Chemistry

Fall 2003

Dr Supplee

Chapter 1- Definitions

• Science– Methodical exploration of nature followed by a

logical explanation of observations

• Scientific Method– A systematic investigation of nature and

requires proposing an explanation for the results of an experiment in the form of a general principle (hypothesis)

Chapter 1 - Definitions

• Hypothesis– Initial explanation of observations

• Theory– Sufficient evidence in support of the hypothesis– Model that scientifically explains the behavior of

nature

• Law– Does not explain behavior– States a measurable relationship under different

experimental conditions

Chapter 1 – Definition Examples

• Hypothesis– Dalton proposed that all matter was composed of

small individual particles (atoms)

• Theory– 100 years later Atomic Theory which

explains the composition of substances as well as the behavior of gases

• Law– Boyle’s Law P1V1= P2V2 at constant temperature

• If volume decreases than pressure increases at constant temperature

Chapter 1- Definitions Summary

Scientific Theory Natural Law

Hypothesis

Experiment

Analyze more data

Analyze initial observations

Chapter 1 – Modern Chemistry

• Organic Chemistry– Chemistry of carbon containing compounds

• (C, H, O, and N)

• Inorganic Chemistry– Chemistry of all other substances

• Biochemistry– Chemistry of substances derived from plant

substances

Chapter 1 – Modern Chemistry

• All three have in common– Analytical Chemistry

• Qualitative (what) and quantitative (how much) analyses

– Physical Chemistry • Theoretical and mathematical explanations of

chemical behavior

Relevance to daily life

Fun experiments

Benefits to society

Applications

Career Opportunities

Interesting topics

CHEMISTRY

Chapter 2- Scientific Measurements

• Introduction to Laboratory– Work alone– Handout– Due 9/15/03

• Measurement Uncertainty– Plus/minus factor ( error)

• Metric versus English Units– Conversion factors

• Significant Figures – Rounding rules

Precision versus Accuracy

True Value

Precise, not accurate

Accurate, not precise

Precision –how close two measurements of the same quantity are to each other

Accuracy – how close an experimental observations to the true value

Chapter 2- Scientific Measurements

• Measurement– a number with units

• Uncertainty– the estimated unit amount – plus/minus associated with measurement

• Mass– Amount of matter an object possesses

• Weight– Force exerted by gravity on an object

Chapter 2- Scientific Measurements

• Volume– Amount of space occupied by a solid, gas or

liquid

Significant Digits/ Figures

• Digits are significant when the do more than hold a decimal place– A place holder zero is NEVER significant

• determines measurement uncertainty (error analysis)

• Does not apply for exact numbers, only measured numbers

Significant Digits/ Figures Rule

• Rule #1– Count the number of nonzero digits left to

right– Do not count place holder zeros

Significant Figure Rounding Rules

• After all calculations are complete determine significant figures and then round– 5 or greater round-up to the nearest whole

number– less than 5 truncate

Scientific Notation• Exponential numbers (power of 10)

Base 10exponent

• The number 10 is raised to the nth power• Numbers greater than 1 the exponent is positive • Numbers less than 1 the exponent is negative• The decimal is placed after the first significant

digit and sets the size of the number by using a power of 10.

Unit Equations, Factors and Conversions• Problem Solving Technique

• Equivalent relationships

• Unit equation– A simple statement of two equivalent

quantities

• Unit Factor– A ratio of two equivalent quantities

Unit Dimensional Analysis Problem Solving• Three steps

1) write down the units asked for in the answer

2) write down the value given in the problem that is related to the required answer

3) Apply a unit factor to convert the units in the given value to the units in the answer

Given Value x Unit = units asked for

Factor

Percent Concept

• amount of a single quantity compared to the entire sample

• one part per 100 parts

one quantity x 100 = %

total sample

Review

Significant Digits/ Figures

• Digits are significant when the do more than hold a decimal place– If the number is less than 1, a place holder

zero is NEVER significant

• determines measurement uncertainty (error analysis)

