous is mainly a visual one; when you add a chemical to water and it appears to dissolve, we say an aqueous solution has been formed.

However, we must be more specific: For ionic compounds, including strong acids, aqueous means that the compound has

dissociated into ions. For weak acids and other covalent compounds, aqueous means that the compound has

dissolved, but not been broken down into smaller parts. It is merely “surrounded” by water molecules.

Looking Back Consider the following reaction:

NaCl(aq) +KNO3(aq) Since both the products of this reaction are aqueous, and since no molecular compound is

formed, we say that there is no reaction. However, since we did bring these two solutions together, why is it incorrect to say that the

products are KCl(aq) and NaNO3(aq)?

Earlier, we defined a chemical change as one that involves transforming one or more chemical compounds into one or more different products.

By combining two solutions which do not make an insoluble product or a molecular compound, we fail to create a new product, as we shall now prove.

Total Ionic Equations Total molecular equations show all species in solution for a given reaction.

The term species refers to the specific ions and/or compounds in the solution. Examples of species:

For NaCl(aq), Na+(aq) and Cl-(aq).

For HNO3(aq), H+(aq) and NO3

-(aq).

For H3PO4(aq), it remains H3PO4(aq). Some simple guidelines for writing down the species.

Strong electrolytes, including strong acids and bases, dissociate into their ions. Weak electrolytes, including weak acids and bases, generally do not dissociate, so the

species is just the original compound. Nonelectrolytes do not dissociate, so the species is again the same as the original

compound.

Total Ionic Equations Example:

NaCl(aq) + AgNO3(aq) AgCl(s) + NaNO3(aq)

Total Molecular Equation: Na+

(aq) + Cl-(aq) + Ag+(aq) + NO3

-(aq) AgCl(s) + Na+

(aq) + NO3-(aq)

Chemistry 120 Net Ionic Equation Notes

2

Spectator Ions

A spectator of sports is someone who watches the game from the sidelines, but does not participate.

Similarly, in chemical reactions, spectator ions “hang out” in a solution but do not actively participate in the reaction itself.

In other words, any ion which is both on the reactants and products side of a reaction is a spectator ion, for it has not undergone a chemical change.

The ions’ main purpose is to maintain constant charge in the solution.

Net Ionic Equations Net ionic equations only show those chemicals which participate in the reaction. Spectator

ions are not included. To write a net ionic equation, first write down the total ionic equation. Then, cancel anything which appears identically on both sides of the reaction. Consider the total ionic equation

Na+


-(aq) AgCl(s) + Na+

(aq) + NO3-(aq)

Now, factor out the spectator ions

Na+


-(aq) AgCl(s) + Na+

(aq) + NO3-(aq)

The net ionic equation is left over.

Ag+

(aq) + Cl-(aq) AgCl(s) Examples Write chemical equations, total ionic equations, and net ionic equations for the following:

a. a solution of barium chloride reacts with a sodium sulfate solution.

b. sodium hydroxide is neutralized by hydrochloric acid.

c. phosphoric acid is reacted with potassium hydroxide.

d. solutions of sodium nitrate and ammonium chloride are brought together.

e. a piece of sodium metal is dropped into cold water.

Type of Substance Type of Electrolyte Ionizes in Water Examples How to Write in Net Ionic Equations

soluble ionic compounds strong completely

NaCl(aq) Na+(aq) + Cl-(aq)

KNO3(aq) K+(aq) + NO3

-(aq)

(NH4)2SO4(aq) 2 NH4+(aq) + SO4

2-(aq)

insoluble ionic compounds nonelectrolyte

(not all compounds in these

categories are nonelectrolytes, but you should not worry about exceptions)

not at all

AgCl(s) AgCl(s)

PbI2(s) PbI2(s)

BaSO4(s) BaSO4(s)

covalent/molecular compounds

H2O(l) H2O(l)

C6H12O6(aq) C6H12O6(aq)

H2(g) H2(g)

strong acids and

strong bases strong completely

HCl(aq) H+(aq) + Cl-(aq)

HNO3(aq) H+(aq) + NO3

-(aq)

NaOH(aq) Na+(aq) + OH-

(aq)

weak acids and

weak bases weak only slightly

H3PO4(aq) H3PO4(aq)

HC2H3O2(aq) HC2H3O2(aq)

Mg(OH)2(s) Mg(OH)2(s)

Net Ionic Equations Write (1) full formula, (2) total ionic, and (3) net ionic equations for each of the following chemical combinations. Write “no rxn.” To the right of the arrow if you expect no reaction to occur. 1. A piece of potassium is dropped in water.

