Chemistry Semester 2 Final Exam Study Guide

Learning goals Checklist

As you prepare for your final exam use the following list to monitor your progress. Only check off a learning

goal when you are confident of your ability to complete the task described.

Unit 0: Being a Chemist

Recognize and apply safe practices when working in the lab.

Identify the appropriate instruments to use when working in the lab

List the types of investigations typically carried out by scientist

Identify the independent, dependent and controlled variables in an experiment

Explain the importance of controlled variables in scientific investigations

Properly record measurements to the correct number of significant figures for the instrument used

Keep and organize data into data tables

Analyze date to determine which type of graph is best

Analyze data and use it to write conclusions that address the purpose, hypothesis, evidence and an

explanation of why the evidence is important.

Unit 1: Matter

Describe density as a physical property that depends only on the type of substance, not the amount of

substance (intensive property)

Distinguish between elements, compounds, and mixtures (heterogeneous and homogeneous)

Explain why elements and compounds are pure substances but mixtures are not

Sort particulate models of elements, compounds, and mixtures into their respective categories

Calculate the mass, volume, and density of an object from real world data, such as: a plot of mass versus

volume to calculate the density of a substance

Predict whether an object floats or sinks in a liquid

Plan and conduct an experiment to classify properties of matter as intensive

Demonstrate how intensive properties can be used to identify a compound

Classify changes in matter as either the result of a chemical change or a physical change

Provide evidence or clues that a chemical change has taken place

Unit 2: Atomic Structure

Communicate information from historical experiments

Write isotope symbols and isotope notation.

Read an isotope symbol and use it to determine information about the isotope such as the element, mass

number, atomic number, number of protons, neutrons and electrons in a neutral atom.

Read a periodic table to determine information about elements such as; chemical symbol, atomic

mass/mass number, atomic number

Use the periodic table to aid in determining the number of protons, neutrons and electrons in a neutral

atom and isotope.

Develop and use models of atomic nuclei

Calculate average atomic mass?

Write the long hand electron configuration of elements

Use the electron configuration to determine the number of core electrons and the number of valence

electrons

Unit 3: Periodic Properties

Explain the development of the Periodic Table

o Mendeleev (using atomic weights)

o Moseley (using atomic numbers)

o Periodic Law: the chemical and physical properties of elements are periodic functions (recurrent

functions) of their atomic numbers

Read and explain the Periodic Table:

o Names and Locations of:

- Groups (vertical – up and down)

- Periods (horizontal – across left to right)

- Metals and their general properties: malleable, ductile, shiny, good conductors, lose electrons to

form positive ions

- Nonmetals and their general properties: Brittle, no metallic luster, poor conductors, gain electrons to

form negative ions

- Metalloids: elements sharing some metallic and some nonmetallic properties

- “Stairway to 7”: boundary line between metals and nonmetals

- Main group elements: groups 1-2 and 13-18 (The “A” groups)

- Transition metals: groups 3-12; scandium (Sc) through zinc (Zn) and all below them

- Lanthanides: elements 58-71, following lanthanum (La)

- Actinides: elements 90-103, following actinium (Ac)

- Alkali metals (group 1): Li, Na, K, Rb, Cs, Fr. Hydrogen (H) is chemically similar.

- Alkaline earth metals (group 2): Be, Mg, Ca, Sr, Ba, Ra

- Halogens (group 17): F, Cl, B, I, At

- Noble or inert gases (group 18): He, Ne, Ar, Kr, Xe, Rn, with complete octets

o s,p,d,f blocks

- Write the long hand and short hand electron configuration of elements AND Ions including those

with exceptions to the rules

o valence electrons

- definition: outermost shell of electrons that determine chemical properties of an element

- be able to determine # of valence electrons for any main group element (A group)

- Draw Lewis dot diagrams

- Importance of valence electrons to physical and chemical properties

- The octet rule: atoms are most stable with 8 valence electrons (2 in the first period)

o Ionization

- Reason why atoms form ions

- Predict what ions will form for an atom based on atom’s position on periodic table

Understand Periodic Trends and use these trends to predict properties of atoms

o Recognize that elements within the same group share similar physical and chemical properties

o Understand the reason why this occurs

o General Period Trends

- Atomic radius

Definition

Know trend (both group and period)

Reason for trend (both group and period) (Coulombic Attraction/ Effective Nuclear Charge)

- Ionization energy

Definition

Basic chemical equation

Know trend (both group and period)

Reason for trend (both group and period) (Coulombic Attraction/ Effective Nuclear Charge)

- Electronegativity

Definition

Know trend (both group and period)

Reason for trend (both group and period) (Coulombic Attraction/ Effective Nuclear Charge)

- Ion Size

Definition

Be able to compare ion size to parent atom

Be able to explain why there is a difference in size between an ion and its parent atom

(Coulombic Attraction/ Effective Nuclear Charge and Octet Rule)

Unit Summaries

The following is a brief overview of what was learned during the unit. Use these as a beginning point on what to study.

Unit 0: Being a Chemist

Unit Essential Questions

1. What is the nature of science?

a. Which questions can be answered by science?

b. How does scientific observation differ from regular every day observation?

c. How do scientist observe things they cannot see?