• Does not apply for exact numbers, only measured numbers

Exact Numbers

• Infinite significant figures

• English to English conversion factors

• Metric to metric conversion factors

Unit Equations, Factors and Conversions• Problem Solving Technique

• Equivalent relationships

• Unit equation– A simple statement of two equivalent

quantities

• Unit Factor– A ratio of two equivalent quantities

Chapter 3 – The Metric System

• Single basic unit for each quantity measured

• Decimal system that uses a system of prefixes to enlarge or reduce a basic unit

Metric System Definitions

• Meter equals one ten-millionth of the distance from the North Pole to the equator

• Kilogram equals the mass of one a cube of water one-tenth of a meter on a side

• Liter equals the volume occupied by a kilogram of water at 4 oC

The Metric System

Physical Quantity Basic Unit Symbol

Length meter m

Mass gram g

Volume liter L

time second s

Metric Prefixes

Prefix Symbol Multiple/Fraction

giga- G 1 x 109

mega- M 1 x 106

kilo- k 1 x 103

deci- d 1 x 10-1

centi- c 1 x 10-2

milli- m 1 x 10-3

micro- 1 x 10-6

nano- n 1 x 10-9

Metric Conversion Factors Practice

• 1 kg =? g

k = kilo = 1000 basic units

1kg = 1000g

• 2s =? ns

n=nano=1 1 x 10-9

2s=2 x 10-9 ns

Unit Conversion Factors

• Ratio of two equivalent quantities

• The quantity in the numerator is equal to the quantity in the denominator

• If 100cm = 1 m, then the factor becomes

100 cm or 1m

1 m 100 cm

Metric- English Conversions

Physical Quantity

English UnitMetric

Equivalent

length 1 inch (in.) 1 in = 2.54 cm

mass 1 pound (lb) 1 lb = 454 g

volume 1 quart (qt) 1 qt = 946 mL

time 1 second (sec) 1 sec = 1.00 s

Unit Analysis

• Recall:– Problem Solving Technique

Units GivenUnit

FactorNew unit

Units asked for

UnitFactor

Practice Problems

• Work in groups of 3-4

• One student from each group puts solution in board and explains to class

Quiz # 4

• See Chemistry Current News Slides

• Presentation to be given on Oct 6, 2003.

Density - Review

• Lab Experiment 2• Physical property• Defined as mass per unit volume• Liquids and solids expressed in g/ml (g/ cm3)

• Gases expresses in grams per liter• Density of water is 1.00 g/ml• Floats in water density <1.00 g/ml• Sinks in water density >1.00g/ml

Estimating Density(page 59 and 60 )

Liquid 1

Liquid 2 = water

Liquid 3

Solid 1 = ice

Solid 2 = rubber

Solid 3 = aluminum

Water, chloroform and ethyl ether are poured into a tall glass cylinder. Three known solids are added. Identify the liquids.

TemperatureFahrenheit, Celsius and Kelvin• Measure of the average energy of

individual particles in a system– Warmer temps = more molecules moving thus

more energy– Cooler temps = slow moving molecules thus

less energy

• Fahrenheit oF • Celsius oC • Kelvin K

Temperature

• oF – Freezing point of water 32 oF– Boiling point of water 212 oF

• oC– Freezing point of water 0 oC– Boiling point of water 100 oC

• K ( SI unit)– Absolute zero 0 K– Equal to -273.15oC

Temperature Conversions

• oF to oC( oF - 32 oF ) x 100 oC / 180 oF = oC

• oC to oF( oC x 180 oF / 100 oC ) +32 = oF

• KelvinoC +273

Heat

• Heat measures the total energy• Temperature measures the average energy• Heat energy units calories or kilocalories• A calorie (cal) is defined as the amount of heat

needed to raise 1 g of water 1 oC • Food Calorie equals 1 kcal = 1000 cal• SI unit = joule (J)

1 cal = 4.184 J

Specific Heat

• Amount of heat required to bring about a given change in temperature

• Observed amount• Unique for each substance• Specific heat of water is high

– Change in temperature is minimal as water gains or losses heat

– Surface of earth is covered in water so water helps to regulate the climates

Specific Heat

• Amount of heat required to raise the temperature of 1 g of substance 1 oC

• Units are cal/g oC

1 g 1 g 1 g1 g

WaterIce Iron Silver

1.0 oC 2.0 oC 9.3 oC 17.7 oC

Specific Heat• gain or loss of heat divided by mass and

temperature change = specific heat

How many calories are required to raise the temperature of a 3 inch iron nail weighing 7.05 g form room temperature to 100 oC?