(1) Full: (2) Total (3) Net Ionic

2. Aqueous solutions of iron (III) sulfate and ammonia are mixed. (1) Full: (2) Total (3) Net Ionic

3. Silver is added to a cupric bromide solution. (1) Full: (2) Total (3) Net Ionic

4. A sodium fluoride solution is mixed with nitric acid. (1) Full: (2) Total (3) Net Ionic

5. Aqueous solutions of potassium phosphate and cobalt (II) chloride are mixed. (1) Full: (2) Total (3) Net Ionic

6. Aqueous ammonia is mixed with hydoiodic acid. (1) Full: (2) Total (3) Net Ionic

7. Phosphoric acid is mixed with an aqueous solution of lithium hydroxide. (1) Full: (2) Total (3) Net Ionic

8. Carbon dioxide is passed into aqueous barium hydroxide. (1) Full: (2) Total (3) Net Ionic

9. Hydrobromic acid is mixed with a sodium carbonate solution. (1) Full: (2) Total (3) Net Ionic

10. Nitric acid is mixed with aqueous potassium bicarbonate. (1) Full: (2) Total (3) Net Ionic

11. Rubidium oxide is added to water. (1) Full: (2) Total (3) Net Ionic

12. Aqueous solutions of lithium acetate and tin (II) chloride are mixed. (1) Full: (2) Total (3) Net Ionic

13. Hydrochloric acid is added to a lithium sulfite solution. (1) Full: (2) Total (3) Net Ionic

Chemistry 120 Chapter Nine Notes

1

Chapter Nine: Stoichiometry Conservation

Recall that, in any chemical reaction, that matter is always conserved. Consider the reaction of zinc with sulfur:

Zn(s) + S(s) ZnS(s) From this reaction we can imply that, assuming the reaction works perfectly, every one atom

of zinc reacts with one atom of sulfur to produce one unit of zinc sulfide. More conveniently, we can say that one mole of zinc atoms (1 mol Zn) reacts with one mole

of sulfur atoms (1 mol S), producing one mole of zinc sulfide (1 mol ZnS)

Common Misconceptions Is it true that if I start a reaction with 5.0 g of zinc and 5.0 g of sulfur that I will end the reaction

with 10.0 g of matter?

Is it true that if I start a reaction with 5.0 g of zinc and 5.0 g of sulfur that I will end the reaction with 10.0 g of zinc sulfide?

In predicting the quantities of products created during a reaction, it is ultimately the number of atoms/molecules reacting that we must consider, not the mass.

The most convenient way to do this is to consider the moles of reactants participating in a reaction.

Coefficients of Reactions The coefficients in chemical equations are useful in this context because they tell us how

many moles of reactants are required to produce a certain number of moles of products. We can say that one mole of zinc reacts with one mole of sulfur, producing 1 mole of zinc

sulfide. The three are equivalent for the purpose of this reaction. 1 mol Zn 1mol S 1 mol ZnS

We can convert between moles of reactants and products using dimensional analysis, the same technique we used for unit conversions.

To convert between moles of zinc and moles of zinc sulfide, we use the following ratio:

Basic Problems Let’s consider the following reaction:

2H2(g) + O2(g) 2H2O(l) To produce 2 mol of water we need to combine_____mol H2 with_____mol O2.

How many moles of each reactant are required to produce 10 moles of water?


2

How many moles of water can be produced at most from 13 moles of hydrogen? How many

moles of oxygen would I need to accomplish this?

