Zumdahls Chemistry

1 AP Chemistry

A. Allan

Chapter 1 Notes - Chemical Foundations 1.1 Chemistry: An Overview

A. Reaction of hydrogen and oxygen

1. Two molecules of hydrogen react with one molecule of oxygen to form two molecules of water

2H2 + O2 à 2H2O 2. Decomposition of water

2H2O à 2H2 + O2 B. Problem Solving in Chemistry (and life)

1. Making observations 2. Making a prediction

3. Do experiments to test the prediction

1.2 The Scientific Method A. General Framework

1. Making observations

a. Quantitative ( measurement)

b. Qualitative (color, phase, shape, etc) 2. Making a prediction

3. Do experiments to test the prediction B. Vocabulary

1. Observation

a. Something that is witnessed and can be recorded 2. Theory (Model)

a. Tested hypotheses that explains WHY nature behaves in a certain way

b. Theories change as more information becomes available 3. Natural Law

a. A summary of observed, measurable behavior

1.3 Units of Measurement A. Measurements

1. Number and Scale (units) are both essential "The number without the units is worthless!" B. SI system

Important SI Units for Chemistry

Mass kilogram kg

Length meter m

Time second s

Temperature Kelvin K

Amount of Substance mole mol

2 C. SI Prefixes mega M 1,000,000 106 nano n 0.000000001 10-9 kilo k 1,000 103 micro Î 0.000001 10-6 hecto h 100 102 milli m 0.001 10-3 deka da 10 101 centi c 0.01 10-2 deci n 0.1 10-1 1.4 Uncertainty in Measurement

A. Recording Measurements (Significant figures) 1. Record all digits that are certain

2. Record the first digit that is uncertain (all measurements have some degree of uncertainty)

3. Uncertainty in the last number is + 1, unless otherwise noted B. Accuracy

1. The agreement of a particular value with the accepted value C. Precision

1. The degree of agreement among several measurements made in the same way

"You can be precise, but not accurate. If you are accurate, you are necessarily precise."

D. Errors

1. Random Errors (indeterminate errors) a. Measurements may be high or low b. Causes:

1) Interpretation of the uncertain digit 2) Procedural ineptness

2. Systematic Errors

a. Always occur in the same direction

b. Caused by poor measurement calibration 1) gun sight set too high/low

2) balance improperly zeroed 3) thermometer improperly marked

3 1.5 Significant Figures and Calculations

A. Rules for Counting Significant Figures

Number Rule Example

Nonzero integers Always significant 6.34 m (3 sig figs) Leading zeroes Never significant 0.00634 m ( 3 sig figs) Captive zeroes Always significant 6.0034 (5 sig figs) Trailing zeroes Significant if after a decimal 63400 (3 sig figs)

0.63400 (5 sig figs) Exact numbers Infinite significance e.g. There is 1 star at the

center of our solar system. There is no doubt about the number "1"

Scientific notation All digits are significant 6.3400 x 106 (5 sig figs) B. Multiplication and Division

1. Keep as many sig figs in your answer as are in the piece of data with the least number of sig figs

2.37 cm x 15.67 cm x 7.4 cm = 274.82046 (keep two sig figs) = 2.7 x 102 cm3 C. Addition and Subtraction

1. Keep the same number of decimal places as the least precise measurement in your calculation

34.039 m + 0.24 m + 1.332 m + 12.7 m = 48.311 m (keep one decimal place) = 48.3 m D. Rules for Rounding

1. Round at the end of a series of calculations, NOT after each step 2. Use only the first number to the right of the last sig fig to decide

whether or not to round

a. Less than 5, the last significant digit is unchanged b. 5 or more, the last significant digit is increased by 1

Note from this section in your book: Detailed solutions and stepwise examples of problems in this text show correct sig figs at each step. Since you will most often do a sequence of calculations and then round to correct sig figs at the end, your answer will often be slightly different (usually only in the last place) than the answer given in the book.

1. Significant figures rules will be observed in all calculations throughout the year in this course. You need never ask, "Do we have to watch our sig figs?" The answer is always "Yes!"

4 1.6 Dimensional Analysis

A. Examine examples 1. pages 18 - 21

B. Unit Conversions Questions 1. What units am I given?

2. What units must be in my answer? 3. What is conversion factor?

Full credit can never be given for working a problem in which you do not do all of the following:

1. Observe significant figures rules

2. Label all steps of your work with the correct units 3. Correctly label and identify your answer

4. Solve the problem in a manner that can be understood by the reader.

1.7 Temperature

A. Celsius (°C) and Kelvin (K)

1. Kelvin = Celsius + 273.15 2. Celsius = Kelvin - 273.15

3. Size of the temperature unit (degree) is the same B. Fahrenheit 1. T_C = (T_F - 32°F)(5°C/9°F) 2. T_F = T_C x (9°F/5°C) + 32°F 1.8 Density A. Density = mass/volume 1.9 Classification of Matter A. Matter

1. Anything that occupies space and has mass B. States of Matter

1. Solids - rigid, fixed volume and shape 2. Liquids - definite volume, no specific shape

3. Gases - no fixed volume or shape, highly compressible C. Mixtures - Matter of variable composition

1. Heterogeneous mixtures

a. Having visibly distinguishable parts 2. Homogeneous mixtures (solutions)

a. Having visibly indistinguishable parts

D. Components of Mixtures can be Separated by Physical Means 1. Distillation

2. Filtration

5 E. Pure substances

1. Elements

a. Cannot be decomposed into simpler substances by physical or chemical means

2. Compounds

a. Constant composition

b. Can be broken into simpler substances by chemical means, not by physical means

The Organization of Matter (Slightly different than your book)

Matter Mixtures: a) Homogeneous (Solutions) b) Heterogeneous Pure Substances Compounds s Elements Atoms Nucleus Electrons Protons Neutrons Quarks Quarks

AP Chemistry A. Allan

Chapter 2 Notes - Atoms, Molecules and Ions 2.1 The Early History

Refer to the Chemistry History Timeline for this chapter

2.2 Fundamental Chemical Laws

A. Law of Conservation of Mass

1. "Mass is neither created nor destroyed"

2. Translation: In ordinary chemical reactions, the total mass of the reactants is equal to the total mass of the products

B. Law of Definite Proportion

1. "A given compound always contains the same proportions of elements by mass"

2. Translation: Compounds have an unchanging chemical formula C. Law of Multiple Proportions

1. "When two elements form a series of compounds, the ratios of the masses of the second element that combine with one gram of the first element can always be reduced to small whole numbers

2. Translation: Sometimes two elements can come together in more than one way, forming compounds with similar, though not identical

formulas

2.3 Dalton's Atomic Theory A. Atomic Theory

1. Each element is made up of tiny particles called atoms 2. The atoms of a given element are identical

3. Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms

4. Chemical reactions involve reorganizations of the atoms. The atoms themselves are not changed in a chemical reaction

B. Avogadro's Hypothesis

1. At the same conditions of temperature and pressure, equal volumes of different gases contain the same number of particles.

