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OSU / Chemistry / CHEM 1210 / When does liquid expand?

When does liquid expand?

When does liquid expand?


School: Ohio State University
Department: Chemistry
Course: Basics of Chemistry
Term: Fall 2016
Tags: General Chemistry and Chemistry
Cost: 50
Name: Chemistry 1210 Final Exam Study Guide
Description: This study guide covers everything that was gone over in lecture for the entire semester and chapters 1-10 of the textbook. The guide is very long but it contains all of the important concept information for all 10 of the covered chapters. Important formulas are included in their respective chapters.
Uploaded: 12/07/2016
38 Pages 66 Views 2 Unlocks

Chemistry 1210 Final Exam Study Guide

Liquid, expands a little when?


Chapter 1: Matter and Measurement

 A. Chemistry- study of matter and how it changes; deals with the  composition, structure, and reactions of matter

 a. Mass- the measure of the quantity of matter

 b. Weight- a result of the gravitational attraction between matter; a  force but is often used interchangeably with “mass”

 c. Atoms- “building blocks” of matter

 d. Molecules- two or more atoms that have joined together e. Composition  

i. The kinds of atoms that matter contains

ii. Can be at the macroscopic (can be seen and handled) or  submicroscopic (atoms and molecules) level

Compounds is a substances made up of, how many elements?

iii. Qualitative (what substances it is made of) or quantitative (the number of substances it is made of) Don't forget about the age old question of What are the four kinds of luck?
If you want to learn more check out What is the film theory?

f. Structure

i. Arrangement; how atoms are put together

ii. Makes a difference in what the final product is

1. EX: a sandwich layered in the usual order is different  

than a sandwich layered with the lettuce and cheese  

on the outside and the bread on the inside

g. Reactions

i. Occurs when the composition and structure changes We also discuss several other topics like What are the major parties in the history of the us?

ii. Exothermic (gives off heat) and endothermic (absorbs  


 B. Matter- anything that has mass and occupies space; physical material a. States of matter

What are the properties of matter?

i. Gas (also known as vapor)

1. No definite volume or shape

2. Takes the shape of its container and expands to fill it

3. Very compressible

4. Expands greatly when heated

ii. Liquid

1. Unchanging volume

2. Takes the shape of its container

3. A little compressible

4. Expands a little when heated If you want to learn more check out What is the purpose of a structuralist theory?

iii. Solid

1. Definite volume and shape

2. Not compressible

3. Barely expands when heated

 b. Pure substance- composition does not change between samples  used; definite properties

i. Can only be separated by chemical methods

 c. Elements- the simplest substances that can not be broken down  any further; consists of atoms that all contain the same number  of protons If you want to learn more check out How genes relate to crime?

i. 118 known elements

d. Compounds- substances made up of 2 or more elements i. chemically combined

ii. can be separated through chemical reactions

iii. ionic compounds

iv. molecular compounds

v. Law of Definite Proportions/ Law of Constant  


1. The elements in a compound are always combined in the same way

e. Mixtures- 2 or more substances physically, not chemically,  combined

i. Each substance retains its original chemical identity

ii. Composition can vary

iii. Homogenous

1. Solutions (can be solid, liquid, or gas)

2. Mixture is uniform throughout

3. Miscible- mix in all proportions If you want to learn more check out Who is edward taylor?

iv. Heterogenous

1. Mixture is NOT uniform throughout

2. Parts vary throughout the mixture

C. Properties of Matter

a. Physical properties

i. Observed without changing the substance

ii. Physical appearance changes but its composition does not iii. EX: color, odor, density, specific heat, physical state,  melting point, boiling point...evaporating water

b. Chemical properties

i. The way a substance reacts with another substance or  transforms into a chemically different substance

ii. EX: flammability…hydrogen burning in air

c. Intensive properties

i. Do NOT depend on the amount of a substance being used ii. Very useful in identifying substances

iii. EX: color, melting point, boiling point, density, specific  heat, temperature

d. Extensive properties

i. Depends on the amount of a substance being used

ii. EX: mass, volume, heat content

e. Physical change

i. Physical appearance changes but its composition does not ii. EX: change in state

1. Fusion- solid to liquid

2. Freezing- liquid to solid

3. Condensation- gas to liquid

4. Vaporization- liquid to gas

5. Sublimation- solid to gas

6. Deposition- gas to solid

f. Chemical change

i. Also called chemical reaction

ii. Substance transforms into a chemically difference  


iii. Composition and/or structure changes

g. Separation of mixtures

i. Filtration- filtering a substance through a filter

1. Used for heterogenous mixtures

ii. Distillation- utilizes the different abilities of substances to form gases

1. Used for separating a homogenous mixture

2. EX: boiling salt water…the water evaporates and the  salt is left behind

iii. Extraction- separation based on solubility differences iv. Chromatography- utilizes the different abilities of  

substances to adhere to the surface of solids

D. Units of Measurement

a. Prefixes

i. Kilo (k)… 10^3 

ii. Deci (d)… 10^-1 

iii. Centi (c)… 10^-2 

iv. Milli (m)… 10^-3 

v. Micro (µ)… 10^-6 

vi. Nano (n)… 10^-9 

b. Mass = gram

c. Length = meter

i. 1 inch = 2.54cm 

d. Volume = m^3, liter (L), milliliter (mL)

i. 1dm = 10cm

ii. 1mL = 1cm^3

e. Temperature

i. 212° F = 100°C 

1. boiling point of water

ii. 98.6°F = 37°C 

1. body temperature

iii. 32°F = 0°C 

1. freezing point of water

iv. Conversion formulas

1. y° C = 5°C/9°F (x° F – 32°F) 

2. y° F = 9°F/5°C (x° C) + 32°F

3. 0° C = 273.15 K 

v. 0 K is the lowest possible temperature

f. Density  

i. D = m/V 

ii. Solids = g/cm^3

iii. Liquids = g/mL

iv. Gases = g/L

v. Specific gravity  

1. Density of substance (g/mL) / density of water (g/mL) = specific gravity

2. No units

3. Density of water = 1.0 g/mL

g. Precision- how close each measurement is to another i. Standard deviation- how much the individual  

measurements differ from the average


h. Accuracy- how close a measurement is to the correct value E. Significant Figures

a. When recording from an instrument, report all the digits that you  know PLUS one that you estimate 

b. Rules of Significant figures

i. All nonzero digits are significant

ii. Zeros between significant digits are significant 1. Captive Zeros

2. EX: 20.006 has 5 significant figures

iii. Zeros to the left of the first nonzero digit are NOT  significant

1. Leading zeros

2. Make sure to locate the decimal point

3. EX: 0.00004 has 1 significant figure

iv. Zeros to the right of the last nonzero digit  

1. The number ends in zero to the RIGHT OF THE  

DECIMAL POINT, zeros are significant

a. Trailing zeros

b. EX: 0.0040 has 2 significant figures

c. EX: 4000.0 has 5 significant figures

2. The number ends in zero to the LEFT OF THE  

DECIMAL POINT, zeros are NOT significant

a. Trailing zeros

b. EX: 4000 has 1 significant figure

c. EX: 4120000 has 3 significant figures

c. Scientific notation

i. 1 nonzero digit to the left of the decimal point

ii. EX: 400 = 4 x 10^2 (1 significant figure) OR 4.0 x 10^2 (2  significant figures) OR 4.00 x 10^2 (3 significant figures) d. Calculations