2. How do we keep safe when working in the lab?

3. How are scientific investigations designed and conducted?

a. What factors do we need to consider when designing investigations in order for the results to be

considered "valid"?

b. Why do scientific measurements sometimes have decimals and other times don't?

c. How do scientist organize data?

i. How do I read a data table?

ii. How do I determine which type of graph to use?

d. How are conclusions written?

Lesson Overviews

Scientific Investigations Use a Variety of Methods o Scientist use diverse methods and do not always use the same set of procedures to obtain data. o New technologies advance scientific knowledge. o Scientific inquiry is characterized by a common set of values that include: logical thinking,

precision, open-mindedness, objectivity, skepticism, replicability of results, and honest and

ethical reporting of findings.

Scientific investigations use a variety of methods, tools, and techniques to revise and

produce new knowledge.

Science Ideas Are Based On Experimental Evidence. o All science disciplines share a common set of rules for how to evaluate evidence and its’ use in

constructing explanations about natural systems. o Science includes the process of evaluating the data gathered for patterns and determining how

and if they fit into the current theory o Scientific arguments are made stronger when there are multiple pieces of evidence used to

support a single explanation or theory.

Scientific Ideas Can Be Changed And/Or Revised When New Evidence Emerges.

o Scientific explanations can be based on probability (the likelihood of an explanation being true).

o Most scientific knowledge is quite durable but is, in principle, subject to change based on new

evidence and/or reinterpretation of existing evidence.

o Scientific argumentation is a kind of a conversation based on logic, used to clarify the strength of

relationships between ideas and evidence that may result in revision of an explanation.

Science Models, Laws, Mechanisms, and Theories Explain Natural Phenomena

o Theories and laws provide explanations in science, but theories do not with time become laws or

facts.

o A scientific theory is a corroborated explanation of some aspect of the natural world, based on a

body of facts that has been repeatedly confirmed through observation and experiment,

The science community validates each theory before it is accepted.

If new evidence is discovered that the theory does not accommodate, the theory is

generally modified in light of this new evidence.

o Models, mechanisms, and explanations collectively serve as tools in the development of a

scientific theory.

o Laws are statements or descriptions of the relationships among observable phenomena.

o Scientists often use hypotheses to develop and test theories and explanations.

Science is a Way of Knowing

o Science is both a body of knowledge that represents a current understanding of natural systems

and the processes used to refine, elaborate, revise, and extend this knowledge.

o Science is a unique way of knowing and there are other ways of knowing.

Other ways include those having to do with religious belief systems, philosophical

systems, ethics and law systems, etc

HYPOTHESIS

THEORY

LAW

o Science distinguishes itself from other ways of knowing through use of empirical standards,

logical arguments, and skeptical review.

o Science knowledge has a history that includes the refinement of, and changes to, theories, ideas,

and beliefs over time.

Scientific Knowledge Assumes an Order and Consistency in Natural Systems

o Scientific knowledge is based on the assumption that natural laws operate today as they did in

the past and they will continue to do so in the future.

o Science assumes the universe is a large single system in which basic laws are consistent.

Science is a Human Endeavor

o Scientific knowledge is a result of human effort, imagination, and creativity.

o Individuals and teams from many nations and cultures have contributed to science and to

advances in engineering.

o A scientists’ background, their religious and cultural upbringing, and their chosen field of study

can influence the nature of their findings.

o Technological advances have influenced the progress of science and science has influenced

advances in technology.

o Science and engineering are influenced by society and society is influenced by science and

engineering.

Science Addresses Questions About the Natural and Material World

o Not all questions can be answered by science.

o Science and technology may raise ethical issues for which science, by itself, does not provide

answers and solutions.

o Science knowledge indicates what can happen in natural systems—not what should happen. The

latter involves ethics, values, and human decisions about the use of knowledge.

o Many decisions are not made using science alone, but rely on social and cultural contexts to

resolve issues

The 6 basic elements of the Scientific Method

1. Observation

2. Question or Problem

- This is written in the form of a question.

Example: Which of these two paper towel brands will absorb more water?

3. Form a Hypothesis

- A hypothesis is a possible answer to your question. It is usually based on the

inference made from the observation!

- It is a testable statement

Example: If brand A is thicker than brand B, then brand A will absorb more water

because it will have more air pockets to trap the water in.

4. Perform an Experiment to test the hypothesis.

Example: How would you test which paper towel will absorb more?

Spread out a certain amount of liquid

Use the paper towel to pick up as much as you can

Measure how much is left over

5. Analyze the Data

- Includes creating tables or graphs to look for trends.

6. Draw a Conclusion

- Here you look at the data from the experiment to see if your hypothesis was correct or

incorrect.

Experimental Design

o Variables:

o Root: Vary

o Vary means to change

o Variable: factors that change.

o Independent Variable: the variable that you purposely change and test.

o Dependent Variable: the factor that may change as a result of your independent

variable.

o Constants (or controlled variables): keeping all conditions of the experiment the same

except for the variables you are testing.