2.4 Early Experiments to Characterize the Atom A. J.J. Thomson and the Electron

1. Determined the charge to mass ratio of the electron 2. Reasoned that all atoms must contain electrons

3. Reasoned that all atoms must contain positive charges B. Robert Millikan and the Oil Drop

1. Oil drop experiments determined the charge on an electron 2. With charge information, and Thomson's charge/mass ratio, he

C. Radioactivity

1. Gamma (È) rays - high energy light 2. Beta (Ã) particles - high speed electrons

3. Alpha (Â) particles - nuclear particle with a 2+ charge D. The Nuclear Atom - Rutherford's Metal Foil Experiment

1. Most alpha particles pass straight through thin metal foil

2. Some particles were greatly deflected ("like a howitzer shell bouncing off of a piece of paper")

a. Could not have been deflected by electrons or single protons b. Must have been deflected by a positively charged object of

substantial mass

1) Supported concept of a small, central, positive nucleus where most of the atom's mass was concentrated 2) Disproved Thomson's "plum pudding" model

2.5 The Modern View of Atomic Structure: An Introduction A. Nucleus

1. Protons - positively charged 2. Neutrons - no charge

3. Small size, high density

a. The mass of all of the cars in the United States in an object that would easily fit in a teaspoon

b. A pea with the mass of 250 million tons B. Electrons

1. Negatively charged

2. The source of varying reactivity of different elements 3. Provide most of the atomic volume

C. Atomic Number

1. Number of protons D. Mass Number

1. Number of protons + number of neutrons E. Isotopes

1. Atoms with the same number of protons (same element) but different numbers of neutrons (mass numbers)

F. Symbols for the Elements 1. Mass Number

23

11Na Element symbol Atomic Number

2.6 Molecules and Ions A. Chemical Bonding

1. Covalent bonding - Sharing of electrons

2. Ionic bonding - Attraction of oppositely charged ions due to a reaction in which electrons are transferred

B. Representing Molecules (Covalently bonded) 1. Chemical formula

a. Symbols for atoms and subscripts 1) H₂0

2) CH₄ 2. Structural formula

a. Bonds represented by lines

Ball and Stick Space Filling

C. Ions

1. Cations

a. Positive ions formed by the loss of electrons 2. Anions

a. Negative ions formed by gaining electrons D. Ionic Bonding

1. Bond formed by the attraction between oppositely charged ions 2. Ionic bonding forms ionic solids (salts)

3. Ions can be monatomic (one atom) or polyatomic (more than one atom)

2.7 An Introduction to the Periodic Table A. Organization

1. Horizontal row is called a "period" (or series) 2. Vertical column is called a "group" or "family"

a. Group 1A - Alkali metals

b. Group 2A - Alkaline earth metals

c. Group 7A - Halogens (Gr, "salt makers") d. Group 8A - Noble gases

B. Naming Elements 104 and beyond

Nil = 0 un = 1 bi = 2 tri = 3 quad = 4 Pent = 5 hex = 6 sept = 7 oct = 8 enn = 9 Element 109 = un (1) nil(0) enn(9) ium = unnilennium

2.8 Naming Simple Compounds A. Ionic Compounds

1. Positive ion is always named first, negative ion second

You were given a list of ions to memorize on the first day of class Tips for memorizing the polyatomics:

a. Find the "ate" ion (sulfate, for instance) sulfate = SO₄

2-b. The "ite" ion always has one less oxygen than the "ate" ion sulfite = SO₃

2-c. The prefix "per" (think hyper, meaning "above") is used with the "ate" prefix to indicate one more oxygen than the "ate" ion

persulfate = SO₅

2-d. The prefix "hypo" (meaning "under" or "below") is used with the "ite" prefix to indicate one less oxygen than the "ite" ion

hyposulfite = SO₂

2-Examples (Just because you can name it doesn't mean it exists!)

Perchlorate _ClO4- Pernitrate _NO4

-Chlorate _ClO3- Nitrate _NO3

-Chlorite _ClO2- Nitrite _NO2

-hypochlorite _ClO- hyponitrite

NO-2. Metals with more than one oxidation state (transition metals) must have a roman numeral to indicate the oxidation state

Fe3+ _{= iron (III) Mn}+2 _{= manganese (II)} B. Binary Covalent Compounds

1. Must contain two elements, BOTH nonmetals a. First element

1) full element name

2) prefix only if there is more than one atom b. Second element

1) named as if it were an anion (-ide suffix) 2) always gets a prefix

mono - 1 penta - 5 octa - 8

di - 2 hexa - 6 nona - 9

tri - 3 hepta - 7 deca - 10

tetra - 4

C. Naming Acids

1. Binary Acids (two elements - hydrogen + one other) a. prefix "Hydro" + root of second element + "ic" suffix 2. Oxyacids

a. If the acid contains an anion whose name ends in "ate":

Use root of anion name and an "ic" ending (H₂SO₄ = sulfuric acid) b. If the acid contains an anion whose name ends in "ite":

Use the root of the anion name and an "ous" ending (H₂SO₃ = sulfurous acid)

AP Chemistry A. Allan

Chapter 3 Notes - Stoichiometry 3.1 Atomic Masses

A. C-12, the Relative Standard

1. C-12 is assigned a mass of exactly 12 atomic mass units (amu) 2. Masses of all elements are determined in comparison to the carbon

-12 atom (12C) the most common isotope of carbon 3. Comparisons are made using a mass spectrometer B. Atomic Mass (Average atomic mass, atomic weight)

1. Atomic masses are the average of the naturally occurring isotopes of an element

2. Atomic mass does not represent the mass of any actual atom 3. Atomic mass can be used to "weigh out" large numbers of atoms

3.2 The Mole

A. Avogadro's number

1. 6.022 x 1023 units = 1 mole

2. Named in honor of Avogadro (he did NOT discover it) B. Measuring moles

1. An element's atomic mass expressed in grams contains 1 mole of atoms of that element

a. 12.01 grams of carbon is 1 mole of carbon b. 12 grams of carbon-12 is 1 mole of carbon-12

3.3 Molar Mass

A. Molar Mass (Gram molecular weight)

1. The mass in grams of one mole of a compound

2. The sum of the masses of the component atoms in a compound a. Molar mass of ethane (C₂H₆):

Mass of 2 moles of C = 2(12.01 g) Mass of 6 moles of H = 6(1.008 g) 30.07 g

3.4 Percent Composition of Compounds A. Calculating any percentage

1. "The part, divided by the whole, multiplied by 100" B. Percentage Composition

1. Calculate the percent of each element in the total mass of the compound

(#atoms of the element)(atomic mass of element) x 100 (molar mass of the compound)

3.5 Determining the Formula of a Compound A. Determining the empirical formula

1. Determine the percentage of each element in your compound 2. Treat % as grams, and convert grams of each element to moles of

each element

3. Find the smallest whole number ratio of atoms

4. If the ratio is not all whole number, multiply each by an integer so that all elements are in whole number ratio

B. Determining the molecular formula 1. Find the empirical formula mass

2. Divide the known molecular mass by the empirical formula mass, deriving a whole number, n

3. Multiply the empirical formula by n to derive the molecular formula

3.6 Chemical Equations A. Chemical reactions

1. Reactants are listed on the left hand side 2. Products are listed on the right hand side 3. Atoms are neither created nor destroyed

a. All atoms present in the reactants must be accounted for among the products, in the same number

b. No new atoms may appear in the products that were not present in the reactants