In the reaction of aqueous solutions of hydrochloric acid and barium hydroxide:

a. How many moles of each reactant do you need to produce 5.75 mol water? b. What is the maximum number of moles of water you can produce with 18.0 mol HCl? How many moles of barium hydroxide would you need to accomplish this?


3

Using the Mass-Mole Relationship

Recall that the molar mass relates the mass of a compound (or element) to the number of moles.

Therefore, if we know the mass of a given reactant, we can easily determine the number of moles of the reactant, and from there the number of moles of product produced.

Remember, you must use the mole-mole relationship to carry out a conversion; the mass alone is not sufficient!

Mass-Mole Problems Consider the reaction of sodium metal with chlorine to make sodium chloride.

a. How many moles of sodium are contained in 15.50 g of sodium?

b. Assuming we have excess (that is, more than enough) chlorine, how many moles of sodium chloride could be produced?

c. How many grams of sodium chloride is this?

When 39.75 g of sodium is reacted with excess chlorine, how many grams of sodium chloride can be produced?


4

When 146.7 g of chlorine is reacted with excess sodium, how many grams of sodium chloride can be produced?

How many grams of each reactant do you need to produce 100.00 g of sodium chloride?

Another Example: Ethanol has molecular formula CH3CH2OH. a. How many grams of carbon dioxide and water can be produced if 55.0 mL of ethanol is

combusted? Note: The density of ethanol is 0.789 g/mL.


5

b. How many oxygen molecules are required to react with this amount of ethanol?

The Ice Cream Sundae Problem

Suppose we are going to make ice cream sundaes. We decide that each of our sundaes must contain

2 scoops of ice cream 50 mL of chocolate syrup 1 cherry

How many sundaes can I make if I have 8 scoops of ice cream, 13 cherries, and 300 mL of chocolate syrup available?


6

Limiting Reactants

The problem we had with our ice cream sundaes is the same type of problem we face in chemical reactions.

In virtually all reactions, we will have a disproportionate amount of reactants. The reactant which will produce less product is called the limiting reactant; in all cases, it

predicts the amount of product formed. The other reactants are said to be in excess, meaning there is more than enough of it. The key to identifying a limiting reagent problem is that it will give you the amounts being

reacted of more than one reactant.

Suppose you with to synthesize zinc sulfide.

You combine 50. grams of zinc with 50. grams of sulfur. What mass of zinc sulfide is predicted to be produced from the zinc?

What mass of zinc sulfide is predicted to be produced from the sulfur?

Which reactant is the limiting reactant?

What mass of zinc sulfide is produced by the reaction of 75.8 g of zinc with 60.5 g of sulfur?


7

Using the Tabulation Method

The tabulation method is an alternative method for solving some stoichiometry problems While most problems do not require using this method, it is essential for more complex

problems, especially those involving equilibrium calculations This method is most helpful when you need to know the amounts of each chemical present at

the end of the reaction

Use the tabulation method to solve this problem: 3.0 mols of N2(g) react with 8.0 mols of H2(g) to produce NH3(g). If the reaction goes to

completion (that is, until you run out of the limiting reactant), how many moles of each substance will be present in the flask at the end of the reaction?

Consider the hydrogenation of propene(C3H6 ): C3H6 + H2 C3H8 78.24 g of propene is treated with 25.0 g of hydrogen. a. What mass of propane (C3H8) is produced?


8

b. Which reactant was in excess, and what mass of it remains after the reaction is

completed?

Percent Yield So far, we have assumed that all reactions will produce exactly the amount of products that

our calculations will predict. The theoretical yield is the amount of a product you expect to get at the end of a reaction. In reality, we never actually get 100% of this amount.

Why might this be so? The amount of a product that is really produced at the end of a reaction is called the actual

yield. The percent yield gives us an idea of how much of a product we actually produced compared

to the amount that should have been produced.

In general, the closer the percent yield is to 100% percent, the better the reaction

Example 8.0 g of hydrogen is combined with excess oxygen, producing 58.5 g of water. What is the

percent yield of this reaction?