Example: You want to answer the questions: Will houseplants grow faster if

you make the room warmer? To answer this question, you decide to grow

the plants at different temperatures. o Independent variable: the temperature of the room

o Dependent variable: how fast plants grow

o Constants (controlled variables): size of container, type of soil, amount of

water, amount of light exposure, use of fertilizer. o An Hypothesis

Is a prediction about the outcome of a scientific investigation. They are based on a

person’s observations and previous knowledge or experience

Written in an If…then…because statement form

this shows a cause and effect relationship

Example: If I wash my dog, then he will be clean because the water and

shampoo will remove all of the dirt.

It is testable

Examples

1. If I give my plants fertilizer, then they will grow as big as my neighbor’s

plants, because that is what my neighbor does.

Testable and properly worded

2. If I get lucky, then my plants will grow bigger because that is how luck

works.

not testable, can’t test “getting lucky” 3. My plants aren’t growing bigger because I don’t water them enough.

Not worded properly.

Data and Data Analysis

o Measurement is important to science because measurement gives meaning to observations.

Consider the statement "That reaction produced a lot of energy." To a biochemist it may mean a

ten degree rise in temperature, to a nuclear physicist it may mean the release of enough energy to

obliterate a mountain!

Quantitative measurements (think of the word "quantity") result in a definite form,

usually involving numbers.

Qualitative measurements are descriptive and are nonnumeric.

Accuracy refers to how close a quantitative measurement is to the accepted value and

precision is how repeatable the results of a quantitative measurement are.

Lets say that I know an object has a mass of 5.00 grams. Two sets of five students

measure the mass of the object giving the following results:

Set 1 Set 2

student 1 : 5.00 g student 1 : 6.50 g

student 2 : 3.50 g student 2 : 6.49 g

student 3 : 8.20 g student 3 : 6.51 g

student 4 : 6.50 g student 4 : 6.50 g

student 5 : 5.78 g student 5 : 6.50 g

- Only student 1 in data set 1 is accurate and the set of data shows little precision. In data set 2 no

one is accurate (remember the accepted value is 5.00g), but the data set shows good precision.

Error is a measure of the accuracy of a measurement. Absolute error (Ea) is the absolute (always

positive) difference between the accepted value and the observed (measured) value.

Absolute error = Observed value ­ Accepted value

Percent error is the absolute error expressed as a percentage of the accepted value.

Percent error = Absolute error / Accepted value x 100%

Percent error is a better expression of the accuracy of a measurement. An absolute error of 0.005 grams may

sound great, but if the accepted value is 0.0001 g, the % Error is 5,000 %!

Significant figures are all those digits in a measurement which can be known precisely and a last digit

which can be estimated. Significant figures are those digits in a measurement which are actually

measured and are an indication of the precision of the instrument used to make a measurement. A

measurement of 1 meter has only one significant figure and has little precision. A measurement of 1.000

meters has four significant figures and has a much greater degree of precision. Consider the

measurement at the arrow on the following centimeter ruler.

A measurement made at the arrow would be 1.94 cm. The digits 1 and 9 are known because we can see

the lines which indicate these numbers. The last digit, 4, is estimated. It would not be correct to round the

measurement up to 2.0 or down to 1.9, because we can see that it is somewhere in between. The ruler has

a degree of precision which is to the 0.01 (hundredths) cm place. What should the measurement indicated

on the following ruler be?

The measurement should be 3.00 cm. This ruler has lines at the tenths of a cm place and therefore the

measurement can be carried one place further ­ the last estimated digit. Contrast this to the ruler below.

A measurement made at the arrow would be 1.5 cm. This ruler has a lesser degree of precision.

Significant figures are used in measurements ­where instruments with certain degrees of precision are used

and where measurements are made. Significant figures are not used in defined quantities. For example, by

definition 1 meter equals 1 000 millimeters. I do not need to write 1.000000.... m, because a meter is always

exactly a meter, it has an unlimited number of significant figures.

The following are the rules for determining which digits in a number are significant and which are not.

1. All nonzero digits are significant.

2. Zeroes appearing between nonzero digits are significant.

3. Zeroes appearing in front of all nonzero digits are not significant (the zeroes in the number 0.0056 are

not significant).

4. Zeroes to the right of a nonzero digit and to the right of a decimal place are significant. (2.000 has 4 sig

figs, while 0.0200 only has three sig figs).

5. Zeroes to the right of nonzero digits and to the left of a decimal: 500 = 1 sig fig 500. = 3 sig figs

500.0 = 4 sig figs

Rules for determining how many numbers to report in a calculation

In addition and subtraction the sum or the difference can have no more decimal places than the number in

the calculation which has the least number of decimal places.

For example, if I subtract 1.2 from 3.40, the final answer is 2.2, not 2.20. If I add 2.20 and 3.1, the

final answer is 5.3

In multiplication and division the answer can have no more significant figures than the number in the

calculation with the least number of significant figures.

For example if I multiply 2.6 and 4.555, the final answer can only have two significant figures.

* Note the difference between addition­subtraction and multiplication­division.

In addition or subtraction it is the number of decimal places that matter, in multiplication and division

it is the total number of significant figures that matter.

The reason for these rules is that measurements must not be presented to a greater degree of precision

than is actually known. For example if two measurements are to be added, 3.2 cm and 4.0001 cm the sum

would be 7.1001 cm. However the first measurement (3.2 cm) is not very precise, it is known only to the tenths

of a cm place. The hundredths, thousandths and ten thousandths place are not known and should therefore not

be reported in the sum.