i. The result of a calculation has to reflect the significant  

figures of the original measurements

ii. Multiplication and Division

1. The answer must possess the same number of  

significant figures as the original number in  

the problem that had the LEAST amount of  

significant figures

a. EX: If dividing/multiplying 907.54 with 43.9,  

your answer must have 3 significant figures  

because 43.9 has 3 significant figures, which is  

less than 907.54 which has 5 significant figures

2. Rounding Rule

a. If the leftmost number that is going to be left  

out because of rounding is less than 5, round  

down (the last number that will be remaining  

doesn’t change)

iii. Addition and Subtraction

1. The last place in your answer will be the last  

place that is common to all of the numbers

a. EX: 5 


 + 97 3.489 


The correct answer is 987 because the last digit  

common to all of the numbers is the one right before  

the decimal point, so that has to be the digit where  

you answer stops


 2. Rounding Rule  

a. If the leftmost number to be discarded is more  

than or equal to 5 and is followed by nonzero  

digits, round up

 3. Rounding Rule

a. if the number that is going to be left out is 5 or 5  

followed by zeros, round even (leave the last digit  

unchanged if it is an even number, or increase it by 1

if it is an odd number)

F. Dimensional Analysis (Factor Unit Method)

a. Conversion factor- a number that has 2 or more units associated  with it

i. Equal to 1

ii. Are exact by definition

b. To get the same number but in a different unit, multiply the given information of one type of unit by a conversion factor  

c. EX: Convert 1.54cm to meters

i. 1cm = 10^-2 m (this is your conversion factor)

ii. 1.54 cm x 10^-2 m = 0.154 m = 0.2 m

 1 1 cm

Chapter 2: Atoms, Molecules, and Ions

A. The Atomic Theory of Matter

a. Dalton’s Atomic Theory

i. Each element is composed of extremely small particles  called atoms

ii. All atoms of a given element are identical, but the atoms of one element are different from the atoms of all other  


iii. Atoms of one element cannot be changed into atoms of a  different element by chemical reactions

1. Atoms are neither created nor destroyed in chemical  


2. Law of Conservation of Mass 

iv. Compounds are formed when atoms of more than one  element combine

1. A given compound always has the same relative  

number and kinds of atoms

a. Law of Constant Composition 

2. Atoms combine in whole numbers

a. Law of Multiple Proportions 

b. Dalton’s Atomic Theory explains the law of constant composition and the law of conservation of mass, and deduced the law of  multiple proportions

B. The Discovery of Atomic Structure

a. Atoms

i. Atomic mass unit (amu) = 1.6603 x 10^-24 g  

ii. Mass of atoms: 1 – 260 amu

iii. Radii

1. Use nm

iv. 1 Angstrom = 10^-10 m = 10^-8 cm

b. Subatomic particles

i. Make up atoms

ii. Electron  

1. Charge = -1.6022 x 10^-19 C; relative = -1 

2. m = 9.1094 x 10^-28 g = 5.486 x 10^-4 amu

3. atoms are neutral until electrons are gained or lost

4. number of electrons in an atom = number of protons  

in the atom

5. electrons are in constant motion

6. Can be traded back and forth between atoms

iii. Proton

1. Charge = 1.6022 x 10^-19 C, relative = +1

2. m = 1.6726 x 10^-24 g = 1.0073 amu

3. make up about ½ of mass of an atom

4. number of protons in an atom of an element is that  

element’s atomic number 

5. The number of protons do not change in a normal  

chemical reaction

iv. Neutron

1. charge = 0 

2. m = 1.6749 x 10^-24 g = 1.0088 amu

3. make up about ½ of mass of an atom

v. Particles of the same charge repel each other; particles of  different charges attract each other

c. Cathode rays and Electrons

i. The radiation produced when a high voltage is applied to  electrodes in a tube

ii. Cause certain materials to give off light (fluoresce)

iii. J.J. Thompson

1. Discovered the electron by observing cathode rays  

as streams of negatively charged particles

2. Measured charge to mass ratio (coulombs)

iv. Robert Millikan- Oil Drop Experiment

1. Determined the charge of an electron: 1.6022 x 10^- 19

 d. Radioactivity- the spontaneous emission of radiation

i. Ernest Rutherford found 3 types of radiation with his gold  foil experiment

ii. Alpha

1. Bent by an electric field

2. Fast moving particles

3. Have positive charge of +2; attracted to a negative  


iii. Beta

1. Bent by an electric field

2. High speed electrons

3. Radioactive equivalent of cathode rays

4. Have a charge of -1; attracted to a positive charge

iv. Gamma

1. Unaffected by an electric field

e. The Nuclear Model of the Atom

i. Atoms contain a dense nucleus

1. Nucleus composed of protons and neutrons

ii. Nucleus surrounded by electrons in mostly empty space 1. Electrons are in a constant motion

iii. Atoms’ diameter = 10^-8 cm

C. The Modern View of Atomic Structure

a. Atomic Number (Z)

i. Number of protons in an atom

ii. Periodic table is arranged by increasing atomic number b. Mass Number (A)

i. A = number of protons + number of neutrons 

ii. Is an integer

c. Elemental/Atomic Symbol

i. Represents an element

ii. EX: Zn = zinc

iii. Number in top left corner is the element’s mass number iv. Number in the bottom left corner is the element’s atomic  number

d. Isotopes

i. Atoms of the same element that have the same number of  protons but a different number of neutrons 

ii. Have the same atomic number, but different mass number iii. Some mass gets lost in isotopic masses because it gets  converted to energy when the atoms come together

e. Atomic Weight

i. Weighted average of the isotopes of an atom

ii. Given in amu

iii. The average atomic mass 

1. Doesn’t give you the actual weight of the atom, but  

gives you an average

D. Periodic Table

a. The arrangement of elements in order of increasing atomic  number

i. Modern periodic law

b. Places elements with similar properties in vertical columns c. Groups or families

i. Vertical columns

ii. Elements within a group share similar properties 

iii. Columns labeled at the top by either numbers, letters, or  roman numerals

iv. Representative Elements (Main group elements)

1. The first 2 columns and the last 6 columns

2. 1A- 8A

3. 1A = alkali metals (except for hydrogen)

4. 2A = alkaline earth metals

5. 7A = halogens

6. 8A = noble gases

v. Transition Metals

1. The middle 10 columns

2. 1B- 8B

d. Periods

i. Horizontal rows

ii. The 2 long rows below the main body of the periodic table  are inner transition elements

1. Lanthanides – part of row 6

2. Actinides - part of row 7

e. The stair step semimetals are transitional metals

E. Metals vs Nonmetals

a. Metals

i. Solids

1. Except Hg

ii. Metallic luster

iii. Malleable and ductile

1. Can be pounded into shapes and drawn out into a  


iv. Conduct heat and electricity well

v. Oxides (compounds):

1. Nonvolatile (ionic compounds)

2. High melting

3. MgO, Na2O

b. Nonmetals

i. Gases or solids

1. Except Br

ii. Variety of color and appearance

iii. Solids are brittle

iv. Not good conductors

1. Insulators

v. Oxides:

1. Volatile

2. Low melting

3. CO, CO2, SO2

F. Molecules and Molecular Compounds

 a. Molecular formula 

i. Chemical formula that indicates the actual numbers of  atoms in a molecule

ii. Subscripts are always some integer multiple of the  

subscripts in the empirical formula

 b. Empirical formula 

i. Chemical formula that gives only the relative number of  atoms of each type in a molecule

ii. The smallest whole number ratio of atoms

c. Molecular elements

i. Diatomic molecule- a molecule made up of 2 atoms

1. EX: H2… has 2 H atoms bonded together

2. All of the halogens are diatomic

 ii. Polyatomics 

1. A molecule made up of more than 2 atoms

2. EX: P4, S8, O3

d. Molecular Compounds

i. Compounds composed of molecules containing more than  one type of an atom (contain atoms of different elements) ii. EX: CO2, H2O

iii. All elements are nonmetals

e. Structural formula

i. Shows which atoms are attached to which

ii. Demonstrates structure

G. Ions and Ionic Compounds

a. Ion- a charged particle that is formed when electrons are added  or removed from an atom

i. Cation- positively charged ion

1. Metals tend to lose electrons and become cations

ii. Anion- negatively charged ion

1. Nonmetals tend to gain electrons and become anions b. Polyatomic ion- atoms joined as a molecule but have a net  charge

i. NH^4+, SO4^2-

c. Predicting ionic charges

i. Many atoms gain/lose electrons to end up with the same  number of electrons as the noble gas closest to them in the periodic table

1. Representative elements (1A- 8A)

ii. Elements in the same isoelectronic series have the same  electron configuration

iii. Periodic Table

1. Cation groups

a. Charge = group number

i. Except for H

b. Special cations:

i. Al = +3

ii. Zn = +2

iii. Ag = +1

2. Anion groups

a. Charge = group number – 8

3. 5A = -3, 6A = -2, 7A = -1

iv. Ionic Compounds

1. Compound made of cations and anions

a. Oppositely charged ions are held together by  

electrostatic attractions

2. Compounds of metals and nonmetals

a. Electrically neurtral

3. Crystalline solids (salts)

4. EX: Ca^2+ and CO3^2- form CaCO3 (calcium  


5. Compounds are NOT molecules

a. They are formula units 

6. Empirical formula- chemical formula that shows the  simplest ratio of ions in a compound

H. Naming Inorganic Compounds

a. Organic compounds contain carbon and hydrogen, all other  compounds are considered inorganic

b. Chemical nomenclature- system of naming substances c. Monatomic ions

i. Cations

1. Formed from metal atoms: name of the elements  

followed by “ion”

2. EX: K^+ is potassium ion

ii. Anions

1. Add “-ide ion” to the end of the element’s name

2. EX: Br^- is bromide ion

d. Stock System

i. Metals that can have more than 1 possible charge

1. Transition and representative metals

ii. Use roman numerals in parenthesis after the element’s  name to represent the ion’s charge

iii. EX: Fe ^3+ …. Iron (III); Sn ^4+ …. Tin (IV)

e. Polyatomic Ions

i. Group of chemically bonded atoms with an overall charge ii. Cations

1. EX: NH4 ^+ ammonium ion

2. H3O^+ hydronium ion

3. Hg2^2+ mercury (I) ion

iii. Anions

1. Ending in “-ide”

a. EX: OH^- hydroxide ion

2. Others

a. EX: C2H3O2^- acetate ion

iv. Oxyanions- polyatomic anions that contain oxygen

1. Suffixes

a. “-ate” (most common)  

b. “-ite” (same charge but 1 fewer O atom)

2. Prefixes

a. “per-“

i. 1 more O atom than “-ate”

b. “hypo-“

i. 1 less O atom than “-ite”

3. EX:

a. ClO4^- perchlorate

b. ClO3^- chlorate

c. ClO2^- chlorite

d. ClO^- hypochlorite

e. Cl^- chloride  

f. ***Overall charge remains the same

4. Addition of H^+ to a -2 or -3 oxyanion

a. Results still charged (anions)

b. EX:

i. HCO3^- bicarbonate/ hydrogen  


ii. HPO4^2- monohydrogen phosphate  

1. Mono- is usually dropped

iii. H2PO4^- dihydrogen phosphate

5. Acids

a. Acid- H^+ produces a neutral compound when  

combined with an anion

b. Not ionic

i. But ionize in H2O to produce H^+  


c. EX: HCl (g) H2O H^+ (aq) + Cl^- (aq)

d. Binary (2 element) acids  

i. Hydrogen + nonmetal

1. –ide  hydro- -ic acid

a. EX: HF (aq) hydrofluoric  


2. Per- -ate  per- -ic acid

3. –ate  -ic acid

4. –ite  -ous acid

5. hypo- -ite  hypo- -ous acid

I. Binary Molecular Compounds

a. Compounds contain 2 different elements

i. Nonmetals or nonmetals and semimetals 

b. Element further left in the periodic table (closest to the metals) is usually listed first

i. Exception when the compound contains oxygen and any  halogen (except for fluorine), then oxygen is always written last

c. If both elements are in the same group, the lower one is listed  first (less electronegative)

d. The name of the second element is given an –ide ending  e. Greek prefixes are used to indicate the number of atoms of each  element

i. Mono- is never used with the first element

1. When the prefix ends with an a or an o and the  

second element begins with a vowel, the prefix is  


ii. Prefixes:

1. Mono

2. Di-

3. Tri

4. Tetra

5. Penta

6. Hexa

7. Hepta

8. Octa

9. Nona

10. Deca

iii. EX: N2O4 Dinitrogen tetroxide

iv. SO3 Sulfur trioxide

J. Simple Organic Compounds

a. Hydrocarbons- compounds that contain only carbon and  hydrogen

b. Alk anes- simplest class of hydrocarbons

i. Each carbon is bonded to 4 other atoms

ii. Methane CH4

iii. Ethane C2H6

iv. Propane C3H8

c. Functional groups

i. Hydrogens are replaced with functional groups- specific  groups of atoms

ii. Alkenes

1. Ethene CH2 (double bond) CH2

iii. Alkynes

1. Ethyne CH (triple bond) CH

d. Alcohols

i. Replace hydrogen atom with an –OH group

ii. Methanol, ethanol, propanol

1. Are structural isomers of each other (have same  

chemical formula but different arrangements of  


e. Ethers

i. Dimethyl ether, diethyl ether

ii. Functional group isomers with alcohols

f. Aldehydes

i. C (double bond) O group … carbonyl

ii. EX: HCHO formaldehyde

g. Ketones

i. EX: CH3COCH3 acetone

h. Carboxylic acids

i. Carboxyl group

ii. EX: H-CO2H formic acid

iii. CH3CO2H acetic acid

i. Amines

i. organic analogues of NH3 (ammonia)

ii. EX: CH3NH2 methyl amine


Chapter 3: Stoichiometry: Calculations with Chemical Formulas and  Equations

 A. Chemical Equations- symbolic representation of a chemical reaction  a. Coefficients- indicate the number of atoms, molecules, or formula units of each substance that is involved in a reaction

b. Each substance always written as:

i. (s) = solid

1. most ionic substances are solids at room  


2. most elements are solids at room temperature

ii. (g) = gas

iii. (l) = liquid

iv. (aq) = aqueous solution (dissolved in water)

c. reactants- chemical formulas to the left of the arrow

d. products- chemical formulas to the right of the arrow

i. substances that were produced from the reaction

e. Balancing Equations

i. Final equation should contain the smallest possible whole number coefficients

ii. Never change subscripts, instead use a coefficient  

iii. The number of atoms of each element before the reaction  has to equal the number of atoms of each element after  

the reaction

iv. Charges on the left side of the equation must equal the  charges on the right side of the equation

v. Tips:

1. Balance the molecule that has the largest number of  

atoms of a single element first

a. Excluding H, O, and polyatomic ions

2. Balance polyatomic ions as a whole unit (if they did  

not change)

3. Balance hydrogen and oxygen last

a. If they are both in an equation, balance the one

that is in an equal number of compounds on  

both sides

i. If they both are, balance the one that is  

in the least number of compounds on  

both sides

B. Simple Patterns of Chemical Reactivity  

a. Elements of the same group often react in similar ways

b. Combination reaction- 2 or more substances react to form one  product

i. EX: 2Mg (s) + O2 (g)  2MgO (s)

ii. Oxides reacting with water:

1. Metal oxides

a. Basic oxides produce basic metal hydroxides

b. EX: K2O + H2O  2KOH

c. OH- causes substances to be basic

2. Nonmetal oxides

a. Acidic oxides produce acids

b. EX: SO3 + H2O  H2SO4

c. Decomposition reaction- one substance undergoes a reaction to  produce 2 or more other substances

i. EX: CaCO3 (s)  CaO (s) + CO2 (g)

d. Combustion reaction- a rapid reaction that produces a flame; A  reaction with O2

i. Will always only produce CO2 and H2O

ii. Incomplete combustion

1. Occurs when there is not enough oxygen present in  

the reaction

iii. Oxide reaction- a combustion reaction that requires  intermediate steps

C. Formula Weights

a. Molecular weight- the sum of the atomic weights of the atoms in  the chemical formula of a molecule

b. Formula weight- the sum of the atomic weights of the atoms in  the chemical formula of an ionic substance

c. Percent composition- the percentage by mass contributed by  each element in a substance

i. = (number of atoms of element)(atomic weight) x  100

 formula weight of substance

D. Avogadro’s Number and The Mole

a. Avogadro’s number = 6.022 x 1023 atoms

i. Represented by NA

b. 1 amu = 6.022 x 1023 grams

c. Molar Mass

i. 1 mole = 6.022 x 1023 particles

1. a mole is always the same number, but different  

samples of 1 mol can have different masses

ii. an element’s atomic weight/molecular weight/formula  weight (amu) is equal to the mass (g) of 1 mole of that  


d. Interconverting masses, moles, and number of particles i. Grams ( use molar mass)  moles (use Avogadro’s number)  formula units

ii. Formula units (use Avogadro’s number)  moles (use molar mass)  grams

E. Empirical Formulas from Analyses

a. The ratio of the numbers of moles of all elements in a compound  gives the subscripts in the compound’s empirical formula b. Calculating the mole ratio of each element in a compound i. Mass % elements  (assume 100g sample) grams of each  element  (use molar mass) moles of each element 

(calculate the mole ratio)  then you can find empirical  


c. Finding empirical formula

i. Divide larger number of moles by the smaller to obtain the  mole ratio (also the atom ratio)

ii. Multiply by the simplest factor to get whole numbers iii. Write formula

d. Molecular Formulas from Empirical Formulas

i. The subscripts in the molecular formula of a substance are  always whole number multiples of the subscripts in its  

empirical formula

ii. Whole number multiple = molecular weight 

 Empirical formula weight e. Combustion analysis  

i. Determines empirical formula for compounds principally  containing C and H

Chapter 4 Notes: Reactions in Aqueous Solution


I. General Properties of Aqueous Solutions

a. Homogenous mixture; solution

b. 2 components: solvent and solute

c. dissolution- process of dissolving

d. Nonelectrolytes

i. dissolve in molecular form

ii. do NOT conduct electricity

iii. EX: pure water

e. Electrolytes

i. Aqueous solution of substance conducts electricity due to  the ions present in the solution

ii. Dissociation- ionic compounds (solutes) dissolve as ions 1. Strong electrolytes dissociate completely

iii. Ionization- polar covalent molecule forms ions as it  


1. Not all polar covalent molecules

2. Strong acids

3. Weak electrolytes

a. Partial ionization

b. Weakly conduct electricity

c. Include weak acids and bases

II. Precipitation Reactions

a. Precipitate- an insoluble product that is formed in a chemical  reaction

b. Solubility differs for different substances

c. Exchange reactions/ Metathesis/ Double Replacement Reactions i. Cations and anions change partners from reactants to form products

ii. Reaction has to have a driving force

1. Precipitate/gas/weak or nonelectrolyte formation

iii. Precipitation reactions

1. Molecular equation  

a. all substances written as molecules

2. Ionic equation  

a. Soluble ionic substances written as individual  


b. Insoluble ionic substances written as formula

c. Strong electrolytes written as individual ions

d. Weak/nonelectrolytes written as formula

3. Net ionic equation

a. Eliminate spectator ions  

III. Acids, Bases, and Neutralization Reactions

a. Acids

i. Proton donors

ii. Ionize in water to form H+ 

iii. Strong acids (ionize completely)

1. HCl, HBr, HI, HClO3, HClO4, H2SO4

iv. Weak acids  

1. Weak electrolytes

2. Any acid that is not strong

b. Bases

i. Accept protons

ii. Produce OH 

iii. Strong bases

1. Hydroxides and oxides of groups 1A and 2A (except  

Be, Mg)

c. Neutralization  

i. Forms weak or nonelectrolytes

ii. Will go to completion if contains strong acid/base

iii. Driving force = formation of water

iv. Gas formation

1. Carbonates and bicarbonates, Sulfites and bisulfites,  sulfides

IV. Oxidation-Reduction Reactions

a. Loss of an electron of one element results in another element  gaining that elctron

b. Oxidation = losing an electron

c. Reduction = gaining an electron

d. Oxidizing agent = substance being reduced

e. Reduction agent = substance being oxidized

 f. Oxidation numbers- the charge that an atom would have if the  electrons in compounds were assigned a certain way

i. Elements = 0

ii. The sum of all of the oxidation numbers in a neutral  compound = 0

iii. Ox. numbers cant be more positive or more negative than  the group number

g. Combination reactions

h. Decomposition reactions

i. Combustion reactions

i. Produce CO2 and H2O

j. Displacement reactions

i. Single replacement- one element replaces another

ii. The Activity Series

1. How we know if a reaction is going to occur or not

V. Concentration of Solutions

a. Molarity = moles of solute/ Liters of solution

b. Dilution

i. M1V1 = M2V2

VI. Solution Stoichiometry

a. Used to determine the molarity of ions in a solution

b. Titrations

i. Determines the concentration of an unknown substance by adding to it a measured volume of solution of a known  


ii. Equivalence point- point where equivalent amounts of  reactants have reacted

iii. End point- point where the indicator changes color

Chapter 5 Notes: Thermochemistry


I. Thermochemistry- the study of energy changes that relate to  chemical reactions

a. Nature of energy

i. Energy = capacity to do work

ii. Work (J) = force x distance

iii. Heat = energy causes the temperature of something to  increase

iv. Calorie (cal) – amount of energy needed to raise the  temperature of 1g of water by 1°C