%% 100 yieldltheoretica

yieldactual yield ×=


9

Example Suppose that you wish to produce 15.00 grams of barium sulfate by reacting excess sulfuric acid

with barium hydroxide. If this reaction is known to usually give an 87.5% yield, how many grams of barium hydroxide should you use?

Stoichiometry and Net Ionic Equations It is often more practical to use net ionic equations in stoichiometry calculations, especially for

problems dealing with solutions There is no significant difference in solving problems using these equations than with the

methods used so far


10

Example 12.0 grams of strontium chloride is dissolved in approximately 500 mL of water. To this is added

5.00 grams of silver nitrate. Determine a. The total ionic equation and net ionic equation describing this process.

b. The mass of solid precipitate theoretically produced, and c. The moles of each ion remaining in solution at the end of the reaction.

Chemistry 120 Molarity Notes

1

Molarity Concentration and Molarity

The concentration of a solution is a measurement of how much of a solute it contains (the solute is the substance dissolved within it.)

The most useful term for concentration relates the number of moles of solute in a liter of solution; it is called the molarity, and abbreviated “M”:

We often use square brackets around a compound or ion to indicate molarity For example [Na+] means “molarity of sodium ion” For example, to prepare a 1.0 M NaCl solution, we take 1.0 mol of NaCl (58.44 g) and add

enough water to it to bring the total volume of the solution to 1.0 L. We cannot simply say “add 1.0 L of water,” because the presence of the salt will effect the

final volume of the solution. Example Explain how to prepare 2.00 L of a 0.550 M solution of potassium chloride.

Mole-Volume Problems Since the molarity of a solution relates the volume to the number of moles of a solute, we can

use it to help us calculate the amount of product formed in a reaction. First, we must convert the volume of a solution to moles of solute, using the molarity. From there, we can use the mole ratios of the chemical equation to figure out how much

product is produced.

solution Lsolute mol)molarity(M =


2

Examples 25.00 mL of a 0.45 M AgNO3 solution is treated with excess NaCl solution. What solid precipitate

is formed, and what should its mass be in grams?

What volume of 0.224 M Ba(OH)2 solution is required to completely neutralize 15.00 mL of a 0.855 M HNO3 solution?


3

25.00 mL of a 0.425 M potassium iodide solution is added to about 90 mL of water. To this is

added 25.00 mL of a 0.724 M lead(II) nitrate solution. What precipitate is formed, and what is its mass?

Millimoles Since we often carry out volume measurements in the lab in milliliters, it is often convenient to

use millimoles (mmol) rather than moles in calculations A millimole is defined in the same way other metric units are

1 mmol = 10-3 mol, or equivalently 1000 mmol = 1 mol

We can define molarity as and use the form which is most convenient

Example Let’s redo an earlier problem, using millimoles instead of moles in the calculation:

What volume of 0.224 M Ba(OH)2 solution is required to completely neutralize 15.00 mL of a 0.855 M HNO3 solution?

solutionmLsolutemmol

solution Lsolute mol)molarity(M ==


4

Titrations

A titration is usually carried out to determine the molarity of an unknown solution. A specific amount of a compound whose concentration is well known, called the primary

standard is measured out. An indicator, a compound with a different color under different chemical environments, is

added to the primary standard. The unknown solution is added until the indicator changes color, signaling the end of the

titration (called the end point). From the collected data it is possible to determine the molarity of the unknown solution. At the end point, we assume that there are “equivalent” amounts of each chemical present

This is not precisely true, but this definition is adequate for now By equivalent, we mean that the ratio of the moles of each reactant added equals the ratio in

the balanced equation Examples

HCl + NaOH H2O + NaCl Ratio of acid to base: 1 to 1, so we need 1 equivalent of acid for every 1 equivalent of base H2SO4 + 2 NaOH 2 H2O + Na2SO4

Ratio of acid to base: 1 to 2, so we need 1 equivalent of acid for every 2 equivalents of base

Examples 10.00 mL of an HCl solution is pipetted into a beaker containing about 25 mL of deionized water.

This solution is then titrated with 0.4575 M NaOH. It is found that 17.85 mL of base is required to reach the end point. What is the molarity of the acid solution?