B. The Meaning of a Chemical Reaction 1. Physical States

a. Solid - (s) b. Liquid - (l) c. Gas - (g)

d. Dissolved in water (aqueous solution) - (aq) 2. Relative numbers of reactants and products

a. Coefficients give atomic/molecular/mole ratios

3.7 Balancing Chemical Equations

A. Determine what reaction is occurring

1. It is sometimes helpful to write this in word form: Hydrogen + oxygen à water

B. Write the unbalanced equation

1. Focus on writing correctly atomic and compound formulas H2 + O2 à H2O

C. Balance the equation by inspection

1. It is often helpful to work systematically from left to right 2H2 + O2 à 2H2O

D. Include phase information

3.8 Stoichiometric Calculations: Amounts of Reactants and Products A. Balance the chemical equation

B. Convert grams of reactant or product to moles

C. Compare moles of the known to moles of the desired substance 1. A ratio derived from the coefficients in the balanced equation D. Convert from moles back to grams if required

3.9 Calculations Involving a Limiting Reactant

A. Concept of limiting reactant (limiting reagent):

" I want to make chocolate chip cookies. I look around my kitchen (I have a BIG kitchen!) and find 40 lbs. of butter, two lbs. of salt, 1 gallon of vanilla extract, 80 lbs. of chocolate chips, 200 lbs. of flour, 150 lbs. of sugar, 150 lbs. of brown sugar, ten lbs. of baking soda and TWO eggs. It should be clear that it is the number of eggs that will determine the number of cookies that I can make."

1. The limiting reactant controls the amount of product that can form

B. Solving limiting reactant problems

1. Convert grams of reactants to moles

2. Use stoichiometric ratios to determine the limiting reactant

3. Solve as before, beginning the stoichiometric calculation with the grams of the limiting reactant

C. Calculating Percent Yield

1. Actual yield - what you got by actually performing the reaction 2. Theoretical yield - what stoichiometric calculation says the reaction

SHOULD have produced

Actual Yield x 100% = percent yield Theoretical Yield

AP Chemistry A. Allan

Chapter 4 Notes - Types of Chemical Reactions and Solution Chemistry 4.1 Water, the Common Solvent

A. Structure of water

1. Oxygen's electronegativity is high (3.5) and hydrogen's is low (2.1) 2. Water is a bent molecule

3. Water is a polar molecule B. Hydration of Ionic Solute Molecules

1. Positive ions attracted to the oxygen end of water

2. Negative ions attracted to the hydrogen end of water C. Hydration of Polar Solute Molecules

1. Negative end of polar solute molecules are attracted to water's hydrogen

2. Positive end of polar solute molecules are attracted to water's oxygen D. "Like Dissolves Like"

1. Polar and ionic compounds dissolve in polar solvents like water 2. Nonpolar compounds like fats dissolve in nonpolar solvents like

____?_____

4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes A. Definition of Electrolytes

1. A substance that when dissolved in water produces a solution that can conduct an electric current

B. Strong electrolytes conduct current very efficiently 1. Completely ionized when dissolved in water

a. Ionic compounds

b. Strong acids (HNO3(aq_{), H2SO4(}aq), HCl(aq)) c. Strong bases (KOH NaOH)

C. Weak electrolytes conduct only a small current 1. Slightly ionized in solution

a. Weak acids (organic acids - acetic, citric, butyric, malic)

b. Weak bases (ammonia) D. Nonelectrolytes conduct no current

1. No ions present in solution a. alcohols, sugars

4.3 The Composition of Solutions A. Molarity

1. Moles of solute per liter of solution

B. Concentration of Ions in Solution

1. Ionic compounds dissociate in solution, multiplying the molarity by the number of ions present

C. Moles from Concentration

1. Liters of solution x molarity = moles of solute D. Solutions of Known Concentration

1. Standard solution - a solution whose concentration is accurately known 2. Preparation of Standard solutions

How much x How strong x What does it weigh?

L x mol/L x g/mol = grams required to prepare the standard

E. Dilution

1. Dilution of a volume of solution with water does not change the number of moles present

2. Solving dilution problems M₁V₁ = M₂V₂ 4.4 Types of Chemical Reactions

A. Precipitation reactions

1. When two solutions are mixed, an insoluble solid forms B. Acid-Base reactions

1. A soluble hydroxide and a soluble acid react to form water and a salt C. Oxidation-Reduction reactions (redox rxns)

1. Reactions in which one or more electrons are transferred

4.5 Precipitation Reactions A. Dissociation

1. Ionic compounds dissolve in water and the ions separate and move independently

AgNO₃(aq) + NaCl(aq) à products

Ag+ (aq) + NO₃-(aq) + Na+(aq) + Cl- à products B. Determination of Products

1. Recombination of ions

a. AgNO₃ NaCl AgCl NaNO₃ 2. Elimination of reactants as products

a. AgNO₃ and NaCl are reactants and can't be products 3. Identifying the precipitate

a. "Switch Partners" of reactant pairs to determine the names of the products.

b. AgCl and NaNO₃ are the products

c. AgCl is insoluble, so it is the white precipitate

d. If there is no insoluble product, the reaction does not occur a. NaCl(aq) + KNO₃(aq) à NaNO₃(aq) + KCl(aq)

Both products are soluble and all ions remain independent in solution; no reaction occurs Na+(aq) + Cl-(aq) + K+(aq) + NO₃-(aq)

Table 4.1 Simple Rules for the Solubility of Salts in Water 1. Most nitrate (NO3-)salts are soluble.

2. Most salts containing the alkali metal ions (Li+, Na+, K+, Cs+, Rb+) and the ammonium (NH₄+) ion are soluble.

3. Most chloride, bromide and iodide salts are soluble. Exceptions are salts containing the ions Ag+, Pb2+, and Hg2+.

4. Most sulfate salts are soluble. Notable exceptions are BaSO₄, PbSO₄, HgSO₄ and CaSO₄.

5. Most hydroxide salts are only slightly soluble. The important soluble hydroxides are NaOH and KOH. The hydroxides of barium, strontium and calcium are marginally soluble.

6. Most sulfide (S2-), carbonate (CO₃2-), chromate (CrO₄2-) and phosphate (PO₄3-) salts are only slightly soluble.

4.6 Describing Reactions in Solution A. The Molecular Equation

1. Gives the overall reaction stoichiometry, not necessarily the actual forms of reactants and products in solution

Na₂CO₃(aq) + Ca(NO₃)₂(aq) à 2NaNO₃(aq) + CaCO₃(s) B. The Complete Ionic Equation

1. Represents as ions all reactants and products that are strong electrolytes

2Na+(aq) + CO₃2-(aq) + Ca2+(aq) + 2NO₃-(aq) à 2Na+(aq) + 2NO₃-(aq) + CaCO₃(s) C. The Net Ionic Equation