1. 1 cal = 4.184 J

b. Kinetic and potential energies

i. Kinetic Energy (J) = 1/2mv2

1. Motion energy

ii. Potential energy

1. Stored energy

2. Chemical energy

3. Electrostatic PE

a. The interaction between charged particles

c. System (what we single out to study) and surroundings  (everything else)

i. EX: contents of a flask(system) and the flask and  

everything outside of it (surroundings)

II. The First Law of Thermodynamics

a. The Law of Conservation of Energy 

i. Energy can neither be created nor destroyed, but can be  converted into a different form

ii. If energy is lost from a system, it must be gained by the  surroundings

b. Internal energy (E)

i. Total energy of system

ii. Actual value cannot be determined, but its change can  c. State function- property that depends only on its initial and final  value

i. Does NOT depend on the path that it took during a reaction ii. Extensive  

d. Heat (q) and work (w)

i. Are not state functions

ii. q is positive = endothermic (heat gained to system)

iii. q is negative = exothermic (heat lost to surroundings) iv. w is positive = work done ON the system

v. w is negative = work done BY the system

e. change in internal energy = q + w

f. exothermic reaction releases heat

g. endothermic reaction absorbs heat

i. requires input of energy

III. Enthalpy

a. Is a state function

b. At constant pressure  

i. w = -P∆V

1. negative because work is done BY the system

ii. ∆E = qp – P∆V

iii. qp = ∆E + P∆V

c. At constant volume

i. ∆E = qv

d. Enthalpy (H) = E + PV

e. ∆H = ∆E + P∆V

f. ∆H = qp

IV. Enthalpies of reaction

a. ∆Hrxn = Hproducts - Hreactants

b. endothermic reactions are the reverse of the exothermic reactions

c. ∆H is negative for an exothermic reaction and is positive for an endothermic reaction

V. Calorimetry

a. Experimental method utilizing the fact that any heat evolved or absorbed by the system will be reflected in the surroundings b. Heat capacity (C)

i. C = q / ∆T

ii. Molar heat capacity = heat capacity per mole

iii. Specific heat = heat capacity per gram

c. Calorimeter = ∆H at constant pressure

d. Bomb calorimeter = ∆E at constant volume

VI. Hess’s Law

a. ∆Hrx = ∑ ∆Hsteps

b. add the chemical equations for all of the steps to get to the overall reaction

c. add the ∆H of all the steps to get the ∆H of the overall reaction d. Tips when solving Hess’s Law problems:  

i. if you reverse an equation, the sign of that ∆H changes

ii. if you multiply an equation by a factor, that ∆H gets multiplied by the same factor

iii. any substances that don’t appear in the final equation must cancel out

VII. Enthalpy of Formation - The enthalpy change for the formation of a compound from its elements (∆Hf)

a. Standard enthalpy change- the enthalpy change when all  reactants and products are in their standard states (∆H°)

i. Standard state- the stable state of a substance in its pure  form at standard pressure and temperature

1. Solid or liquid

a. Pure substance at 1 atm

2. gas

a. pressure at 1 atm

3. species in solution

a. concentration of 1 M

b. standard enthalpy of formation (∆Hf°)- 1 mole of a compound is  formed from its elements with all of the substances in their  standard states

i. element in its standard state = 0

Chapter 6: Electronic Structure of Atoms

A. The wave nature of light

a. long wavelength = low frequency

b. frequency = speed of light/wavelength

B. Energy and Photons

a. Plank’s Theory = energy changes are quantized… discrete  energy changes

C. Photoelectric Effect

a. A minimum frequency of light shining on a metal surface causes  it to emit electrons

b. Radiant energy is quantized

c. Energy of a photon = Plank’s constant x frequency  

i. High frequency = high energy

D. Line spectra

a. Each line is associated with a different energy and color b. Different elements give distinct line spectra

c. Emission spectrum: its showing the light that’s being given off i. In the visible portion

d. Rydberg Equation

i. 1/wavelength = (Rydberg constant)(1/n12 – 1/n22)

ii. n2 is bigger than n1

e. Bohr Model

i. Electrons are restricted to certain energy levels  

corresponding to spherical orbits, with certain radii, about  the nucleus

ii. Ground state

1. The closer to the nucleus, the more negative

2. N = 1 is the closest to the nucleus

3. Lowest energy level; most stable state

iii. N = principle quantum number

1. N= 1 : Ground state for hydrogen

iv. Excited state

1. Electrons jump orbits when it absorbs energy

2. N is greater than 1

3. Higher energy, less stable, less negative energy

v. Zero point of energy

1. N = infinity  

2. Electron is completely separated from the nucleus

vi. Emission of energy = light

1. Electron emits light as a photon as it falls back to its  ground state  

E. Wave behavior of matter

a. De Broglie = matter should have wave properties

i. Wavelength = Plank’s constant/(mass)(velocity of particle) b. Heisenberg Uncertainty Principle

i. The wave-particle duality of matter makes it impossible to  precisely measure both the position and momentum of an  object

ii. ∆(Delta) does not mean change in the equation…means  uncertainty

F. Quantum Mechanics

a. Schröinger’s Wave Equation

i. The total energy of an H-atom is the sum of kinetic and  potential energy

b. Quantum numbers

i. N= energy level and distance from nucleus

1. Identifies the shell

2. Larger n means its further from the nucleus and at a  higher energy level

ii. L = shape

1. L = n – 1

2. Identifies subshell

iii. M = orientation of an orbital  

1. M = anything from –L to L, including 0

iv. Ms = spin  

G. Orbitals  

a. All s orbitals are spherical

b. Contour representation

i. Represent volume of space in which there is a high  

probability of finding the electron

c. All p orbitals have 2 lobes pointing in opposite directions  H. Many-electron atoms

a. Electron spin

i. Spinning charge creates a magnetic field

ii. Can only be either up or down (represented by +1/2 or  -1/2)

b. Pauli Exclusion Principle

i. No 2 electrons in an atom can have the same set of 4  quantum numbers

I. Electron Configuration

a. Orbitals filled in order of increasing energy until all electrons  have been used

b. For representative elements:

i. Period number = n value of valence shell

ii. Group number = number of valence electrons

c. Paramagnetic substance= has unpaired electrons; attracted by a  magnetic field

d. Diamagnetic = has no unpaired electrons; not attracted by a  magnetic field

e. Hund’s Rule

i. Electrons occupy different orbitals of a subshell until all are singularly occupied before electron pairing occurs

f. Exceptions:

i. 24Cr and 29Cu do not follow proper electron configuration  rules

Chapter 7: Periodic Properties of the Elements


I. Development of the Periodic Table

a. Electronic structure of atoms correlates with the properties of the elements

i. Reflected by the arrangements of the elements

b. Nobles Gases

i. Stable

ii. ns2np6 

c. Representative elements

i. Electron added to s and p orbitals

d. d-Transition elements

i. electron added to d orbitals

II. Effective Nuclear Change

a. The positive net charge attracting an electron

b. Zeff = Z – S 

i. Z = number of protons  

ii. S = screening constant

1. Average number of electrons between the nucleus  

and any one electron

2. Usually close to the number of core electrons

c. Core electrons screen/shield valence electrons from the full positive charge

i. Core electrons (inner shell) vs valence electrons (outer  shell)