The specific heat of iron is 0.108 cal/g oC

Solution

• Specific Heat = Heat/ (mass x t)

cal/g oC = cal / g x oC

• 0.108 cal/g oC = cal / 7.05 x (100-25oC)

• Solving for Heat ( energy required )

• Rearrange (0.108 cal/g oC) x 7.05 g x 75 oC

= 57 cal

Chapter 4 Matter and Energy

• Matter is any substance that has mass and occupies volume

• Physical State changes– Melting solid into liquid– Sublimation solid into gas– Condensation gas into liquid– Deposition gas to solid– Freezes liquid to solid– Vaporization liquid to gas

Increasing temperature

ice watersteam

melting

freezing

vaporizing

condensing

Sublimation

Deposition

Chapter 4 Matter and Energy

Property Solid Liquid Gas

Shape Definite Indefinite Indefinite

Volume Fixed Fixed Variable

Compressibility negligible negligible significant

Elements, Compounds and Mixtures

• Properties of matter may be consistent throughout or they may vary

• Melting point– Gold (Au) 1064 oC– Quartz 1000 – 1600 oC

• Gold is homogenous – properties consistent

• Quartz is heterogeneous – properties vary

Mixtures

• Heterogeneous – Usually Solids– Separated into pure substances by physical methods

which take advantage of different physical properties– Properties are not the same throughout the sample

• Homogeneous– Gases or liquids – Separated into pure substances by either chemical or

physical methods which take advantage of different physical properties

– Properties a the same for any given sample, but can vary sample to sample

Mixtures• Alloy

– Homogeneous mixture of two or more metals– Gold ( Au)

10 K 14 K 18 K 42 % 75%

• Substance– Matter with definite composition and constant properties– Compound or an element

• Compound– Broken down into elements by chemical reactions

• Element– Cannot be broken down further by chemical reactions

Matter

Mixtures Substances

Heterogeneous Homogeneous Compounds Elements

Physical Separate

Names and Symbols of the Elements

• 81 stable elements that occur in nature

• Only 10 account for 95% of the mass of the earths crust, water and atmosphere

ElementMass Percent

ElementMass Percent

O 49.5 Na 2.6

Si 25.7 K 2.4

Al 7.5 Mg 1.9

Fe 4.7 H 0.9

Ca 3.4 Ti 0.6

All other elements 0.5 %

Names and Symbols

• Names are from various sources– Hydrogen (hydro, Gr. = water former)– Carbon (carbo, Lt. =coal)– Calcium (calcis, Lt. = lime)

• Chemical Symbols– Dalton in 1803 proposed that elements are composed

of indivisible spherical particles or atoms (atomos, Gr. = indivisible)

– Suggested the use of circles with markings for symbols ( pg 83)

– Berzelius in 1813 proposed our current system of symbols = using the first letter of the name and if the first letter is already in use two letters

Metals, Nonmetals and Semimetals

• Predict by position in Periodic Table• Metals

– solid element– Bright metallic luster– Good conductor of heat and electricity– High density– High melting point– Malleable ( thin sheets)– Ductile ( fine wire)

Metals, Nonmetals and Semimetals

• Nonmetals– Solid or gas element – Dull appearance– Low density– Low melting point– Poor conductor of heat and electricity– Crush to a powder, if solid

• Semimetal– metalloids– midway between metal and nonmetal– Semiconductor

Periodic Table of the Elements(page 86)

• Atomic number– Number of protons

• Metals are placed on the left

• Nonmetals on the right

• Separated by semimetals starting at B

• Solids are to the left (most all elements)

• Gases to the right

Compounds & Chemical Formulas

• 1799 Proust

• Law of Definite Composition

• Law of Constant Proportion

“Compounds always contain the same elements in a constant proportion by mass.”

