2012 Midterm Review: AP Chemistry
Thane Jones
Unit 1- Chemical Foundations
Chemistry is the study of matter, its composition, properties, and the changes it undergoes.
Scientific Method
 Make observations
 Qualitative data
 Descriptive words
 Quantitative data
 Numerical values
 Making a Hypothesis

An educated guess to explain an observation
 Experimentation
 Gathering information
 Outcomes to support your experimentation
 Theory
 A set of tested hypotheses that gives an overall explanation of some natural phenomenon and
attempts to explain why it happens
 Law
 Same observations applies to many different systems, and it summarizes
Matter – Don’t trust it, it makes up everything.
Matter exists in pure substances and mixtures
There are 2 types of pure substances: Elements and Compounds.
 Elements are the simplest form of matter
 They are made up of all the same types of atoms
 They cannot be broken down
 Compounds are two or more elements
 They are bonded together in a fixed ratio
 They can only be broken down by chemical means
There are 2 types of mixtures, too.
 Homogeneous (solutions)
 A mixture that is the same composition throughout
 Gasoline
 Soda
 Ketchup
 Heterogeneous
 A mixture that has a varying composition
 Chocolate chip cookie
 A cake with icing
 Dirt
There are three different states of matter at room temperature.
 Solid
 Rigid/Can’t flow
 Particles are in a fixed position
 Doesn’t change much with temperature/pressure
 Definite shape and volume
 Liquid
 Flows
 Particles close but not fixed
 Indefinite shape
 Definite volume
 Little change with temperature/pressure
 Gas
 Flows
 Particles are far away from one another
 Indefinite shape and volume
 Large volume change with temperature increase