ii. Valence shell electrons DO NOT experience the full nuclear  charge

iii. BUT core electrons do not completely shield the valence  electrons

1. p electrons DO NOT shield s electrons

2. s electrons DO shield p electrons

a. probable that these electrons are closer to the  


i. penetration

d. Coulomb’s Law

i. F = [k(QeQn)] / r2 

ii. Primary interaction of electrons and the nucleus is due to  charge

e. Zeff increases across a row

i. Increases by about 1 (moving left to right across the  periodic table)

ii. Each additional atom has 1 more proton and 1 more  unshielded valence electron than the one before it

f. Zeff increases slightly down a column 

i. Valence shell electrons can penetrate better

III. Atomic and Ionic Radii

a. Atomic radii

i. Nonbonding Radii

1. Based on crystal structures or gas phase collisions

2. Distance from nucleus to electron cloud in a single  


ii. Bonding Radii

1. Distance between covalently bonded atoms in a  


a. Bond length

2. B.a.r = ½ bond length or ½d

a. d = distance between nuclei

3. Size increases down a column/group 

a. N increases  atomic size increases

b. The outermost electron occupies a higher  

energy level with each descending element

i. Determines the average radius and size  

of orbital

4. Size decreases across a period 

a. Moving left to right

b. Electron is being added to the same shell

c. Zeff is increasing, so the whole shell is being  

pulled closer to the nucleus

5. Zef increases, radius decreases

b. Ionic Radii

i. Determined from the crystal structure of ionic compounds ii. The average interatomic distance from multiple  


iii. Cations

1. Smaller than the parent atom/neutral element

2. Size decreases with increasing atomic charge

iv. Anions

1. Larger than the parent ion/neutral element

v. Isoelectric series

1. Group of ions that all contain the same number of  


IV. Ionization Energy (I.E.)

a. Ionization = the removal of an electron

b. Ionization energy = the energy required to remove an electron  from a gaseous atom or an ion

i. The electron is removed from the highest energy level 1. The highest n and l

ii. Depends on the average distance from the nucleus  c. First ionization energy (I1)

i. The energy required to remove the highest energy electron from a neutral atom

ii. K  K+ + e 

d. Second ionization energy (I2)

i. The energy required to remove the next highest energy  electron from an ion

ii. K+  K2+ + e 

e. Successive ionization energies increase in magnitude i. The number of electrons decreases

1. Less repulsion

ii. The number of protons is the same

1. Greater attraction

f. I.E. increases as you move up a group/column 

i. Atomic radius decreases

ii. Electron held more tightly

g. I.E. increases as you move across a period 

i. Decreasing atomic radius

h. Exceptions

i. Be > B

ii. N > O

iii. ½ filled and filled subshells are more stable

iv. elements at the end of each transition series (Zn, Cd, Hg)  have higher I.E. than the following element

1. pseudo-noble-gas

v. highest I.E. for noble gases are filled s and p subshells i. Electron Configuration of Ions

i. Representative ions

1. Metals form cations

a. S-block (groups 1A and 2A)

i. Noble gas configuration

ii. All valence electrons removed

b. P-block (groups 3A-5A)

i. Lose p electrons easily

ii. Often require too much energy to remove

all of the valence electrons

2. Nonmetals  

a. Monoatomic anions

i. Charge = group number – 8

ii. Transition metal ions  

1. Often only the highest energy electron lost (outer s  subshell electron)

2. Often form 2+ cations

a. Lose both s-subshell electrons

3. For ions of higher charge, d-subshell electrons are  


V. Electron Affinity (EA)

a. The energy associated with the gain of an electron by gaseous  atom or ion  

b. First EA

i. The energy released for most neutral atoms and all  

positive ions

ii. Greater attraction for an electron  more negative EA iii. Cl + e-  Cl 

c. Second EA

i. Second electron must be forced onto a negatively charged  ion which requires energy

ii. O- + e-  O2- 

d. More negative as you move across a period 

e. More negative as you move UP a column 

f. Exceptions

i. 2nd period

1. F = -328 kJ/mol

2. Cl = -349 kJ/mol

ii. Group 2A

1. Full s-subshell

2. Added electron goes into p subshell

iii. Group 5A

1. ½ filled p valence subshell

2. added electron pairs with another electron in  

occupied p orbital

a. repulsed  

iv. Group 8A

1. Filled valence shell

2. Electrons goes into the next highest shell

VI. Metals, Nonmetals, Metalloids

a. Metallic character decreases moving across a period 

b. Metallic character decreases moving up a column 

c. Metals

i. Charges of 1A and 2A = group number

ii. Charges of p-block = group number OR (group number – 2) iii. Transition metals vary

iv. Form ionic compounds with nonmetals

1. Metal = oxidized

2. Nonmetal = reduced

v. Metal oxides = basic

vi. have luster, are malleable, ductile, and are good  

conductors of heat/electricity  

d. Nonmetals

i. Form ionic compounds with metals

ii. Form molecular compounds with other nonmetals and  metalloids  

iii. Oxides = acidic

iv. Are NOT lustrous, malleable, or ductile

v. Poor conductors of heat/electricity

vi. Vary widely in color and appearance, are dull  

e. Metalloids

i. Have properties of both metals and nonmetals

ii. Don’t form ions easily  

1. IE is too high to form cations

2. EA if not negative enough to form anions

iii. Form molecular compunds with nonmetals

VII. Summary of Periodic Trends

a. Metallic character and atomic radius increase 

i. Moving right to left across table

ii. Moving down a column

b. Ionization energy increases and electron affinity becomes more  negative 

i. Moving left to right across table

ii. Moving up a column

c. Zeff increases, radius decreases 

- IE increases -  Metallic character increases

- EA more negative -  Atomic radius increases

- Zef

VIII. Group 1A and 2A: Active Metals

a. Group 1A: Alkali Metals

i. Soft and low densities

ii. Larger ion size

iii. Valence electron configuration: ns1 

iv. Easily form M+ ions

1. One electron easily lost

2. Have Noble gas configuration

v. React rapidly with O2 and H2O

1. Reactivity with H2O increases as you move down the  group

vi. Preparation

1. Reduction of salt requires energy

vii. React with nonmetals to form ionic compounds

viii. React with oxygen to form oxides (Li2O), peroxides (Na2O2), and superoxides (KO2)

b. Group 2A: Alkaline Earth Metals

i. Harder, denser, less reactive than group 1A

ii. Smaller size and larger Zeff than group 1A

iii. Mg and Ca are the most abundant and the most important iv. React with oxygen to produce oxides