Chemical Formulas

• Most elements occur in nature as collection of individual atoms

• Diatomic molecules– Oxygen (O2), Hydrogen (H2), Nitrogen (N2)– Halogens Chlorine (Cl2),Bromine Br2, Iodine (I2)

• A chemical formula expresses the number and type of each atom in a compound

• The number of the each atom is indicated with a subscript. The number 1 in the subscript is implied and therefore is omitted.

• Parentheses are used to help clarify the structure of the compound and

Examples

• Water H2O– 2 hydrogen atoms and 1 oxygen atom

• Calcium Chloride CaCl2– 1 calcium atom and 2 chlorine atoms

• Propylene Glycol C3H8O2 – 3 carbon atoms, 8 hydrogen atoms, 2 oxygen atoms

• Lead acetate PbC2H3O2– 1 lead atom, 2 carbon atoms, 2 oxygen atoms, 3

hydrogen atoms• 4-amino-2-hydroxytoluene C7H9NO

– 7 carbon atoms, 9 hydrogen, 1 nitrogen, 1 oxygen

Physical and Chemical Properties

• Substances ( Compounds or Elements)– Physical and chemical properties are the

consistent throughout– No two substances have all the same physical

and chemical properties

2 Na + Cl2 2 NaCl

metal yellow gas white solid

Physical and Chemical Properties

• Physical properties are measured and observed without changing the chemical composition of the substance

• Examples– Appearance, color, melting point, density, solubility,

boiling point, freezing point

• Chemical properties describes how a substances reacts with other substances (reaction chemistry) and always involve a chemical change

Periodic Table and Reaction Chemistry• Elements with similar reaction chemistry

are “Grouped” into families– Columns in the Periodic table

• Group IA Alkali Metals• Group IIA Alkaline Earth Metals• Group VIIA Halogens• Group VIIIA Noble Gases• Group IIlB to IIB Transition Metals • Lanthanides• Actinides

Physical and Chemical Changes

• Physical Change – chemical composition does not change, but the physical state does

• Example ?• Chemical Change – chemical composition

changes and the physical state may or may not change ( formation of a new substance)

• Example ?

Physical and Chemical Changes

• Chemical changes are often detected by– Gas formation bubbles or odor– Color change permanent– Release of energy light or

heat– Precipitate formation solids

Conservation of Mass and Energy

• Matter is nether created or destroyed during a chemical reaction

• Energy can not be created or destroyed. It can however, be converted from one form to another

• Total mass and energy of the universe is constant

Potential vs. Kinetic Energy

• Key concepts for understanding chemical reactions

• Potential energy is stored energy of matter at rest

• Kinetic energy is energy that is a result of motion

• Potential energy can be (and is) converted into kinetic energy

Kinetic Energy as a function of Physical State

Particles Solid Liquid Gas

Kinetic Energy

Very low High Very high

Movement None Restricted unrestricted

Chemical Reactions

• Potential Chemical Energy • Kinetic Heat Energy• Exothermic reactions

– Reactions that give off or produce heat – Reactants have more potential energy than products

• Endothermic reactions – Reactions that take in or absorb heat – Reactants have less potential energy than products

Example

Exothermic

Reactants Products + heat

( high P.E.) (low P.E. ) ( K.E.)

Endothermic

Reactants + heat Products

( low P.E.) + (K.E.) (high P.E. )

Forms of Energy

• Six Forms– Light– Heat– Chemical– Electrical– Mechanical – Nuclear

Energy Examples

• Radioactive Ur vaporizing water

nuclear heat

• Steam driving a turbine

heat mechanical

• Lead acid battery

chemical electrical

Chapter 5 – Models of the Atom

• Atom – indivisible

• Dalton– Proposed that all matter was composed as

tiny particles– Based on the Laws of Conservation of Mass

and Definite Proportion– Compounds are simply the combination of two

or more atoms of different elements

Dalton’s Atomic Theory

1. An element is composed of tiny, indivisible particles called atoms

2. All atoms of an element are identical and have the same properties

3. Atoms of different elements combine to form compounds

4. Compounds contain atoms in small whole number ratios

5. Atoms can combine in more than one ratio to form different compounds

Thomson Model (Plum Pudding Model)