Compressible
There are two types of changes.
 Physical
 Doesn’t change the chemical composition
 Can change the appearance
 If Chuck Norris chews on a piece of Plutonium instead of gum, it’s still Plutonium when he spits it
out, but it’ll look squished
 Chemical
 The chemical composition changes (imagine that) and a new substance forms
 Appearance can still change
 Chemical formula changes
 If Chuck Norris drinks 2 glasses of 100M Hydrochloric Acid and Sodium Hydroxide before he goes
to bed, it won’t still be HCl and NaOH in his stomach.
There are 6 things that indicate a chemical change.
 Color change
 Gas production
 Odor change
 Evolution of heat or light
 Precipitate forms
 Temperature change
Measurement
Measurements are a form a quantitative data, which include an amount and a unit.
 26 kg
 1,000 watts
 9001 KJ
The Metric System
Physical Quantity
Mass
Length
Time
Temperature
Amount
Unit
Grams
Meters
Seconds
Kelvin
Mole
Abbreviation
g
m
s
K
mol
The Metric System is based on powers of 10
Prefix
Factor
Abbreviation
Mega10^6
M
Kilo 10^3
K
Deci10^-1
d
Centi10^-2
c
Milli10^-3
m
Micro10^-6
µ
Nano10^-9
n
Pico10^-12
p
You can change between the prefixes by multiplying or dividing your value by the factor of the target prefix.
There are 3 different temperature scales.
 Celsius
 °C
 °C = K - 273
 Fahrenheit
 °F = 1.8°C + 32
 Kelvin
 K
 K = °C +273
When measuring a value, you want to be as precise as possible.
We achieve this by estimating one decimal place further than the
most accurate indicator on you measuring device. If using an electronic
measurement device, you must estimate the drifting number, or, if the last
number is constant, use that one.
There is a big difference between accuracy and precision.
 Accuracy is how close to the target you are
 Precision is how close you are to other measurements
You must be able to calculate density!
(I use the Department of Motor Vehicles (DMV) to help me remember)
 D = M/V
 M = DV
 V= M/D
Be sure not to mix up mass and density!!
Significant Figures
There are 5 rules for significant figures. (Bolded numbers are significant, italics are not)
 All nonzero numbers are significant – 1,234,567,890
 Zeros in between two non-zero numbers are significant – 1.0234
 Zeros to the right of a decimal point and right of a non-zero number are significant - .023
 Zeros to the left of a decimal point but to the right of a non-zero number are significant – 2,000,000.
 Zeros to the left of a decimal point act as placeholders and are not significant – 2,000
Scientific Notation
We use scientific notation to express large numbers without having to write all of the zeros
 Move the decimal in between the first and second non-zero number
 Write all significant figures
 Count how many places you moved
 If you moved to the left, use a positive exponent of 10
 If you moved to the right, use a negative exponent of 10
Jay Batchu
Chapter 2: Atoms, Molecules, and Ions
Recall a very brief history of Atomic Theory and the definition of an isotope
Atomic theory is a scientific theory detailing the nature of matter. It states that all substances are composed of minuscule
particles known as atoms. It declares that atoms of the same kind are uniform in size, shape, weight, and other properties
(except isotopes: atoms with a differing number of neutron compositions).
Know and understand the five main aspects of Dalton’s Atomic Theory
1.
2.
3.
4.
5.
Elements are composed of small particles called atoms.
Atoms of a given element are identical in size, mass, and other properties. (This section of Dalton’s Theory is now
proven to be false).
Atoms cannot be divided, created, or destroyed. (This section of Dalton’s Theory is now proven to be false).
Atoms of different elements combine in whole number ratios to form compounds.
Atoms are combined, separated, or rearranged in chemical reactions.
Recall some of the experiments that led to the identification of sub-atomic particles
1.
2.
J.J. Thomson’s Experiment: He observed the deflection of charges in a cathode ray tube and formulated the concept
that states atoms are composed of positive and negative charges. The negative charges are the electrons. He also
said that they were sprinkled throughout the positively charged atoms like raisons in raison cookies. He also
generated the analogy of The Plum Pudding Model.
Millikan’s Experiment: He calculated the charge of an electron by observing the behavior of suspended charged oil
drops in an electric field.
3.
Rutherford’s Experiment: He fired alpha particles at gold foil and observed how they scattered. He concluded that all
of the positive portions of an atom are concentrated at its center and that most of an atom is empty space. This is
how the conclusions about nucleus and protons are made.