The system of units used for measuring in science is the SI (after the French "Le Systèmes International

d'Unités). It is commonly referred to as the metric system, but is actually a revised version of the metric

system. The advantages of this system is that it is used worldwide, making it easier for scientists to compare

and share data, and that it is based on the decimal system, making it easy to convert from one unit to another.

o The basic unit of length is the meter.

o The basic unit of volume is the cubic meter (m 3 ).

o The basic metric unit of mass is the kilogram.

Volume is the amount of space an object occupies.

Mass is the amount of matter that an object has.

Weight is the pull of gravity on an object.

Metric prefix Symbol Meaning kilo­ k 1000 deci­ d 0.1 (one­tenth) centi­ c 0.01 (one­hundredth) milli­ m 0.001 (one­thousandth)

micro­ µ 10­ 6 (one­millionth) nano­ n 10­9 (one­billionth)

A liter is the most commonly used metric unit of liquid volume and is equal to 1000 cm 3 (a 10 cm x 10cm x

10 cm cube).

A kilogram is the basic metric unit of mass and it is equal to 1000 grams (a gram is so small that it is too

inconvenient to use as the basic metric unit of mass).

Density is the ratio of the mass of an object to its volume. Density is calculated by dividing the mass of an

object by its volume. For an example, if a 10.0 kg object occupies a volume of 2.0L its density is 10.0kg/2.0L

which is 5.0 kg/L. Note the units which indicate the ratio of mass to volume­ kilograms per liter.

Temperature is a measurement of the average kinetic energy of the particles of a substance. Alcohol or

mercury filled thermometer ­ the liquid is contained in an evacuated glass rod. The liquid expands and

contracts to indicate the temperature on a scale.

The Celsius scale of temperature has the freezing point of water at 0° C and the boiling point of water

(at 1 atmosphere of pressure) at 100° C.

The Kelvin scale is based on absolute zero, the theoretically lowest possible temperature where all

molecular motion would cease.

The Kelvin scale is 273 degrees higher than the Celsius scale. The following equations can be used to

convert between these scales.

K = ºC + 273 and ºC = K­273

Unit 1: Matter

Unit Essential Questions

Why can intensive properties be used to identify substances?

a. How do you know if something is matter?

i. Why do we place matter into different categories?

b. How can matter change?

i. How do physical changes differ from chemical changes?

c. What information can we gather about all forms of matter?

i. How are physical properties different from chemical properties?

ii. It's called chemistry so why are physical properties studied anyway?

Lesson Overviews

Chemistry is the study of matter and the changes that matter undergoes.

o Pure chemistry is chemistry done for the sake of gaining new information.

An example would be any type of research, such as an analytical chemist determining

the composition of an unknown compound, or a biochemist determining the chemical

structure of a newly discovered virus.

o Applied chemistry makes practical use of existing chemical knowledge. An example would be manufacturing rayon or nylon(synthetic fibers) and producing

clothes. Vocabulary

o Matter anything which takes up space and has mass

o Mass a measure of the amount of matter an object has

o Substance a particular type of matter which has a uniform and definite composition (e.g. sodium

(Na), water (H2O).

o Intensive Properties are those that

o Extensive Properties are those properties that a determined by what a substance is made of. All

chemical properties are extensive properties

o Physical properties are those properties which can be observed without the substance changing

into a new substance. (brittleness of chalk, color of the sky).

o Chemical properties are those properties which cannot be observed without the substance

changing into a new substance. (Wood burning)

o Four states of matter :

Solid has a definite shape and volume

Liquid has a definite volume, but not a definite shape (takes the shape of its container)

Gas has neither a definite shape or volume

Plasma highly energized state of matter in which the electrons have separated from the

nuclei creating a mixture of electrons and positive ions (the sun is made of plasma, only

here the matter is so ionized it is almost pure nuclei one tablespoon has an estimated mass

of about 2 tons).

o Gases are substances which are in the gaseous state under normal environmental conditions.

o Vapors are the gaseous states of substances which are normally solids or liquids.

o Physical changes are changes in

the physical properties of a

substance (do not result in a new

substance)

ex. Grinding a crystal

(changing size).

o Chemical changes are changes

which result in a new substance.

For example, burning wood

changes it into mostly water

and carbon dioxide.

A mixture is a physical combination of

two or more substances.

o Homogeneous mixtures have a

uniform composition (salt water),

o Heterogeneous mixtures are not uniform in composition (a bucket of dirt).

o A solution is a homogeneous mixture where one substance (the solute) is completely and evenly

dispersed in another substance (the solvent).

Salt dissolving in water is a physical change the salt is still salt and the water is still

water.

o A phase is any part of a system with uniform composition and properties.

A homogeneous mixture, such as a solution is made up of only one phase.

A heterogeneous mixture consists of more than one phase.

o Law of Conservation of Energy: energy can neither be created nor destroyed. The amount of

energy in the universe is constant.

o The Law of Conservation of Mass: in any physical or chemical reaction mass is conserved –

mass is neither gained nor lost.

o A chemical reaction is a reaction in which a chemical change takes place one or more

substances become new substances.

o A physical reaction results in the changes of the physical properties of one or more substances,

but no new substances are produced.

o Reactants are the substances that are changed in a chemical reaction and the products are the

substances produced in a chemical reaction.