• Cathode Ray experiment– Glass tubes containing a low pressure amount of gas emitted

light when electricity was applied to one end of the tube (Fluorescence – light energy)

– Ray emanates form the negative cathode in the tube, the radiation is referred to as a cathode ray

– Placed tube in magnetic field, light or ray curved towards the positive

– Concluded that cathode rays were composed of tiny negatively charged particles ( electrons, e- )

– Further experiments showed that certain rays contained small particles that had an equal but opposite in sign charge to electrons protons (p+)

Relative Charges and Masses

Subatomic Particle

SymbolRelative Charge

Relative Mass

Electron e- -1 1/1836

Proton p+ +1 1

Thomson Model of the Atom

+

Homogeneous sphere“plum pudding”

- +

--

--

-+

-+

+

Rutherford Model of the Atom

• Alpha ray –– particles identical to He w/o electrons ( He +2 ) – most passed through thin Au foil– suggested that most of the atoms were empty space,

with electrons moving about a center ( nucleus)– Nucleus contains protons and is tiny and dense

• Beta ray - • Gamma ray – not affected by magnetic fields

Rutherford Model of the Atom

++ ++ + 1 X 10 -8 cm

1 x 10 -13 cm

- -

- - - -

Neutrons “heaviness of the atom”

n 0

Subatomic Particles

Particle Symbol LocationRelative Charge

Relative Mass

Electron e-Outside nucleus

-1 1/1836

Proton p+Inside nucleus

+1 1

Neutron n0 Inside nucleus

0 1

Atomic Notation

Sy symbol

A

ZAtomic Number ( protons)

Mass Number(protons and neutrons)

Isotopes

• Same element with different amount of neutrons– Number of Protons and Electrons are the

same

• Mass of the element is different, so the mass number is different

• For example, Hydrogen and deuterium

H H11

21

Isotopes

• Naming– Some are common (hydrogen, deuterium, tritium)– Name of the element followed by the mass number– For example,

carbon -14oxygen -18cobalt-60

Hint: should be able to determine number of protons electrons, neutrons atomic notation, isotope names

Atomic Mass

• Atoms are to small to weigh on balance

• Masses are relative to each other

• Specifically, mass is relative to carbon-12

• Carbon-12 has 12 amu (atomic mass unit)

• So, an amu = 1/12 the mass of carbon

• Weighted average

Atomic Mass

• Weighted average of all the naturally occurring isotopes

• Given the natural abundance of the atom, the amu can be calculated for a given atom

• Since it’s a calculated value no atom will actually weigh this number

Periodic Table

• Does not tell about the number of naturally occurring isotopes

• Atomic Number = Protons• Atomic Mass = weighted average = protons and

neutrons (whole numbers)• Mass number is in parentheses, then the

element is unstable. Mass is given for the best known isotope.

Hint: Should be able to tell form periodic table which elements are stable and which are not.

Wave Nature of Light

• Wavelength – the distance the light wave travels to complete one cycle

• Frequency - the number of wave cycles completed in 1 s

• Velocity of light is constant = 3.00 x 103 m/s• If wavelength decreases, frequency increases• Low frequency – low light energy –long

wavelengths• High frequency-high energy light –short

wavelengths ( page 124)

Light- A Continuous Spectrum

• White light passes through a prism it separates into all the colors– ROY G BIV

• Light – radiant energy that is visible

• Visible spectrum = 400-700 nm

• Radiant energy spectrum – a continuous spectrum of visible and invisible light that ranges form short to long wave lengths

Radiant Energy Spectrum

Cosmic Rays

400 nm 500 nm 600 nm 700 nm

Visible Spectrum

Gamma Rays

X RaysUV / VIS

Infrared

Microwaves TV Radio

Wave Length Increases