5

4.2554 grams of solid oxalic acid (H2C2O4) is dissolved in about 50 mL of deionized water. It is found that 37.55 mL of a KOH solution is required to titrate this acid to the end point. What is the concentration of the base solution?

Ion Concentration For many chemical calculations, it is essential that we know the concentration of each ion in

the solution For example, consider a 1.00 M CuCl2 solution What is the molarity of Cu2+

in this solution?

What is the molarity of Cl-?

Examples 10.00 mL of 2.45 M NaCl solution is combined with 10.00 mL of 1.50 M BaCl2 solution. What is

the concentration of each ion in solution after the two have been combined?


6

How could we prepare 500. mL of a solution with [Ni2+] = 0.100 M? We will use nickel (II) nitrate

as the source of the nickel ion. (Why is it not possible to add only Ni2+ ion to the solution without adding any other ion?)

15.00 mL of a 0.1555 M HCl solution is combined with 10.00 mL of a 0.1250 M Sr(OH)2 solution. What is the concentration of each ion in solution after the reaction has occured?

Chemistry 120 Chapter Twelve Notes

1

Chapter Twelve: Gases Pressure

An important parameter which must be taken into account when studying gases is pressure. Pressure is the amount of force applied to an area:

Consider laying on a bed of nails. Do you want there to be many nails or few nails in the bed?

Pressures of Gases The pressure exerted by a gas in a sealed container results from collisions of many individual

gas particles with the walls of its container. When we say a gas is at high pressure, its individual particles are colliding frequently with

the sides of its container. Later we will see that many other parameters can have an effect on the pressure of a gas

sample, such as the volume of the enclosing container.

Air Pressure The air in the atmosphere exerts a substantial amount of pressure on everything on the

surface of the earth. Since air is composed of matter (nitrogen, oxygen, etc.), it has mass. Air reaches up to at least 20 miles above the surface of Earth. The weight of the air pushes down on everything on the surface, including us. In a similar fashion, the weight of ocean water exerts tremendous pressures at the lower

depths of the oceans.

Measuring Pressure A device used for measuring pressure is called a manometer; a device specifically designed

to measure air pressure is called a barometer. To build a simple barometer, we completely fill a thin cylinder with mercury

Mercury is very dense and is therefore ideal for this application Then, we partially fill a deep dish with mercury and invert the mercury tube in this pool of

mercury. Some of the mercury will descend out of the tube as gravity pulls it down. However, some of the mercury will stay in the tube, as the air pressure keeps this mercury

supported. The weight of the air on the surface of the mercury pool opposes the force of gravity as it

tries to pull the mercury down out of the column. On an “average” day, the height of the mercury column will be close to 760 mm. What would the height of the column be on average at a higher altitude: greater than 760

mm, or less?

The “standard” air pressure is set at this value of 760 mm, as we shall shortly see.

areaforcepressure =


2

Measuring Pressure

Measuring the pressure of the gas in a container involves a similar principal. One method of measuring this pressure involves attaching the sample to a manometer, which

has a U-shaped tube partially filled with mercury. The gas exerts pressure on its side of the U-tube, and the atmosphere exerts pressure on the

other open end. The difference in the heights of the columns can tell use the difference in the pressures of the

gases. The side of the U-tube with the lower height indicates that that side is at a higher pressure,

since it is ultimately pushing back the force of the other gas. To determine the pressure of the gas, the difference in the height is measured and

added to the atmospheric pressure value if the height is lower on the gas sample side. subtracted from the atmospheric pressure value if the height is higher on the gas sample

side.

Units of Pressure

Unfortunately, many different units are used to measure pressure. The units we will employ most often include

Atmospheres (atm) Though not the SI unit of pressure, this is a very common unit. The value of the air

pressure at sea level on an average day is close to 1 atm. Millimeters of Mercury (mm Hg)

This unit is derived from the way one takes a measurement with a barometer or manometer. Recall that on an average day, the height of a mercury column in a barometer is about 760 mm Hg.

1 atm = 760 mm Hg (approximately, though treat this as exact). Torr (torr)

This unit is essentially the same as the mm Hg, so 1 atm = 760 torr (exactly).