1. Includes only those components that take part in the chemical change 2. Spectators are eliminated

Ca2+(aq) + CO₃2-(aq) à CaCO₃(s) 4.7 Stoichiometry of Precipitation Reactions

A. Determine what reaction takes place

B. Write the balanced net ionic equation for the reaction C. Calculate the moles of reactants

D. Determine which reactant is limiting

E. Calculate the moles of product or products F. Convert to grams or other units, as required

4.8 Acid-Base Reactions (Neutralization Reactions) A. Definitions

1. Br∅nsted:Acids are proton donors, bases are proton acceptors

2. Arrhenius: Acids produce H+ ions in water, bases produce OH- ions in solution

B. Net ionic equation for acid-base reactions 1. H+(aq) + OH-(aq) à H₂O(l)

2. The hydroxide ion can be assumed to completely react with even a weak acid in solution

C. Stoichiometry Calculations for Acid-Base Reactions

1. List the species present in the combined solution before any reaction occurs; decide what reaction will occur

2. Write the balanced net ionic equation for this reaction 3. Calculate the moles or reactants

a. For reactions in solution, use volumes of the original solutions and their molarities

4. Determine the limiting reactant where appropriate 5. Calculate the moles of the required reactant or product 6. Convert to grams or volume of solution as required D. Acid-Base Titrations

1. Vocabulary

a. Titrant - Solution of known concentration b. Analyte - Solution of unknown concentration

c. Equivalence point - Point at which the amount of titrant added to analyte results in perfect neutralization

d. Indicator - a substance added at the beginning of the titration that changes color at the equivalence point

e. Endpoint - the point at which the indicator changes color 2. Requirements for a successful titration

a. the exact reaction between titrant and analyte must be known b. the reaction must proceed rapidly

c. the equivalence point must be marked accurately (select the appropriate indicator)

d. the volume of titrant required to reach the equivalence point must be known accurately

e. for acid-base titrations, the titrant should be a strong acid or a strong base

4.9 Oxidation-Reduction Reactions (redox) A. Electron transfer (LEO says GER)

1. Gain electrons = reduction 2. Lose electrons = oxidation B. Examples of redox rxns

1. Photosynthesis 2. Combustion of fuels

Rules for Assigning Oxidation Numbers Summary 1. the oxidation number of the atom of a free element

is zero

Element = 0

2. the oxidation number of a monatomic ion equals its charge

3. In compounds, oxygen has an oxidation number of -2, except in peroxides, where it is -1

Oxygen = -2

4. In compounds containing hydrogen, hydrogen has an oxidation number of +1

Hydrogen = +1

5. In compounds, fluorine is ALWAYS assigned an oxidation number of -1

Fluorine = -1

6. The sum of the oxidation states for an electrically neutral compound must be zero

C. Noninteger Oxidation states 1. Fe3O4 - Magnetite

a. Oxidation number for each iron averages to +8/3 b. Magnetite contains two Fe3+ ions and one Fe2+ D. Characteristics of Oxidation-Reduction Reactions

1. the oxidized substance: a. loses electrons

b. increases oxidation state c. is the reducing agent 2. the reduced substance

a. gains electrons

b. decreases oxidation state c. is the oxidizing agent

4.10 Balancing Oxidation-Reduction Equations

There are no notes for this section. The only way to master the balancing of redox equations is to actually balance them. While there are some minor variations in the processes used for acidic and basic solutions, the skills involved are identical. We will practice balancing numerous redox equations as a class.

AP Chemistry A. Allan

Chapter 5 - Gases 5.1 Pressure

A. Properties of gases

1. Gases uniformly fill any container 2. Gases are easily compressed

3. Gases mix completely with any other gas 4. Gases exert pressure on their surroundings

a. Pressure = force/area B. Measuring barometric pressure

1. The barometer

a. Inventor - Evangelista Torricelli (1643) 2. Units

a. mm Hg (torr)

(1) 760 torr = Standard pressure b. newtons/meter2 = pascal (Pa)

(1) 101,325 Pa = Standard pressure c. atmospheres

(1) 1 atmosphere = Standard pressure

5.2 The Gas Laws of Boyle, Charles, and Avogadro A. Boyle's Law (Robert Boyle, 1627 - 1691)

1. the product of pressure times volume is a constant, provided the temperature remains the same

k PV =

a. P is inversely related to V

b. The graph of P versus V is hyperbolic

c. Volume increases linearly as the pressure decreases (1/P) 2. At constant temperature, Boyle's law can be used to find a new

volumes or a new pressure a. P1V1 = k = P2V2 ∴

V

P

V

P

V

V

P

P

3. Boyle's law works best at low pressures

4. Gases that obey Boyle's law are called Ideal gases B. Charles' Law (Jacques Charles, 1746 - 1823)

1. The volume of a gas increase linearly with temperature provided the pressure remains constant

a. V = bT V/T = b (1) V1/T1 = b = V2/T2

T

V

T

V

T

T

V

V

b. Temperature must be measured in degrees Kelvin (1) K = °C + 273

(2) 0 K is "absolute zero" C. Avogadro's Law (Amedeo Avogadro, 1811)

1. For a gas at constant temperature and pressure, the volume is directly proportional to the number of moles, n

a. V = an V/n = a (1) V1/n1 = a = V2/n2 ∴

n

V

n

V

5.3 The Ideal Gas Law

A. Derivation from existing laws 1. V = k/P V = bT V = an

2. Constants k, b, a are combined into universal gas constant, R

P

nRT

V

=

_PV

=

_nRT

mol

K

atm

L

R

∗

∗

=

0

.

0826

B. Limitations of the Ideal Gas Law

1. Works well at low pressures and high temperatures 2. Most gases do not behave ideally above 1 atm pressure 3. Does not work well near the condensation conditions of a gas C. Solving for new volumes, temp or pressure (n remaining constant)

1. Combined law (from general chem) 2.

T

V

P

R

n

T

V

P

T

V

P

T

V

P

T

T

P

P

V

V

( )( )( )













=

P

Tn

a

b

k

V

5.4 Gas Stiochiometry

A. Standard temperature and pressure (STP) 1. 0 °C, 273 K

2. 760 torr, 1 atm B. Molar volume

1. One mole of an ideal gas occupies 22.42 liters of volume at STP C. Things to remember

volume

mass

Density

=

mass

molar

ce

subs

of

grams

n

=

tan

5.5 Dalton's Law of Partial Pressures (John Dalton, 1803) A. Statement of law

1. "For a mixture of gases in a container, the total pressure exerted is the sum of the pressures each gas would exert if it were alone"

2. It is the total number of moles of particles that is important, not the identity or composition of the gas particles

B. Derivation 1.

_P

_P

_P

_P

V

T

R

n

P

V

T

R

n

P

V

T

R

n

P

_P

V

T

R

n

V

T

R

n

V

T

R

n

(

)

V

T

R

n

n

n

P

V

T

R

n

P

1. The ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture

2. For an ideal gas, the mole fraction (x):

P

P

n

n

x

5.6 The Kinetic Molecular Theory of Gases (KMT) A. Postulates of the KMT Related to Ideal Gases

1. The particles are so small compared with the distances between them that the volume of the individual particles can be assumed to be zero 2. The particles are in constant motion. Collisions of the particles with the

walls of the container cause pressure

3. Assume that the particles exert no forces on each other.

4. The average kinetic energy of a collection of gas particles is assumed to be directly proportional to the Kelvin temperature of the gas

B. Explaining Observed Behavior with KMT 1. P and V (T = constant)

a. As V is decreased, P increases:

V decrease causes a decrease in the surface area. Since P

is force/area, the decrease in V causes the area to decrease, increasing the P

2. P and T (V = constant)

a. As T increase, P increases

The increase in T causes an increase in average kinetic energy. Molecules moving faster collide with the walls of the container more frequently, and with greater force

3. V and T (P = constant)

a. As T increases, V also increases

Increased T creates more frequent, more forceful collisions.