Know the three particles that make up the atom and their relative charges, masses and positions in the atom
Name of Sub-Atomic
Relative
Mass
Position
Particle
Charge
Proton
+1
1.67262158 x 10-27 kilograms or
Nucleus (center)
1.00727638 amu
Neutron
Nonexistent
1.6749 x 10-27 kilograms or
Nucleus (center)
1.0086649156 amu
Electron
-1
9.10938188 x 10-31 kilograms or
Electron cloud (orbiting
.0005446623 amu
the center)
Be able to use the atomic number and mass of an atom to calculate the number of protons, neutrons and electrons present
(and isotopes).
Protons = Atomic Number Neutrons = Mass Electrons = Atomic Number
Protons
Note: These formulas can also be used for isotopes with respect to the potentially different atomic mass
Understand the meaning of the terms Molecule and Ion
A molecule is a combination of atoms. Ex. H20 is a molecule of consisted of two hydrogen and one oxygen atoms. An ion is an
atom or a group of atoms that carries an electric charge. Ex. O2- is an oxide ion with a charge of negative two. An anion is an
ion with a negative charge and a cation is an ion with a positive charge.
Learn the lists of common anions, cations, and polyatomic ions
NOTE: Refer to the last sheet of the “LET’S GET TO WORK” packet received at the beginning of the school year for a more
detailed list.
Cation
Anion
Polyatomic Ions
3+
+
2+
2+
2+
3+
+
2+
+
23Al ,NH4 ,Ba ,Ca ,Cr ,Cr ,Cu ,Cu ,F I , Br , N , Cl ,
NH4 ,C2H3O ,OCN-,CN-,H2PO4-,OH2+
3+ +
+
2+ +
2+
2+
22e ,Fe ,H ,H3O ,Pb ,Li ,Mg ,Mn ,M
,ClO4-,MnO4-,SCN-,CO32-,CrO42-,Cr2O72S ,F,O
3+
2+
2+
2+ +
+
+
2+
2
n ,Hg2 ,Hg ,NO ,K ,Ag ,Na ,Sr ,Sn
,HPO42-,SO42-,SO32-,S2O32-,PO43-,BO33+
4+
2+
,Sn ,Zn
Know how to combine those anions and cations in the correct proportions to form ionic compounds with no net charge
Use the crisscross method. Swap the charges of the ions and simplify if necessary.
Switch charges
Simplify by dividing by 2
Be able to name binary ionic compounds of a metal and nonmetal
The cation (metal) will come first and is followed by the anion (nonmetal). The anion will gain the suffix “ide” to the end of
the stem. For example, in the compound NaCl, Na+ is the cation and Cl- in the anion. So NaCl is to be called Sodium Chloride.
Be able to name binary molecular compounds of two nonmetals.
The Greek prefix (mono – 1, di – 2, tri – 3, tetra – 4, penta – 5, hexa – 6, hepta – 7, octa – 8, nona – 9, and deca - 10)
corresponding to the number of atoms of the first subject (the less electronegative) comes (if mono is the corresponding
prefix, it is avoided) and then it is followed by the name of the first subject and then it is followed by the Greek prefix
corresponding to the number of atoms of the second subject and finally the name of the second subject comes with an “ide”
ending. For example, CO can be described as carbon monoxide. CO2 can be described as carbon dioxide. P2S3 can be
described as diphosphorus trisulfide.
Be able to name simple binary acids
Binary acids are certain molecular compounds in which hydrogen is combined with a nonmetallic element. The names of
binary acids begin with hydro- and is followed by the second subject with a suffix of ic. For example, HS is called hydrosulfic
acid. HCl is called hydrochloric acid.
Be able to name ionic compounds containing polyatomic ions including oxyanions
The cation’s name comes first and then the anion’s name comes with an ending of either –ate or –ite but usually -ate. –ate
is used to designate greater amount of oxygen atoms and –ite is used to designate a fewer amount of oxygen atoms present
in a polyatomic ion. If the anion has more oxygen atoms than the –ate ion, the prefix per- is added but if the anion has less
oxygen atoms than the –ite ion, the prefix hypo- is added. For example, KClO is potassium hypochlorite. KCLO 2 is potassium
chlorite. KClO3 is potassium chlorate. KClO4 is potassium perchlorate.
Be able to name oxyacids and compounds containing oxyanions
Acids are substances whose pH is less than seven. Acids are compounds (anions) bounded to a hydrogen atom. To name an
acid, the suffix –ic or –ous is added to the end of the anion and it is followed by the word acid. –ic is used to replace –ate and
–ous is used to replace -ite when using acids. For example, H2SO3 is called sulfurous acid. Oxyacids are acids that contain an
oxygen atom bonded to a hydrogen atom and at least one other element. For example, H 2SO4 (sulfuric acid) and H3PO4
(phosphoric acid) are oxyacids.
QUESTIONS:
1. What is hypochlorous acid? A. HClO2 B. HClO C. HClO3 D. HCLO4 (B. There is only one oxygen)
2. What is dihydrogen monoxide? A. H3O B. HCN C. H2C2 D. H2O (D. There are supposed to be two hydrogen atoms and
one oxygen atom)
Ketan Choski
Chapter 4- Solution Stoichiometry
Predict whether a substance is a strong, weak, or non-electrolyte