1. Except for Ba, which forms a peroxide

v. React with water

1. Reactivity increases as you move down the group

2. Heat is generated

3. Be = no reaction

4. Mg = slow reaction

5. Ca = more rapid reaction

vi. React with H2 to form ionic hydrides

IX. Nonmetals  

a. Hydrogen

i. Stable

ii. Isotopes

1. 1H (protium), 2H (deuterium), 3H (tritium)

iii. high IE

iv. gains an electron

1. forms hydrides (H-)

b. Oxygen

i. O2 = odorless, colorless gas

ii. Allotropes: O2 and O3 (ozone)

iii. Ions

1. O2- (oxides)

2. O22- (peroxides)

a. Unstable

3. O2- (superoxides)

c. Sulfur

i. S8 – ring, yellow solid

ii. S2- (sulfides)

iii. Acid rain

d. Nitrogen

i. N2 – odorless, colorless gas

ii. Very stable

iii. Compounds with H

1. Hydrazine, rocket fuel

iv. Compounds with O

1. Formed in combustion

2. Acid rain

e. Phosphorous

i. P4 – strained tetrahedral, solid

ii. In rocks, sand, soft drinks (generally as phosphates) iii. ADP = adenosine diphosphate

iv. ATP = adenosine triphosphate

f. Halogens

i. High EA

ii. Form X- ions

iii. Participate in redox reactions

iv. Compounds  

1. CFCs: refrigerants, greenhouse gases

2. Teflon

v. F2 is very reactive

1. Exothermic

g. Carbon  

i. Solid

ii. Covalent bonding  

iii. EX: diamonds, graphite

iv. Hydrocarbons  

h. Silicon

i. Solid

ii. Semiconductors

iii. SiO2 – sand, glass

iv. Silicates: SiO42-

v. Silicones: SiOR2

i. Boron

i. Octet exception

ii. Rocket fuel

iii. Diborane

j. Noble gases

i. monatomic

ii. full s and p subshells

iii. high IE

1. decrease moving down a group

iv. mostly unreactive

1. except Xe

Chemistry 1210: Chapter 8 Notes

Basic Concepts of Chemical Bonding

 I. Chemical Bonds- forces of attraction which hold atoms or ions  together

a. 3 Types

i. Ionic

1. Metals bonding with nonmetals

ii. Covalent

1. Nonmetals bonding with other nonmetals or  


iii. Metallic

1. Metals bonding with other metals

II. Lewis Symbols and The Octet Rule

a. Valence electrons are involved in chemical bonding

b. Representative elements

i. Electron in highest energy s and p subshells

ii. Number of valence electrons = group number

c. Lewis Symbols

i. Represent the electrons in the s and p orbitals of the  valence shell as dots arranged around the symbol of the  element

ii. Separate atoms show electrons as they appear in the  orbital diagram

iii. Bonds: place electrons around a symbol singly before  pairing them

d. Octet Rule

i. Formerly known as Noble Gas Rule

ii. Atoms tend to gain , lose, or share electrons in order to  achieve noble gas configuration (to have full s and p  


III. Ionic Bonding

a. Electrostatic attraction between cations and anions resulting  from a complete electron transfer

b. Bonds are very strong

i. Have high melting points

ii. Solids are brittle

c. Energetics  

i. An ionic compound is an array of cations and anions

1. Packed so attractive forces between ions of opposite  charges are maximized and repulsive forces between

ions of the same charged are minimized

ii. Lattice Energy

1. Measure of the strength of attraction between  

cations and anions

2. Energy required to separate 1 mole of ionic solid to

gaseous ions

3. LE approx. equals Q1Q2/d

a. Q1 and Q2 = the charges on the two ions

b. D = distance between the ions (sum of ionic  


4. LE increases as the charges on the ions increases  

and as their radii decreases

a. Greater charges = bigger LE

b. Smaller radii = bigger LE

c. Charge is the more important factor

i. Only if they are the same charge,  

consider the size

iii. Born-Haber Cycle

1. Based on Hess’s Law

2. Analyze the formation reaction for an ionic solid as a  

series of steps

IV. Covalent Bonds

a. Not a complete electron transfer

i. Pair of electrons are shared

b. Single Bond

i. 2 electrons shared by 2 atoms

ii. attraction of electrons for both nuclei holds the molecules  together

iii. Lewis structure

1. replace one bonding pair with a single line

c. Coordinate Covalent Bonds

i. Both electrons in the shared pair are donated by one atom ii. This bond is indistinguishable from other bonds once it  forms

Chem Chapter 9 Notes

9.4: Covalent Bonding and Orbital Overlap

I. Wave interference = electron behaves like a wave

 a. Constructive interference- waves add together to produce a  bonding orbital

 i. In phase combination- between 2 orbitals

 ii. Bonding orbital- large area spanning both atoms in which 2 electrons are free to move

1. The atoms share the electrons  

2. This orbital is lower in E than the 2 atomic orbitals  

 b. Destructive interference- waves subtract from each other and  get an anti-bonding orbital

 i. Out-of-phase combination 

1. Leads to no electron density between the 2 nuclei at  

the nodal plane  

 ii. Anti-bonding orbital 

1. This orbital is higher in E than the 2 atomic orbitals  

or the bonding orbital

 iii. Node- point where the wavefunction equals zero between  the 2 atoms  

II. Sigma Bonds

a. high electron density concentrated between nuclei along the  inter-nuclear axis

b. bond is attraction of electron nuclei and repulsion of nuclei  c. bond strength depends on:

i. degree of orbital overlap

ii. relative energies of the atomic orbitals which form the  


d. result from overlap of 2 s orbitals, s and p orbitals (end on), 2 p  orbitals (end on), s and hybrid orbitals, or 2 hybrid orbitals (end  on)

i. s + s/ s + p/ p + p = sigma bond

e. Morse Potential Energy curve

i. Nucleus-nucleus repulsion important

ii. Balance between attractive and repulsive forces

iii. PE decreases with increasing orbital overlap

III. Pi Bonds

a. Electron density above and below the inter-nuclear axis b. Results from sideways overlap of parallel p orbitals  

c. The side-side overlap in pi bonds is less efficient than for sigma  bonds

i. Pi bonds are weaker than sigma bonds

IV. Bonding Theories

a. Both Valence Bond theory and Molecular Orbital theory use the  idea of orbital overlap to create bonds

i. Main difference = when the orbitals are allowed to merge

9.5: Hybrid Orbitals- Valence Bond Theory

I. Valence Bond theory

a. Bonds are created by orbital overlap to produce sigma or pi  bonds

b. Explains the many observed molecular geometries

i. Pure s and p atomic orbitals are combined to produce a set  of hybrid orbitals on atoms

ii. These hybrid orbitals then form bonds between atoms  

producing the correct geometry  

c. Hybrid orbitals are responsible for electron domain geometries II. sp Hybrid Orbitals

a. should not form bonds

i. No singly occupied orbitals

b. As it forms, bonds can absorb enough energy to promote one 2s  electron to a 2p orbital

c. The s and p orbitals hybridize (mix) to form 2 degenerate sp  hybrid orbitals

i. These sp hybrid orbitals have 2 lobes similar to a p orbital ii. One of the lobes is larger and more rounded as is the s  orbital

iii. These 2 orbitals align themselves 180° from each other 1. Linear

III. sp2 Hybrid Orbitals

a. trigonal planar, 120°

b. As it forms, bonds can absorb enough energy to promote one 2s  electron to a 2p orbital  

c. The s and p orbitals hybridize (mix) to form 3 degenerate sp2 hybrid orbitals

i. These sp2 hybrid orbitals have 3 lobes with more p orbital  character than for sp hybrid orbitals  

IV. sp3 Hybrid Orbitals

a. tetrahedral, 109.5°

b. As it forms, bonds can absorb enough energy to promote one 2s  electron to a 2p orbital  

c. The s and p orbitals hybridize (mix) to form 4 degenerate sp3 hybrid orbitals

i. These sp3 hybrid orbitals have 4 lobes with more p orbital  character than for sp2 hybrid orbitals  