You must know these units and conversions. Other units of pressure you may encounter, but need not memorize the conversion for,

include Pascals (Pa) (the SI unit of pressure) and kilopascals (kPa) Pounds per square inch (psi) Bars (bar) and millibars (mbar)

1 atm = 101,325 Pa = 14.7 psi = 1013 mbar


3

Kinetic-Molecular Theory

Recall that our only previous description of gases stated that gases completely fill and take the shape of their containers.

The Kinetic-Molecular Theory goes much farther in describing how gases ideally behave. There are five principle statements that comprise it, which should be understood thoroughly

as well as memorized. Note that some textbooks give slight modifications of these

The Statements of the Kinetic-Molecular Theory The particles of a gas are so small compared to the distances between them that the volume

of the individual gas particles is considered negligible (zero). Gas particles are in constant, rapid, and random motion. Gas particles exert neither attractive nor repulsive forces on each other. The average kinetic energy of a collection of gas particles is directly proportional to the

temperature (in Kelvin) of the gas. From this we can state that all gases at the same temperature have the same average

kinetic energy. Collisions between gas particles are perfectly elastic

By this, we mean that the gas molecules do not lose energy when they collide The Gas Laws

In considering gases, there are three important measured values which we will often consider.

Pressure (P) Volume (V) Temperature (T)

Note that the units of temperature will always be expressed in Kelvin in the gas laws. The gas laws describe the mathematical relationship between these parameters

Boyle’s Law

Suppose we have a gas in a sealed piston (a container whose volume can change), and we know its pressure, volume, and temperature.

Keeping the temperature the same, we decrease the volume of the container. What will happen to the pressure?

Hint: Will the gas molecules collide with the sides of the container more or less frequently?

Pressure and volume are inversely proportional. That is, if one increases, the other must decrease. Remember that the temperature must remain constant for this to be true.

We can use this fact to derive Boyle’s Law: P1V1 = P2V2 (at constant T) where the 1 stands for the beginning values and the 2 the final values.

Derivation:


4

Example A gas occupies 29.5 L, and exerts a pressure of 640. mm Hg. If the container’s volume is

expanded to 40.0 L (keeping the temperature constant), what will be the pressure of the gas in mm Hg and in atm?

Charles’ Law Suppose we have a gas in an expandable container, and, without changing its pressure, we

raise its temperature. How will the behavior of the gas molecules change?

In order to keep the pressure constant, what will happen to the volume of the gas?

Volume and temperature are directly proportional. That is, if one increases the other increases. Pressure must remain constant for the process.

Mathematically, this is stated as (constant P)

Example A gas occupies 35.8 L at 1.13 atm and 42.3 °C. What would be the temperature of this gas, in

Celcius, if the volume were decreased to 25.0 L and the pressure was left unchanged?

2

2

1

1

TV

TV

=


5

Gay-Lussac’s Law

Suppose a gas is kept in a rigid container whose volume cannot change. If the gas is heated, how will the pressure be effected?

Pressure and temperature are directly proportional, assuming the volume remains unchanged. Mathematically, this is stated: (constant V)

The Combined Gas Law The three previous laws can all be combined into one convenient form, the combined gas

law.

You may use this law in place of the other three if you wish, since the constant values will always divide out.

Example 72.6 L of a gas at 35.7 °C and 0.872 atm is heated to 50.0 °C and allowed to expand to 87.5 L.

What is the pressure of the gas under these new conditions?

Avogadro’s Hypothesis Amadeo Avogadro is credited with stating one of the most fundamental statements describing

gases: “Equal amounts (moles) of gases occupy equal volumes under the same conditions”

This greatly simplifies our treatment of gases, as we can conclude that, regardless of what composes the gas, if we know how many moles of it we have and its conditions, we can find its volume.

2

2

1

1

TP

TP

=

2

22

1

11

TVP

TVP

=


6

Avogadro’s Law

Mathematically we can use Avogadro’s hypothesis to derive the mathematical relationship (constant P, T)

n is the number of moles of the gas, regardless of what the gas actually is.