Strong electrolyte- Strong acids, strong bases, soluble ionic compounds
Weak electrolyte- Weak acids, weak bases (covalent)
Non-electrolytes- Covalent compounds
Understand what aqueous means and its importance in reactions

Aqueous- Any substance that is extensively ionized when dissolved into water. Dissolved substances can react with
other ions to form stronger insoluble compounds.
Predict the ions that an electrolyte dissociates into


Cation- First atom/polyatomic ion becomes element and positive charge, e.g. Mg+2
Anion- Last atom/polyatomic ion becomes element and negative charge, e.g. Cl-
Know how to calculate the moles of a solution using molarity

M = Moles/1L
o E.g. 10M * 2L = 20 mol
Be able to calculate the concentrations of ions when they dissociate

Use dissociation reaction to determine moles and then divide by volume to determine molarity
Identify substances as acids, bases, and salts



Acids- Always begin (or cation for ionic acids) with H
Bases- Anion is always OH- for ionic bases
Salt- Able to be broken down into anion and cation
Be able to calculate dilution problems

Dilution- Lowering a solution’s concentration by adding solvent


M1V1=M2V2
M = molarity, V = volume
Define a neutralization reaction

An acid and a base react with each other to form a salt and water
Understand the relationship between strength of an acid or base with its degree of dissociation

Greater dissociation = greater strength, determined by reaction with water
Understand the titration process

Titration – determining the concentration of an unknown substance (titrant) using the concentration of a known
substance (analyte)
Be able to perform mole calculations using titration data

Multiply volume of titrant by molarity, multiply by mole ratio, divide by volume to determine concentration
Using solubility rules, predict if a precipitate forms in a double placement reaction, predict its products, and write a
balanced molecular chemical equation


Determine products by swapping cations or anions, use solubility rules to predict if either of the products is a
precipitate
Balance equation based on charge of all ions involved in the reaction
Predict the products and write a balanced chemical equation for an acid-base reaction

HX + YOH -> H2O + YX
Be able to identify spectator ions and write ionic and net ionic equations



Split reactants and products into to get complete ionic equation
Cross out ions that remain in the products. The remaining equation is the net ionic equation
The crossed out ions are the spectator ions
Define oxidation and reduction and recognize their relationship


Oxidation is the loss of electrons, reduction is the gain of electrons
Substance oxidized loses electrons to substance reduced
Become familiar with some common oxidizing and reducing agents to help you identify a redox reaction


Single replacement- Metal is oxidized, cation is reduced.
Synthesis, decomposition, and combustion reactions are all redox
Assign oxidation numbers to atoms
1.
2.
3.
4.
5.
6.
Elements in their standard states have a oxidation state of 0
Oxidation state for monatomic ions is the same as the charge when in a compound
Oxygen is assigned an oxidation state of –2 in its covalent compounds except as a peroxide (-1)
As a cation or in a covalent compound, hydrogen is assigned an oxidation state of +1. -1 when hydride
Fluorine is always –1
The sum of the oxidation states within a compound must be equal to zero and equal to the charge when in a
polyatomic ion
Determine whether a reaction is redox or not