V must increase proportionally to increase the surface area,

and maintain P 4. V and n (T and P constant)

a. As n increases, V must increase

Increasing the number of particles increases the number of collisions. This can be balanced by an increase in V to maintain constant P

5. Dalton's law of partial pressures

a. P is independent of the type of gas molecule

KMT states that particles are independent, and V is

assumed to be zero. The identity of the molecule is therefore unimportant

C. Root Mean Square Velocity

1. Velocity of a gas is dependent on mass and temperature. 2. Velocity of gases is determined as an average

a. M = mass of one mole of gas particles in kg b. R = 8.3145 J/K•mol (1) joule = kg•m2/s2

M

T

R

u

3

D. Mean Free Path

1. Average distance a molecule travels between collisions a. 1 x 10-7 m for O2 at STP

5.7 Effusion and Diffusion A. Effusion

1. Movement of a gas through a small opening into an evacuated container (vacuum)

2. Graham's law of effusion

2

1

gas

for

effusion

of

Rate

gas

for

effusion

of

Rate

M

M

1. The mixing of gases

2. Diffusion is complicated to describe theoretically and mathematically

5.8 Real Gases and van der Waals Equation (Johannes van der Waals, 1873) A. Volume

1. Real gas molecules do have volume

2. Volume available is not 100% of the container volume a. n = number of moles

b. b = is an empirical constant, derived from experimental results

Ideal Real

V

T

R

n

P

b

n

V

T

R

n

P

1. Molecules of real gases do experience attractive forces

a. a = proportionality constant determined by observation of the gas

















V

n

a

P

P

C. Combining to derive van der Waal's eqn

















V

n

a

b

n

V

T

R

n

P

and then rearranging…

(

V

n

b

)

n

R

T

V

n

a

P

















5.9 Chemistry in the Atmosphere

A. Composition of the Troposphere

Composition of dry air (sea level

Component Mole Fraction

Nitrogen 0.78084 Oxygen 0.20948 Argon 0.00934 Carbon dioxide 0.000345 Neon 0.00001818 Helium 0.00000524 Methane 0.00000168

B. Photochemical Smog - the problem of nitrogen oxides (NOx)

1. Auto exhaust contains small amounts of NO, which is quickly oxidized 2NO(g) + O2(g) à 2NO2(g)

2. Radiant energy causes NO2 to decompose

NO2(g) à NO(g) + O(g)

3. Free oxygen atoms combine with oxygen molecules to form ozone O(g) + O2(g) à O3(g)

4. Ozone may absorb light energy and decompose to excited oxygen atoms and excited oxygen molecules

O3(g) à O2* + O*

5. Excited oxygen atoms react with water to form the hydroxyl radical O* + H2O à 2OH

6. Hydroxyl can react with NO2 to form nitric acid

OH + NO2 à HNO3

C. Coal and acid rain

1. Most coal, especially cheap coal, contains sulfer S (in coal) + O2 (g) à SO2 (g)

2. Sulfur dioxide is oxidized in air

2SO2 (g) + O2(g) à 2SO3 (g)

3. Acid rain forms at the SO3 combines with water in the air SO3 (g) + H2O (l) à H2SO4 (aq)

AP Chemistry A. Allan

Chapter Six Notes - Thermochemistry 6.1 The Nature of Energy

A. Definition

1. Energy is the capacity to do work (or to produce heat*)

a. Work is a force acting over a distance (moving an object) b. *Heat is actually a form of energy.

(1) chemicals may store potential energy in their bonds that can be released as heat energy

B. Law of Conservation of Energy

1. Energy can be converted from one form to another, but cannot be created or destroyed

a. Potential energy

(1) energy due to position or composition b. Kinetic energy

(1) energy due to the motion of an object (2)

_KE

_m

_v

2 1

=

C. Heat and Temperature

1. Temperature reflects random motion of particles in a substance 2. Temperature indicates the direction in which heat energy will flow 3. Heat is a measure of energy content

4. Heat is what is transferred during a temperature change D. State Functions

1. A property of a system that depends only on its present state.

2. State functions do not depend on what has happened in the system, or what might happen in the system in the future

3. State functions are independent of the pathway taken to get to that state.

Example: a liter of water behind a dam has the same potential energy for work regardless of whether it flowed downhill to the dam, or was taken uphill to the dam in a bucket. The potential energy is a state function dependent only on the current position of the water, not on how the water got there.

E. Chemical Energy

1. Exothermic reactions

a. Reactions that give off energy as they progress

b. Some of the potential energy stored in the chemical bonds is converted to thermal energy (random KE) through heat c. Products are generally more stable (stronger bonds) than

reactants

2. Endothermic reactions

a. Reactions in which energy is absorbed from the surroundings b. Energy flows into the system to increase the potential energy of

c. Products are generally less stable (weaker bonds) than the reactants F. Thermodynamics 1. System Energy

w

q

E

∆

(1) q is positive in endothermic reactions (2) q is negative in exothermic reactions b. w = work

(1) w is negative if the system does work (2) w is positive if work is done on the system 2. Work done by gases

V

P

w

∆

a. by a gas (through expansion) (1) ∆V is positive

(2) w is negative b. to a gas (by compression)

(1) ∆V is negative (2) w is positive

6.2 Enthalpy and Calorimetry A. Enthalpy

PV

E

H

1. In systems at constant pressure, where the only work is PV, the change in enthalpy is due only to energy flow as heat (∆H = heat of rxn)

H

H

H

∆

a. ∆H is negative for exothermic rxns b. ∆H is positive for endothermic rxns

B. Calorimetry - science of measuring heat 1. Heat capacity (C)

a. ratio of heat absorbed to increase in temperature

increase

e

Temperatur

absorbed

heat

C

2. Specific Heat Capacity

a. Energy required to raise the temp of 1 gram of a substance by 1 °C

3. Molar heat capacity

a. Energy required to raise the temp of 1 mole of a substance by 1 °C

C. Constant Pressure Calorimetry (solutions) 1. Calculating Heat of Rxn, ∆H

a. ∆H = specific heat capacity x mass of sol'n x increase in temp

T

m

s

H

∆

∆

2. Heat of rxn is an extensive property - dependent on the amount of substance

a. ∆H α moles of reactant D. Constant Volume Calorimetry

1. Volume of bomb calorimeter cannot change, so no work is done

2. The heat capacity of the calorimeter must be known, generally in kJ/°C

2.