For example, in the summary chemical reaction of photosynthesis, water and carbon

dioxide are the reactants and glucose and oxygen are the products.

6 CO2 + 6 H2O C6H12O6 + 6 O2

reactants products

Factors which indicate a chemical change include:

o A change in color or odor

o A change in temperature - energy is absorbed or released

o A gas is formed – formation of bubbles

o A solid is produced from a mixture of two liquids (a precipitate))

A precipitate is an insoluble solid produced by a chemical reaction.

Unit 2: Atomic Structure

Unit Essential Questions

1. Why has our concept of the atom changed over time?

a. How did the results of historical experiment regarding atomic theory contribute to the current model

of the atom?

b. How do we determine the structure and function of things we cannot see?

2. How can nuclei of the same element differ?

3. How does the abundance of various isotopes effect an elements atomic mass?

4. How can we model the arrangement and location of electrons in the atom?

Lesson Overviews

Atomic Structure History

An atom is the smallest part of an element that retains the properties of that element. The concept of an

atom goes a long way back. It was first suggested by an ancient Greek named Democritus (the Greek

word "atomos" means indivisible). Democritus theorized that if you took an object and cut it in half

again and again you would eventually end up with some particle which could not be further divided.

In the early 1800's an English scientist by the name of John Dalton started relating what chemists could

see to the concept of the atom. He came up with an atomic theory which could be stated as follows:

o All elements are composed of tiny indivisible particles called atoms.

o Atoms of the same element are identical and differ from atoms of other elements.

o Atoms of different elements can combine together in simple whole number ratios to form

compounds.

o Chemical reactions are the rearranging of the combinations of atoms of elements in

compounds. The atoms themselves remain unchanged.

Atoms are composed of three primary particles, protons, neutrons and electrons.

Particle Symbol Location

Relative

electrical charge

Approximate

relative mass in

amu*

Actual mass

in grams

Electron e- orbits the

nucleus 1- 1/1840 9.11 x 10-28

Proton p+ nucleus 1+ 1 1.66 x 10-24

Neutron nº nucleus 0 1 1.66 x 10-24

*amu (atomic mass unit = 1.66 x 10-24 grams)

The electron was discovered by the English physicist Sir Joseph J. Thomson around 1897 with the use

of a cathode ray tube. A cathode ray tube is similar to your TV. It has an anode (negative electrode) and

a cathode (positive electrode). These are enclosed in an evacuated (air removed) glass container and

when a charge is applied the electrons flow from anode to cathode through the open space of the glass

container. Thomson observed these particles and determined that the particles:

• move at a very high speed

• have a negative charge

• have a mass of about 1/2000 of a hydrogen atom (smallest atom)

• were the same regardless of which gas was used in the container or the metal used as the

electrode

The particles were eventually named "electrons." His model of the

atom was the “plum-pudding” model. This discovery shattered

Dalton's notion of an atom. To Dalton atoms were tiny, solid

particles, not containing smaller particles.

The nucleus was discovered by Ernest Rutherford in 1911.

Rutherford set up an apparatus in which he aimed alpha particles

(type of radiation made up of helium nuclei) at a very thin sheet of gold foil. If the atoms were made up of

evenly dispersed protons and electrons, as believed at the time, the alpha particles should go straight through

unhindered. What happened was that some of the alpha particles went through; some were deflected as they

passed through and some bounced back. Rutherford concluded that the positive mass of the atom (protons) must

be concentrated in a very small area and most of the rest of the atom must be empty space. The area where the

protons were located was called the nucleus.

The nucleus of an atom is extremely small as compared

to the entire size of the atom. It can be compared to the

size of a marble in a football stadium where the marble

represents the nucleus and the stadium the entire atom.

Most of the space of the atom is occupied by the orbiting

electrons.

Structure of the Nucleus:

o Protons

- Positive charge, mass of 1.673x10-27 kg

- The number of protons in the nucleus determines the atom's identity and is called the atomic number

o Neutrons

- No charge, mass of 1.675x10-27 kg

The different historical models are described as follows:

o Dalton's model of the atom - solid, tiny, indivisible particles.

o Thomson's model - often describe as the "plum pudding" model - electrons are scattered

throughout the atom.

o Rutherford's model - includes the solid nucleus in the center of the atom.

Sir James Chadwick discovered the neutron in 1932.

Millikan's Oil Drop Experiment:

Electron

o Electron is negatively charged

o Found outside the nucleus in the “electron cloud” in

orbitals

o Mass is about 1/2000th of a hydrogen atom

▪ Electron mass is 9.109 x 10-31 kg

Electron Orbital “Shapes”

S-orbitals P-orbitals D-orbitals

Atomic number - indicates the number of protons and defines the

element (atomic number 6 is always carbon, atomic number 7 is always

nitrogen etc.).