The Ideal Gas Law All the previous laws allow us to compare a gas under one set of conditions with its values

under other conditions. The Ideal Gas Law summarizes the previous laws in a fundamentally different way: its power

is that it shows how the four parameters (P, V, T, and n) can be used to determine one another.

The Ideal Gas Law is, mathematically, PV = nRT where R is a constant, called the universal gas constant.

If you know any three parameters from the ideal gas law, the fourth can be determined by substitution into this equation.

Examples What volume is one mole of a gas expected to occupy at exactly 1 atm and exactly 0 °C?

Note: These pressures and temperatures are called STP (standard temperature and pressure).

If a balloon containing methane has a volume of 7.25 L at 754 mm Hg and 296.4 K, then how many moles of methane does the balloon contain? How many grams is this?

2

2

1

1

nV

nV

=

KmolatmL 0.0821 R⋅

⋅=


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A 1.45 g piece of dry ice (solid CO2) is put in a balloon and warmed to 298 K, causing all the CO2

to sublimate (turn to gas). The volume of the balloon after the sublimation is completed is measured to be 3.29 L. What is the pressure exerted by the gas inside the balloon?

What is the density of pure carbon dioxide gas at 1.05 atm and 300.0 K? Report your answer in grams per liter.

Dalton’s Law of Partial Pressures In a mixture of gases (like air) each substance in the mixture exerts its own indivual pressure,

called the partial pressure of that gas. The total pressure of the mixture of these gases is simply the sum of these partial pressures.

So, air pressure is derived from the partial pressures of all the gases making up the air, including N2, O2, Ar, H2O, CO2, and others.

Mathematically, this is stated as Ptotal = Pgas a + Pgas b + Pgas c + …

There are as many terms on the right side of the equation as there are gases in the mixture.


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Example Argon, neon, and helium gases are all combined in a 40.0 L container. The total pressure of the

three gases is measured to be 1.45 atm. Argon and neon each exert a partial pressure of 0.65 atm. What is the partial pressure of helium in this mixture?

Collecting a Gas Over Water One common way to collect a gas produced by an experiment is over water. From the reaction vessel, the gas enters a hose which leads to a container filled with water. The gas then bubbles up through the water into a tube filled with water that penetrates the

waters surface. The gas displaces the water in the tube, and is collected there. However the gas is “wet”, meaning it contains some water vapor. To find the pressure of the “dry” (i.e. without water) sample, we must subtract out the

pressure of the water vapor mixed with it. The pressure exerted by the water pressure can be found in a table, so long as we know

the temperature of the water. If the level of the water in the tube is the same as the level of the water in the container the

tube is in, the following relationship exists: Pair = Pgas + Pwater vapor

Example A sample of argon is collected over water at 23 °C. After collecting the gas, the height of the gas

in the collecting tube is lowered to that of the water container. According to a barometer, the air pressure is 763 mm Hg. What is the pressure exerted by argon in atmospheres?


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Gas Stoichiometry

The Ideal Gas Law aids us in predicting the quantity relationships when reactants or products are gases.

An important value to keep in mind is that 1 mole of a gas at STP will occupy 22.4 liters. We will use the Ideal Gas Law for similar calculations under different conditions.

Examples 54.8 L of hydrogen gas is combined with excess oxygen at STP. What mass of water should be produced? 72.9 L of hydrogen (at STP) is reacted with 28.0 L of nitrogen (at STP), producing ammonia. What volume will the ammonia produced occupy at 1.25 atm and 325 K?


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65.0 grams of nitrogen gas and 15.0 grams of hydrogen gas are combined in a 10.0 L container. The gases react, producing ammonia gas. Assume that the reaction goes to completion. a. What is the partial pressure of each gas in the container at the end of the reaction, assuming the total pressure in the container is 2.50 atm? b. What is the temperature of the gas sample in the flask under the conditions in part a? c. Suppose that all the ammonia produced in this reaction is dissolved into water to make 555 mL of solution. How many milliliters of 3.00 M H2SO4 solution would be required to completely neutralize the ammonia solution?

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