∆

_E

q

_w

_w

₀

∆

_E

q

A. Statement of Hess's Law

1. In going from a particular set of reactants to a particular set of

products, the change in enthalpy (∆H) is the same whether the reaction takes place in one step or in a series of steps

One step:

Two step

N2 (g) + O2 (g) à 2NO (g) ∆H2 = 180kJ

2NO (g) + O2 (g) à 2NO2 (g) ∆H3 = -112kJ

N2 (g) + 2O2 (g) à 2NO2 (g) ∆H2 + ∆H3 = 68kJ

B. Characteristics of Enthalpy Changes

1. If a reaction is reversed the sign on ∆H is reversed N2 (g) + 2O2 (g) à 2NO2 (g) ∆H = 68kJ

2NO2 (g) à N2 (g) + 2O2 (g) ∆H = - 68kJ

2. The magnitude of ∆H is directly proportional to the quantities of reactants and products in a reaction. If the coefficients in a balanced reaction are multiplied by an integer, the value of ∆H is multiplied by the same integer

C. Using Hess's Law

1. Work backward from the final reaction

2. Reverse reactions as needed, being sure to also reverse ∆H 3. Remember that identical substances found on both sides of the

summed equation cancel each other

6.4 Standard Enthalpies of Formation A. Standard State

1. For a compound a. Gaseous state

(1) pressure of 1 atm b. Pure liquid or solid

(1) standard state is the pure liquid or solid c. Substance in solution

(1) concentration of 1 M 2. For an element

a. the form in which the element exists at 1 atm and 25°C B. Standard Enthalpy of Formation (∆Hf°)

1. The change in enthalpy that accompanies the formation of one mole of a compound from its elements with all elements in their standard state C. Calculating enthalpy change

1. When a rxn is reversed, the magnitude of ∆H remains the same, but its sign changes

2. When the balanced eqn for a rxn is multiplied by an integer, the value of ∆H must be multiplied by the same integer

3. The change in enthalpy for a rxn can be calculated from the enthalpies of formation of the reactants and products

(

)

_n

_H

(

)

H

n

H

∑

∆

∑

∆

∆

4. Elements in their standard states are not included a. For elements in their standard state, ∆Hf° = 0

6.5 Present Sources of Energy A. Fossil Fuels

1. Energy derived from these fuels was initially captured from solar energy by photosynthesis

2. Combustion of fossil fuels always produces H2O and CO2

B. Petroleum and Natural Gas 1. Petroleum

a. Thick, dark liquid composed of hydrocarbon 2. Natural gas

a. Methane, with smaller amounts of ethane, propane and butane

Some Common Hydrocarbons

CH4 Methane C2H6 Ethane C3H8 Propane C4H10 Butane C5H12 Pentane C6H14 Hexane C7H16 Heptane C8H18 Octane C. Petroleum Refining

1. Original refining isolated kerosene (gasoline was a waste product) 2. Tetraethyl lead added as an "anti-knock" agent

Petroleum Fraction Major Uses

C5 - C10 Gasoline

C10 - C18 Kerosene, Jet fuel

C15 - C25 Diesel fuel, Heating oil, lubricating oil

> C25 Asphalt

D. Coal

1. Four stages of Coal

2. Carbon content increases over time

3. Value of coal is proportional to carbon content

Mass Percent of Each Element

Type of Coal C H O N S

Lignite 71 4 23 1 1

Subbituminous 77 5 16 1 1

Bituminous 80 6 8 1 5

Anthracite 92 3 3 1 1

E. CO2 and Earth's Climate

1. CO2 is a by-product of cellular respiration

2. CO2 is a by-product of burning fossil fuels

3. CO2 is a greenhouse gas

4. Atmospheric CO2 increased 16% from 1880 to 1980

6.6 New Energy Sources A. Coal Conversion

1. Gasification

a. Reduce length of hydrocarbon molecules to create liquid or gaseous fuels

b. Produce Syngas (CO and H2)

2. Coal Slurry

a. Coal dust suspended in water used as a heavy fuel oil replacement

3. Coal limitations

a. Mining of coal has a serious environmental impact B. Hydrogen as a fuel

1. Freeing hydrogen from compounds requires substantial energy a. CH4 (g) + H2O (g) à 3H2 (g) + CO (g) ∆H = 206 kJ

b. H2O (l) + H2 (g) + 1/2 O2 (g) ∆H = 286 kJ

2. Hydrogen is difficult to transport

a. Hydrogen in contact with metal produces free hydrogen atoms b. Hydrogen attempts penetrate the metal and make it brittle 3. Hydrogen is not dense

a. The fuel equivalent of 20 gallons of gasoline occupies a volume of 238,000 liters.

b. Liquid hydrogen is stored under great pressure and is potentially explosive

C. Other Energy Alternatives 1. Shale

a. Must be heated to extract fuel molecules, and produces immense amounts of waste rock

2. Ethanol from fermentation

a. Mixture of ethanol and gasoline - gasohol b. Ethanol is renewable

3. Methanol 4. Seed oils

AP Chemistry A. Allan

Chapter 7 Notes - Atomic Structure and Periodicity 7.1 Electromagnetic Radiation

A. Types of EM Radiation (wavelengths in meters)

10-12 10-10 10-8 4 to 7x10-7 10-4 10-2 1 102 104

Wavelength increases Frequency decreases Energy decreases

Speed is constant = 2.9979 x 108 m/sec B. Properties of EM Waves

1. Wavelength (λ)

a. Distance between two consecutive peaks or troughs in a wave b. Measured in meters (SI system)

2. Frequency (ν)

a. Number of waves that pass a given point per second b. Measured in hertz (sec-1)

3. Speed ( c )

a. Measured in meters/sec

4. Relationship of properties a.

λ

ν

_c

gamma xrays UV visible IR micro Radio waves

7.2 The Nature of Matter

A. Max Planck and Quantum Theory

1. Energy is gained or lost in whole number multiples of the quantity hv Frequency = v

Planck's constant = h = 6.626 x 10-34 J•S

∆

_E

_nhv

2. Energy is transferred to matter in packets of energy, each called a quantum

B. Einstein and the Particle Nature of Matter

1. EM radiation is a stream of particles - "photons"

λ

hc

hv

E

2. Energy and mass are inter-related

_E

_mc

1. Light travels through space as a wave 2. Light transmits energy as a particle

3. Particle's have wavelength, exhibited by diffraction patterns

mv

h

λ

a. large particles have very short wavelengths

b. All matter exhibits both particle and wave properties

7.3 The Atomic Spectrum of Hydrogen A. Continuous spectra

1. Contains all wavelengths of light B. Bright line spectra

410nm 434nm 486nm 656nm

1. Excited electrons in an atom return to lower energy states 2. Energy is emitted in the form of a photon of definite wavelength 3. Definite change in energy corresponds to:

a. Definite frequency b. Definite wavelength

λ

hc

hv

E

∆

7.4 Bohr Model (Neils Bohr, 1913) A. Quantum Model

1. The electron moves around the nucleus only in certain allowed circular orbits

2. Bright line spectra confirms that only certain energies exist in the atom, and atom emits photons with definite wavelengths when the electron returns to a lower energy state

3. Energy levels available to the electron in the hydrogen atom

_       − = −

n

Z

J

E

10

178

.

2

B. Calculating the energy of the emitted photon 1. Calculate electron energy in outer level 2. Calculate electron energy in inner level 3. Calculate the change in energy (∆E)

∆E = energy of final state - energy of initial state 4. Use the equation:

E

hc

∆

=

λ

to calculate the wavelength of the emitted photon C. Energy Change in Hydrogen atoms

1. Calculate energy change between any two energy levels

        − − = −

∆

n

n

J

E

₁₀

1

1

178

.