Mass number - equals the total number of neutrons and protons in the

nucleus of an atom for the most common isotope this equals the atomic

mass rounded off to the nearest whole number

Atomic mass - the average mass of an atom of an element (in amu)

Calculation of the number of particles in an atom of an element:

- number of protons equals the atomic number

- number of neutrons equals the mass number minus the atomic number (remember

virtually all the mass is from the neutrons and protons in the nucleus-each with an amu of

1)

- number of electrons equals the number of protons in a neutral atom

Consider the following periodic table information for carbon, nitrogen and sodium:

Carbon's atomic number is 6, has an average mass of 12.011 amu and carbon's most common isotope has a

mass number of 12 amu. Therefore, the most common type of carbon atom has 6 protons, 6 neutrons and 6

electrons.

Nitrogen's atomic number is 7, has an average mass of 14.007 amu and nitrogen's most common isotope has a

mass of 14 amu. Therefore the most common type of nitrogen atom has 7 protons, 7 neutrons and 7 electrons.

Sodium's atomic number is 11, has an average mass of 22.990 amu and nitrogen's most common isotope has a

mass of 23 amu. Therefore the most common type of sodium atom has 11 protons, 12 neutrons and 11 electrons.

An isotope is a particular form of an atom of an element.

o Isotopes have a different number of neutrons and therefore differ in mass (that is why there is a non-

whole number atomic mass which is an average of the various isotopes).

o The number of protons is always the same in isotopes since they are different forms of the same

element (must be same atomic number).

For example, the following chart shows three isotopes of hydrogen :

Isotope Atomic

Number

Number of

protons

Number of

Neutrons

Number of

electrons

mass

(amu)

Hydrogen-1 1 1 0 1 1

Hydrogen-2

(deuterium) 1 1 1 1 2

Hydrogen-3

(tritium)

1 1 2 1 3

Note hydrogen has three isotopes, each with a whole number mass, yet hydrogen as an element has a

characteristic average mass called the atomic mass (1.0079 amu).

An atomic mass unit (amu) is 1/12 the mass of a carbon-12 atom, or 1.66 x 10-24 grams

Calculating the Average Atomic Mass

Electron Configurations

o Electron configuration: the complete description

of the orbitals occupied by all the electrons in an

atom or an ion

o Aufbau “building up principle” For an atom in its

ground state, electrons are found in the energy

shells, subshells, and orbitals that produce the

lowest possible energy for the atom.

o Hund’s rule states that electrons pair in an orbital

only after each orbital in a subshell is occupied by a

single electron

o Valence electrons are electrons in an atoms outermost shell; electrons in the inner shells are

called core electrons.

o Groups 1A and 2A are s-block elements and groups 3A-8A are p-block; transition metals are

d-block elements

o Pauli Exclusion Principle states that only two electrons can fit into any given orbital and that these

two electrons will have opposite spins

o Orbital Notation

Orbital notation is basically just another way of expressing the electron configuration of an atom. It is very

useful in determining quantum numbers as well as electron pairing. The orbital notation for sulfur would be

represented as follows:

Unit 3: Periodic Properties

Unit Essential Questions

1. How does the number and location of electrons in an atom effect the chemical and physical properties of the element?

2. How can we use the periodic table to make predictions about the properties of undiscovered elements?

Lesson Overviews

o Mendeleev's Periodic Table (1869)

o Organization

o Vertical columns in atomic weight order

a. Mendeleev placed elements in rows with similar properties

o Horizontal rows have similar chemical properties

o Missing Elements

o Gaps existed in Mendeleev’s table

a. Mendeleev predicted the properties of the “yet to be discovered” elements

- scandium, germanium and gallium agreed with his predictions

o Unanswered Questions

o Why didn't some elements fit in order of increasing atomic mass?

o Why did elements exhibit periodic behavior?

o Moseley and the Modern Periodic Table (1911)

o Protons and Atomic Number - The periodic table was found to be in atomic number order, not atomic

mass order

o The Periodic Law

- The physical and chemical properties of the elements are periodic functions of their atomic numbers

-Elements with similar properties are found at regular intervals within the periodic table

o Organization of the Table

o Groups or Families

Vertical columns containing elements with similar chemical properties

o Periods (series)

Horizontal rows of elements

o Metals and Nonmetals

A stair-step line on the table separates the metals from the nonmetals

o Metalloids (Semimetals) straddle the line and have properties of both metals and nonmetals

o Metals and man-made metal elements

Group 1 – Alkali metals (the most reactive metal elements) (except hydrogen (H) also in

this group)

Group 2 – Alkaline earth metals (very reactive metal elements)

Group 17 – Halogens (the most reactive nonmetal elements)

Group 18 – Noble gases (the least reactive elements – inert and very stable)

Lanthanide and Actinide Series (Inner Transition Metals)

o Types of Elements

o Metals

o Luster

o Good conductors of heat and electricity

o Malleable

o Ductile

o High tensile strength

o Nonmetals

o Many nonmetals are gases at room temperature

o Solid nonmetals tend to be brittle and non-lustrous

o Poor conductors of heat and electricity

o Metalloids

o Some properties of metals and some properties of nonmetals

o Solids at room temperature

o Semiconductors of electricity

o Noble Gases

o All are gaseous members of group 18

o Generally unreactive and stable

The periodic law states that when the elements are arranged according to increasing atomic number there is a

periodic pattern in their physical and chemical properties.