V. Hypervalent Compounds- Hybrid Orbitals?

a. Are hybrid orbitals necessary for ED geometries of trigonal  bipyramidal and octahedral?

i. Some say yes, some say no

b. VB theory requires you to hybridize the nd orbitals with the ns  and np  

i. sp3d = 5 hybrids; trigonal bipyramidal

ii. sp3d2 = 6 hybrids; octahedral

9.6: Multiple Bonds

I. Multiple bonds

a. Overlap of hybrid orbitals with s or p or other hybrid orbitals (end to end)

i. Sigma bonds

b. Electron density is symmetric about the internuclear axis of a  sigma bond

i. Groups can rotate about the bond without breaking it

1. Free rotation about sigma binds

c. Single bonds are sigma bonds

d. Multiple bonds require pi bonds

II. Double Bonds

a. Double bond = 1 sigma bond + 1 pi bond

b. Pi bond is perpendicular to plane

c. Trigonal planar around the atom

d. No free rotation

e. Cis-Trans isomers

i. No free rotation about double bond

III. Triple bonds

a. Triple bond = 1 sigma bond + 2 pi bonds

b. Linear around each atom

IV. Summary of VB Theory

a. For a given atom, all hybrid orbitals have the same energy b. Total number of hybrid orbitals = total number of atomic orbitals  used

c. Unhybridized p orbitals can form pi bonds or remain empty d. When atoms share more than 1 pair of electrons, 1 pair forms a  sigma bond and the rest form pi bonds

e. Hybrid orbitals form localized sigma and pi bonds

f. Geometries match VESPR theory  

g. Problems:

i. Doesn’t explain O2 paramagnetism  

1. Simple VB theory predicts any molecule with an even

number of electrons should be diamagnetic  

ii. Doesn’t explain excited states

iii. Doesn’t explain delocalized bonding well

iv. Doesn’t explain magnetic or spectral properties

V. Resonance and Delocalized Bonding

a. Localized sigma and pi bonds can’t explain resonance

i. But delocalized pi bonding can

9.7: Molecular Orbital Theory

I. MO Theory

a. Diamagnetic = NOT attracted to a magnet

b. Paramagnetic = attracted to a magnet

c. Orbitals are combinations of atomic orbitals from ALL atoms in a  molecule

d. The molecular orbital can span more than 2 atoms

e. Each molecular orbital can still only contain 2 electrons f. In VB theory, orbitals are first mixed on individual atoms and  then bonded together as needed

i. In MO theory, the orbitals of all atoms mix and then are  used to form the lowest energy molecular orbitals

II. Bonding and Anti-Bonding Sigma Orbitals

a. Can either add or subtract the atomic orbitals (electron density  of the atoms)

b. Bonding Molecular Orbitals

i. Add atomic orbitals: electron density concentrated in the  region between nuclei

ii. MO has cylindrical shape and is obtained from 1s AO

c. Anti-bonding Molecular Orbitals

i. Subtract atomic orbitals: cancel each other in region  

between nuclei

ii. Node between nuclei: electron density is zero

iii. * signifies an anti-bonding MO

d. Effectiveness of Overlap and Mixing  

i. How well 2 AOs overlap and mix depends on the type of  orbital and their relative energy

1. The closer in energy, the better the overlap

2. ∆E* > ∆E

ii. the orbitals must have correct phases to overlap

III. MOs formed from p AOs

a. Sigma MOs

i. 1 set of 2p atomic orbitals (2pz) overlap end-to-end (head on)

1. sigma2p and sigma2p*

b. Pi MOs

i. Other 2 sets (2px and 2py) of parallel 2p atomic orbitals overlap sideways

1. 2pi2p and 2pi2p*

IV. MO Energy-Level Diagrams, Electron Configuration, and Orbital Diagrams

a. Rules for determining MO Diagram

i. Number of MOs formed = number of AOs combined

ii. AOs combine by both in-phase (addition) and out-of-phase (subtraction) interactions

iii. AOs of similar energy combine most effectively  

iv. The more effectively AOs overlap the larger the energy difference between the MOs

v. AOs may NOT form MOs if there is no other orbital for it to interact with

1. EX: nbe- or lpe 

vi. Only occupied MOs affect the E of the molecule

vii. Pauli Exclusion Principle and Hund’s Rule are obeyed b. Period 1 Homonuclear Diatomics

i. Electron configuration

1. nb = number of bonding electrons

2. na = number of anti-bonding electrons

ii. bond order

1. number of bonds which exist between atoms

2. BO = ½(nb – na)

iii. Bond energy

1. Energy necessary to break 1 mole of bonds

2. Measure bond strength and stability

c. Period 2 Homonuclear Diatomics

i. Core MOs

1. 1s AOs

2. do NOT overlap effectively

3. remain mostly like the 1s AOs on atoms

4. contribute almost nothing to bonding

ii. valence MOs

1. 2s AOs

2. overlap effectively

3. valence AOs combine to give valence MOs

Chemistry Chapter 10 Notes: Gases and Kinetic Molecular Theory

10.2: Properties of Gases

A. Amount (mass or moles)

a. Low molar masses

b. Independent of volume, pressure, and temperature

B. Volume of gas = volume of container

a. Dependent on pressure and temperature

C. Both pressure and volume depend on temperature

D. Temperature must be in Kelvin

E. All gases are miscible (mix completely)

a. Homogenous mixture

F. Pressure = force/area

a. Depends on volume and temperature

b. Unit = pascal; 1 Pa = 1 N/m2 

c. Gas particles exert pressure by colliding with the walls of the  container

i. More collisions/area = higher pressure

ii. Barometer = measures pressure of atmosphere

iii. Manometer = measure pressure of gas/ gas above a liquid  in a vessel

iv. Standard Atmospheric Pressure

1. At 0° C at sea level  

2. 1 atm = 760 mm Hg = 760 torr  

10.3: Gas Laws

A. Boyle’s Law

a. P1V1 = P2V2

b. Constant temp and fixed amount of gas

B. Charles’s Law

a. V2/T2 = V1/T1

b. Constant pressure and fixed amount of gas

C. Avogadro’s Law

a. V = (k3)(n)

b. V1/n1 = V2/n2

c. Equal gas volumes, same T and P, contain same number of  particles

d. Determination of molecular weight

i. Mass 1 B (amu)/ mass 1 A (amu) = mass B (g)/mass A (g) 10.4: Ideal Gas Law

A. PV = nRT

B. R = 0.0821 (L)(atm)/(mol)(K)

a. Universal gas constant

C. Ideal gas hypothetically behaves according to the Ideal Gas Law under  all conditions

D. Standard temperature and pressure

a. T = 0° C = 273.15 K

b. P = 1 atm

E. Standard molar volume

a. 1 mole gas = 22.41 L

F. super combined gas law

a. PV/nT = R

b. P2V2/ n2T2 = P1V1/ n1T1

10.5: Further Applications

 A. Determine MW and molecular formula

 a. MF = (EF)n

 b. n = MF/EFW

 B. Determine EF and EFW from percent composition data  C. Determine MW

 a. PV = nRT, M = m/n, D = m/V, D = PM/RT

 D. Stoichiometry

 a. V2n2 = V1n1

 b. V = (k)(n)

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