2

D. Shortcomings of the Bohr Model

1. Bohr's model does not work for atoms other than hydrogen 2. Electron's do not move in circular orbits

7.5 The Quantum Mechanical Model of the Atom A. The Electron as a standing wave

1. Standing waves do not propagate through space 2. Standing waves are fixed at both ends

3. Only certain size orbits can contain whole numbers of half wave lengths

a. fits the observation of fixed energy quantities - "quanta" B. The Schrödinger Equation (Erwin Schrödinger)

1. For the motion of one particle, in the along the x axis in space:

ψ

ψ

ψ

π

dx

V

E

d

m

h

8

2. Solution of the equation has demonstrated that E (energy) must occur in integer multiples

3. Your book has done you the favor of greatly simplifying this and presents the general equation:

ψ

ψ E

H

H

= a set of mathematical functions called an "operator"

ψ

4. Orbitals

1. Orbitals are not circular orbits for electron

C. Heisenberg Uncertainty Principle (Werner Heisenberg)

1. "There is a fundamental limitation on how precisely we can know both the position and momentum of a particle at a given time"

π

4

)

(

mv

h

x

∆

∆

∆

x

=

∆

(mv

)

accurately we can know its momentum, and vice-versa D. Physical Meaning of a Wave Function

1. Square of the absolute value of the wave function gives a probability distribution.

2

ψ

2. Electron density map indicates the most probable distance from the nucleus

3. Wave functions and probability maps do not describe: a. How an electron arrived at its location

b. Where the electron will go next

c. When the electron will be in a particular location

7.6 Quantum Numbers

A. Principal Quantum Number (n) 1. Integral values: 1, 2, 3, ….

2. Indicates probable distance from the nucleus a. Higher numbers = greater distance

b. Greater distance = less tightly bound = higher energy B. Angular Momentum Quantum (l)

(this was called the "orbital quantum number" in your general chem book) 1. Integral values from 0 to n - 1 for each principal quantum number n 2. Indicates the shape of the atomic orbitals

Table 7.1 Angular momentum quantum numbers and corresponding atomic orbital numbers

Value of l 0 1 2 3 4

Letter used s p d f g

C. Magnetic Quantum Number (ml)

1. Integral values from l to -l, including zero

2. Magnetic quantum number relates to the orientation of the orbital in space relative to the other orbitals

D. Spin Quantum Number 1. Covered in section 7.8

Table 7.2 Quantum numbers for the first four levels of orbitals in the hydrogen atom n l Orbital designation ml # of orbitals 1 0 1s 0 1 2 0 2s 0 1 1 2p -1, 0, 1 3 3 0 3s 0 1 1 3p -1, 0, 1 3 2 3d -2, -1, 0, 1, 2 5 4 0 4s 0 1 1 4p -1, 0, 1 3 2 4d -2, -1, 0, 1, 2 5 3 4f -3, -2, -1, 0, 1, 2, 3 7

7.7 Orbital Shapes and Energies A. Size of orbitals

1. Defined as the surface that contains 90% of the total electron probability

2. Orbitals of the same shape (s, for instance) grow larger as n increases B. s Orbitals

1. Spherical shape

2. Nodes (s orbitals of n=2 or greater) a. Internal regions of zero probability C. p Orbitals

1. Two lobes each

2. Occur in levels n=2 and greater

D. d Orbitals

1. Occur in levels n=3 and greater 2. Two fundamental shapes

a. Four orbitals with four lobes each, centered in the plane indicated in the orbital label

dxz dyz dxy dx2- y2

b. Fifth orbital is uniquely shaped - two lobes along the z axis and a belt centered in the xy plane

dz2

E. f Orbitals

1. Occur in levels n=4 and greater 2. Highly complex shapes

3. Not involved in bonding in most compounds F. Orbital Energies

1. All orbitals with the same value of n have the same energy a. "degenerate orbitals" (hydrogen only!)

2. The lowest energy state is called the "ground state"

3. When the atom absorbs energy, electrons may move to higher energy orbitals - "excited state"

7.8 Electron Spin and the Pauli Principle A. Electronic Spin Quantum Number

1. An orbital can hold only two electrons, and they must have opposite spins

2. Spin can have two values, +1/2 and -1/2 B. Pauli Exclusion Principle (Wolfgang Pauli)

1. "In a given atom no two electrons can have the same set of four quantum numbers"

7.9 Polyelectronic Atoms

A. Internal Atomic energies

1. Kinetic energy of moving electrons

2. Potential energy of attraction between nucleus and electrons 3. Potential energy of repulsion between electrons

B. The Electron Correlation Problem

1. Electron pathways are not known, so electron replusive forces cannot be calculated exactly

2. We approximate the average repulsions of all other electrons

C. Screening or Shielding

1. Electrons are attracted to the nucleus 2. Electrons are repulsed by other electrons

3. Electrons would be bound more tightly if other electrons weren't present

D. Variations in energy within the same quantum level

1. Atoms other than hydrogen have variations in energy for orbitals having the same principal quantum number

2. Electrons fill orbitals of the same n value in preferential order Ens < Enp < End < Enf

3. Electron density profiles show that s electrons penetrate to the nucleus more than other orbital types

a. Closer proximity to the nucleus = lower energy

7.10 The History of the Periodic Table

No notes. Read this section. Notice especially Table 7.3 and 7.4. The periodic table is as useful for predicting properties of undiscovered elements today as it was in Mendeleev's time

7.11 The Aufbau Principle and the Periodic Table A. The Aufbau Principle

1. "As protons are added one by one to the nucleus to build up elements, electrons are similarly added to these hydrogen-like orbitals

B. Hund's Rule

1. "The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals

***Note: We will review configuration and orbital notation in class. This was covered extensively in General Chem, and will come back to you quickly

C. Period Table Vocabulary 1. Valence Electrons

a. Electrons in the outermost principal quantum level of an atom b. Elements in the same group (vertical column) have the same

valence electron configuration 2. Transition metals

a. What we have called the "d" block 3. Lanthanide and Actinide Series

a. The sets of 14 elements following lanthanum and actinium b. What we have called the "f" block

4. Main-group, or Representative Elements a. Groups 1A through 8A

b. Configurations are consistent 5. Metalloids (semi-metals)

a. Found along the border between metals and nonmetals b. Exhibit properties of metals and nonmetals

7.12 Periodic Trends in Atomic Properties

A. Ionization Energy - the energy required to remove an electron from an atom 1. Ionization energy increases for successive electrons

2. Ionization energy tends to increase across a period

a. electrons in the same quantum level do not shield as effectively as electrons in inner levels

b. irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove

3. Ionization energy decreases with increasing atomic number within a group

a. electrons farther from the nucleus are easier to remove B. Electron Affinity - the energy change associated with the addition of an

electron

1. Affinity tends to increase across a period

2. Affinity tends to decrease as you go down in a period

a. electrons farther from the nucleus experience less nuclear attraction

b. Some irregularities due to repulsive forces in the relatively small p orbitals

C. Atomic Radius

1. Determination of radius

a. half of the distance between radii in a covalently bonded diatomic molecule - "covalent atomic radii"

2r

2. Periodic Trends

a. Radius decreases across a period

(1) increased effective nuclear charge due to decreased shielding

b. Radius increases down a group

(1) addition of principal quantum levels

7.13 The Properties of a Group: the Alkali Metals

A. Easily lose valence electron (Reducing agents) 1. React with halogens to form salts

2. React violently with water

a. Lithium is not the most reactive because the heat of reaction is insufficient to melt lithium and expose all of its surface area B. Large hydration energy