Although the periodic table is arranged by increasing atomic number (which indicates the number of protons),

the electrons configuration is what really determines the physical and chemical properties of the elements. The

periodic table can be divided into four groups based on electron configuration:

The Noble gases (Group 0) - have their outermost s and p orbitals filled which creates a stable and non-

reactive (inert) element.

The representative elements - Group A elements - have their s and p orbitals being filled. These

include :

o Group 1A - Li, Na, K etc. - all very reactive with one electron in the outer s orbital

o Group 2A - Be, Mg, Ca etc. - all quite reactive with 2 electrons filling their outer s orbital

o Group 3A - Aluminum group - 3 electrons in outer energy level (2s and 1p) properties vary from

metallic to metalloid

o Group 4A - Carbon group - 4 electrons in outer energy level (2s and 2p) - properties vary from

nonmetallic to metalloid to metallic down the group

o Group 5A - Nitrogen group - 5 electrons in outer energy level (2s and 3p) - properties vary from

nonmetallic to metalloid to metallic

o Group 6A - Oxygen group - 6 electrons in outer energy level (2s and 4p) - properties vary from

nonmetallic to metalloid

o Group 7A - Halogens - all have 7 electrons in the outer energy level (2s and 5p) - properties

vary from nonmetallic to metalloid. Very reactive due to the outer energy level being almost

filled.

The transition metals - elements whose d orbitals are being filled - found in the "d-block." These are

also called the Group B elements

The Inner transition metals - These are the Lanthanide and Actinide series, element whose f orbitals

are being filled.

The s, p, d, and f groups can be identified on the diagram below. The f block (inner transition metals) is

usually shown separated and below the rest of the table.

Periodicity is the property of having periodic properties. The periodic table shows periodicity in the following

properties:

Atomic size - atoms of elements tend to increase as you go down a group (due to a greater number of

energy levels) and atoms of elements tend to decrease in size across a period (greater positive nuclear

charge which draws in electrons -energy levels are constant across a period).

Ionization energy - the energy required to remove an electron from the gaseous state of the atom.

Ionization energy decreases as you go down a group due to the outer electrons being further from the

positive charge of the nucleus and being shielded from the nucleus' positive charge by the inner energy

levels. Ionization energy increases as you move across a period. This is due to the increase in nuclear

charge without the increase in number of energy levels.

Electron affinity - this is the energy change associated with the addition of an electron to a gaseous

atom. This trend is not as consistent as the others, but in general electron affinity decreases down a

group and increases across a period.

Ionic size - The size of ions increases as you go down a group, and decreases as you move across a

period for the metals and for the nonmetals, for the same reasons as atomic size. However, the metallic

ions are positive (have lost electrons, which make up the space or size of the atom) and are much smaller

than the negative nonmetallic ions (have gained electrons which create the volume of atoms or ions).

Electronegativity -the tendency of an atom to gain an electron(s) when combining with another

element.

Electronegativity decreases down a group (due to shielding) and increases as you move across a period (due to

the increase in nuclear charge).

Ionic Radii

o Cations

1. Positive ions

2. Smaller than the corresponding atom (they have lost electrons)

a. Protons outnumber electrons (greater effective nuclear charge)

b. Less shielding of electrons

o Anions

1. Negative ions

2. Larger than the corresponding atoms (they have gained electrons)

a. Electrons outnumber protons (weaker effective nuclear charge)

b. Greater electron-electron repulsion

Valence Electrons

o Valence Electrons

The electrons available to be lost, gained, or shared in the formation of

chemical compounds

Main group element valence electrons are in the outermost s and p

sublevels

The three major forces affecting periodicity are:

Nuclear charge - the greater the number of protons in the nucleus, the greater the positive charge and

the stronger the electrons are held.

Shielding - the effect of inner energy levels reducing the strength of the nuclear charge on the electrons

in the outer energy levels.

Electron configuration - atoms are most stable when their outer orbitals are filled (especially the s and

p orbitals). This causes the Noble gases to be inert.

Properties of the representative groups of elements:

Noble gases - Group 8A, helium, neon, argon, krypton, xenon and radon

o inert (unreactive) because of stable electron configuration (filled s and p orbitals)

Alkali metals - Group 1A, lithium, sodium, potassium, rubidium, cesium and francium

o very reactive (one electron away from a filled s and p orbital)

o low density

o low melting point

o good electrical conductivity

o react with water to form strong bases (sodium hydroxide, lithium hydroxide etc.)

Alkaline earth elements - Group 2A, beryllium, magnesium, calcium, strontium, barium and radium

o very reactive (2 electrons away from a filled s and p orbital)

o react with water to form hydroxides

o used to form metal alloys

Aluminum group - Group 3A - 3 electrons in outer energy level (2s and 1p) properties vary from

metallic to metalloid

o aluminum is the most useful metal of this group being lightweight and strong to make boats,

aircraft etc.

Group 4A - Carbon group - 4 electrons in outer energy level (2s and 2p) - properties vary from

nonmetallic to metalloid to metallic down the group

o diamond and graphite are forms of pure carbon

o silicon and germanium are semiconductors used in electronics

o tin and lead are useful metals

Group 5A - Nitrogen group - 5 electrons in outer energy level (2s and 3p) - properties vary from

nonmetallic to metalloid to metallic

o nitrogen and phosphorus are elements necessary to form proteins and nucleic acids in living

things

Group 6A - Oxygen group - 6 electrons in outer energy level (2s and 4p) - properties vary from

nonmetallic to metalloid

o oxygen is the most abundant element on the earth

o sulfur has many industrial uses (sulfuric acid is the most widely used industrial chemical)

Group 7A - Halogens - all have 7 electrons in the outer energy level (2s and 5p) - properties vary from

nonmetallic to metalloid.

o Very reactive due to the outer energy level being almost filled.

o iodine is used as an antiseptic

o chlorine is a bleaching agent and disinfecting agent

o fluorine, as the fluoride ion, is used to maintain the health of our teeth

o fluorine is used to make Teflon

Lewis Dot Notation

Dot formulas use dots surrounding the symbol of the element to represent the valence electrons. The dot

formulas for period 2 and 3 would appear as follows.

Note electrons are usually shown as far apart as possible -they have the same charge and therefore repel each

other.

Gilbert Lewis, in 1916, proposed the octet rule: Atoms react by changing their number of electrons so as to

acquire the stable electron configuration of a noble gas (s2p6).

- The Noble gases (Group 0) have a stable electron configuration (s2p6) with 8 electrons filling the outer

s and p orbitals.

▪ This stability comes from the low energy state of this configuration and also accounts for the low

reactivity of these elements (most elements react with other elements to get to a lower, more stable

energy state).

• For example the halogens (Group 7A) have 7 valence electrons (s2p5) and want to gain one

electron to get the low energy, stable electron configuration of the noble gases.

- The elements in group 6 (s2p4) want to gain 2 electrons to get the low energy, stable electron

configuration of the noble gases.

- The Group 1A elements (s1) want to lose their outer electron to empty their outer shell and get a stable

electron configuration. For example if sodium (1s22s22p63s1) loses its 3s1 electron it will have filled s

and p orbitals in its outer energy level.

An exception to the octet rule is the electron configuration of helium.

Helium(1s2) is a noble gas, only it has only one orbital, the s orbital. It is filled and therefore stable and

elements close to it (lithium, beryllium and sometimes hydrogen) try to acquire its electron configuration by

losing or gaining electrons).

The pseudo noble-gas electron configuration has the outer three orbitals filled, the s, p and d- s2p6d10 (18

electrons total) and so is fairly stable. Elements that attain this electron configuration are at the right side of the

transition metals (d block).

An ion is a charged atom that is formed by the gaining or losing of electrons. When an atom gains or loses

electrons it is no longer neutral because the number of electrons (negative charges) and protons (positive

charges) are not equal. For example when a sodium atom loses an electron to get the noble gas electron

configuration of neon it gets a charge of +1and becomes the sodium ion, because it now has 11 protons (+) and

only 10 electrons(-). When a chlorine atom gains one electron to get the noble gas electron configuration of

argon it gets a charge of -1 and becomes the chloride ion, because it has 17 protons (+) and 18 electrons (-).

Ionic compounds are compounds formed when the elements bond together with ionic bonds - bonds formed by

the electrostatic charges (+ and -) formed when there is a transfer of electrons from one element to another.

Atoms bond together in order to achieve a lower, more stable electron configuration (noble gas or pseudo

noble-gas electron configurations). They can do this by transferring electrons from one element to another,

creating ions which then bond together due to their electrostatic charges. This type of bond is called an ionic

bond because it involves ions. This type of bonds occurs between metals (have few electrons in their outer

energy level so they want to get rid of those electrons) and nonmetals (have many electrons in their outer

energy level so they want to take in more electrons to get the noble gas electron configuration).

A common example of an ionic bond is between sodium (a metal) and chlorine (a nonmetal). Sodium

(1s22s22p63s1) needs to lose one electron (the 3s1) to achieve the noble gas

electron configuration of neon. Chlorine (1s22s22p63s23p5) needs to gain one

electron to get the noble gas electron configuration of argon. A Bohr model of

this bond would look like:

An electron dot representation of this electron transfer and bond would look like :

A quantum mechanical representation of the electron transfer would look like:

Characteristics of the Main Bond Types

o Ionic compounds (bonds) are : o bond is formed by electrostatic attraction of opposite charges

o crystalline solids at room temperature

o brittle

o electrolytes - conduct electricity when dissolved or melted

Metal atoms tend to lose electrons to become cations. Metallic

bonds are bonds between metallic cations. These bonds are formed

by the mutual attraction the cations have for free floating electrons.

It may be pictured as follows. When a hammer strikes a malleable

metal the cations will slide past each other and still be bonded

together by their mutual attractions for the free floating electrons.

When a hammer strikes an ionic crystal it forces the like charged ions

to come into close proximity. This creates a repulsive force which

may shatter the crystal.

o Metallic Bonds:

o metals are malleable - can be hammered into a shape as the cations slide past each other

o metals are ductile - can be drawn out into a wire as cations slide past each other

o metals are good conductors of electricity - the weakly held electrons are able to flow o Not soluble in water

A Comparison of Ionic and Covalent Compounds

Covalent Compounds