1. Positive ionic charge makes ions attractive to polar water molecules C. Radius and Ionization energy follow expected trends

AP Chemistry A. Allan

Chapter 8 Notes - Bonding: General Concepts 8.1 Types of Chemical Bonds

A. Ionic Bonding

1. Electrons are transferred 2. Metals react with nonmetals

3. Ions paired have lower energy (greater stability) than separated ions B. Coulomb's Law 1. _       ⋅ = −

r

Q

Q

E

b. Q1 and Q2 are numerical ion charges

c. r = distance between ion center in nanometers d. negative sign indicates an attractive force C. Bond Length (covalent)

1. Distance at which the system energy is at a minimum 2. Forces at work

a. Attractive forces (proton - electron)

b. Repulsive forces (electron - electron, proton - proton)

3. Energy is given off (bond energy) when two atoms achieve greater stability together than apart

D. Covalent Bonds

1. Electrons are shared by nuclei 2. Pure covalent (non-polar covalent)

a. Electrons are shared evenly 3. Polar covalent bonds

a. Electrons are shared unequally b. Atoms end up with fractional charges

(1) δ+ or δ -8.2 Electronegativity

A. Electronegativity

1. The ability of an atom in a molecule to attract shared electrons to itself B. Electronegativity Trends

1. Electronegativity generally increases across a period (why?) 2. Electronegativity generally decrease within a family (why?) C. Characterizing bonds

1. Greater electronegativity difference between two elements means less covalent character and greater ionic character

2. We will not use the subtraction of electronegativities to determine ionic character. This text uses a practical definition to identify ionic

compounds:

Any compound that conducts an electric current when melted is an ionic compound.

8.3 Bond Polarity and Dipole Moments A. Dipolar Molecules

1. Molecules with a somewhat negative end and a somewhat positive end (a dipole moment)

2. Molecules with preferential orientation in an electric field + + +

-3. All diatomic molecules with a polar covalent bond are dipolar B. Molecules with Polar Bonds but no Dipole Moment

1. Linear, radial or tetrahedral symmetry of charge distribution a. CO2 - linear

b. CCl4 - tetrahedral

2. See table 8.2 in your text

8.4 Ions: Electron Configurations and Sizes

A. Bonding and Noble Gas Electron Configurations 1. Ionic bonds

a. Electrons are transferred until each species attains a noble gas electron configuration

2. Covalent bonds

a. Electrons are shared in order to complete the valence configurations of both atoms

B. Predicting Formulas of Ionic Compounds

1. Placement of elements on the periodic table suggests how many electrons are lost or gained to achieve a noble-gas configuration

a. Group I loses one electron, Group II loses two, Group VI gains two, Group VII gains one….

2. Formulas for compounds are balanced so that the total positive ionic charge is equal to the total negative ionic charge

O

Al

1. Anions are larger than the parent atom 2. Cations are smaller than the parent atom 3. Ion size increases within a family

4. Isoelectronic ions

a. Ions with the same number of electrons

8.5 Formation of Binary Ionic Compounds A. Lattice Energy

1. The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

M+ (g) + X- (g) à MX (s)

2. Energy change is exothermic (negative sign)

Example: Formation of lithium fluoride

Process Description Energy Change (kJ)

Li(s) à Li(g) Sublimation energy 161 Li(g) à Li+

(g) + e- Ionization energy 520 1/2F2 à F(g) Bond energy (1/2 mole) 77

F(g) + e- à F-(g) Electron affinity -328 Li+(g) + F-(g) à LiF(s) Lattice energy -1047 Li(s) + 1/2F2(g) à LiF(s) ∆H -617

3. The formation of ionic compounds is endothermic until the formation of the lattice

4. The lattice formed by alkali metals and halogens (1:1 ratio) is cubic except for cesium salts

B. Lattice Energy Calculations

1. _       =

r

Q

Q

k

a. k = a proportionality constant dependent on the solid structure and the electron configuration

b. Q1 and Q2 are charges on the ions

c. r = shortest distance between centers of the cations and the anions

2. Lattice energy increases as the ionic charge increases and the distance between anions and cations decreases

8.6 Partial Ionic Character of Covalent Bonds A. Calculating Percent Ionic Character

% 100 x Y X of moment dipole calculated Y X of moment dipole measured character ionic Percent     − = _{+ −} B. Ionic vs. Covalent

1. Ionic compounds generally have greater than 50% ionic character 2. Ionic compounds generally have electronegativity differences greater

than 1.6

3. Percent ionic character is difficult to calculate for compounds containing polyatomic ions

8.7 The Covalent Chemical Bond: A Model A. Strengths of the Bond Model

1. Associates quantities of energy with the formation of bonds between elements

2. Allows the drawing of structures showing the spatial relationship between atoms in a molecule

3. Provides a visual tool to understanding chemical structure B. Weaknesses of the Bond Model

1. Bonds are not actual physical structures

2. Bonds can not adequately explain some phenomena a. resonance

8.8 Covalent Bond Energies and Chemical Reactions A. Average Bond Energies

Process Energy Required (kJ/mol)

CH4(g) à CH3(g) + H(g) 435 CH3(g) à CH2(g) + H(g) 453 CH2(g) à CH(g) + H(g) 425 CH(g) à C(g) + H(g) 339 Total 1652 Average 413 B. Multiple Bonds

1. Single bonds - 1 pair of shared electrons 2. Double bonds - 2 pairs of shared electrons 3. Triple bonds - 3 pairs of shared electrons

Multiple Bonds, Average Energy (kJ/mole)

C=C 614 N=O 607

C≡C 839 N=N 418

O=O 495 N≡N 941

C=O 745 C≡N 891

C≡O 1072 C=N 615

4. As the number of shared electrons increases, the bond length shortens (see table 8.5)

C. Bond Energy and Enthalpy (using bond energy to calculate approximate energies for rxns)

1. ∆H = sum of the energies required to break old bonds(endothermic) +

sum of the energies released in forming new bonds (exothermic)

2. ∆H =

∑

∑

8.9 The Localized Electron Bonding Model A. Lone electron pairs

1. Electrons localized on an atom (unshared) B. Bonding electron pairs

1. Electrons found in the space between atoms (shared pairs) C. Localized Electron Model

1. "A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms

D. Derivations of the Localized Model

1. Valence electron arrangement using Lewis structures

2. Prediction of molecular geometry using VSEPR (valence shell electron pair repulsion)

3. Description of the type of atomic orbitals used to share or hold lone pairs of electrons

8.10 Lewis Structures

A. Electrons and Stability

1. "the most important requirement for the formation of a stable compound is that the atoms achieve noble gas configurations 2. Duet rule

a. Hydrogen, lithium, beryllium, and boron form stable molecules when they share two electrons (helium configuration)

3. Octet Rule

a. Elements carbon and beyond form stable molecules when they are surrounded by eight electrons

B. Writing Lewis Structures 1. Rules

a. Add up the TOTAL number of valence electrons from all atoms b. Use a pair of electrons to form a bond between each pair of

bound atoms. Lines instead of dots are used to indicate each pair of bonding electrons

c. Arrange the remaining atoms to satisfy the duet rule for hydrogen and the octet rule for the second row elements