Reactions where a participant changes its number of electrons are redox
When all participants keep the same number of electrons, the reaction is not redox
Multiple Choice #1
What is the oxidation number of chromium in CrO42-ion?
A)
B)
C)
D)
E)
+14
+7
+3
+12
+6
E because oxygen has a oxidation state of -2, and the sun of oxidation numbers must be -2. Therefore, Cr is +6
Multiple Choice #2
Complete neutralization of 50 mL of .10 M NaOH solution requires 80 mL of HCl solution. What is the molarity of the HCl
solution?
A) 1.0
B) 0.063
C) 1.3
D) 0.75
E) 2.7
B because .05L NaOH * .10 M = .08L HCl * X. X= .063 M
Reed Williams
Chapter 4 Solution Stoichiometry
An intro into Solution Stoichiometry- solution stoichiomtery is the mathematical study of reactions that take place within
solutions and their properties. Notably, topics such as molarity and solubility will be covered along with acid-base, REDOX,
and precipitation reactions. Understanding this topic is paramount in acing the reactions section of the AP exam.
Basic definitions
Soluble- can be dissolved
Solute- what gets dissolved
Solvent- what does the dissolving
Aqueous- dissolved in water. Remember, like dissolves like (polar dissolves polar).
Molarity- concentration of solution. Abbreviated with M
Dilution- adding more solvent to a solution. Decreases the molarity.
ConductivityElectrolyte- compound that conducts electricity. A strong electrolyte will dissolve completely while a weak one will only
partially ionize ( break into ions).
Strong electrolyte- soluble ionic compounds/ strong acids and bases
Weak electrolyte- covalent compounds
As a side note, acids are compounds which form H+ ions when dissolved ( usually begin with H). Bases form OH- when
dissolved ( usually end with OH). A salt is an ionic compound featuring a metal and nonmetal.
Types Of ReactionsPrecipitation reaction- This type of reaction results in a precipitate (insoluble salt). After completing the initial steps of a single
or double replacement reaction, the memorized solubility rules will be recalled. Those compound which are soluble will split
into ions and those that appear on both sides will be crossed out ( called spectator ions). The product that remains intact in the
net ionic equation, what is left with spectators discarded, is the precipitate.
Acid-Base reaction- Often called a neutralization reaction because the acid neutralizes the base, this process can be defined as:
Acid + Base--> salt + water. This reaction is demonstrated through a Titration. In a titration, an acid is slowly dripped into a
base until the equivalence point has been reached and the indicator changes color. Stoichiomtery may be required to convert
from moles of one substance to another before utilizing the molarity equation ( expressed in equation section of review).
REDOX/ Oxidation reaction - Involves the transfer of electrons. Memorize a few common elements: O is -2, H is 1 with nonmetal and -1 with metal, F is -1 , etc. The sum of oxidation numbers is a compound will always equal 0. Oxidation is the
loss of electrons while reductions is the acquisition of electrons. However, oxidizing agent gets reduced and reducing agent
gets oxidized. Half Reactions are often used to simplify the balancing of a REDOX reaction. 1. Write equation of oxidation
and reduction. 2. Balance except for H and OH. 3. Balance O using water. 4. Balance H with H+. 5. Balance with e- 6.
Multiply to get equal number of e- 7. Add equations and cancel identical entities. 8. Check that charges and elements are
balanced.
A few types of Redox reactions exist:
Combination reactions- A +B--> AB
Single replacement- AC + B --> BC +A
Decomposition- AB--> A+ B
Double replacement- AB+ CD--> AD + CB
Combustion - hydrocarbon and oxygen form water and carbon dioxide
In s single replacement reaction, only elements higher on the activity series will replace. On AP exam, all reactions are
assumed to go.
Important equations :
Molarity= moles of solute/ liter of solution. Make sure you have converted to liters. With this equation, the molarity can be
calculated. It can also be used to find one of the other two pieces.
M1 x V1= M2 x V2. This equation is used to calculate molarity after a dilution. Since moles= M x V and the number of moles
remains unchanged, a missing variable can quickly be calculated.
Reference Materials
Solubility rules must be memorized.
Chapters 5 & 9 in Princeton review book. Problems on pg. 74 and 142
Chapter 4 in textbook. Problems on pg. 179-185. Practice is the best way to study these types of questions.
The notes for chapter 4 feature good explanations and review problems.
Practice Questions:
When sodium chloride is added to a saturated aqueous solution of silver chloride, which of the following precipitates would be
expected to appear? A. Sodium B. silver C. Chlorine D. Sodium chloride E. Silver chloride
Answer: E. using the memorized solubility rules, you should know that silver chloride is insoluble
2. When 300. ml of a 0.60 M NaCl solution is combined with 200. Ml of a 0.40 M MgCl2 solution, what will be the molar
concentration of Cl- ions? A. .20 M B. .34 M C. .68 M D. .80 M E. 1.0 M
Answer: C. Use MV= moles to find moles of Cl- in each compound ( remember that there are two in MgCl2). Add the two to
find total moles. Put moles over liters (.500 ml) and divide. Since M= moles of solute/ liters of solution, you should get .68M
as the final answer.
Harry Nguyen
Unit 5: Gases
reference pages 188-229
Properties of Gases:
1. small particles spaced apart
2. indefinite shape and volume
3. high expansion and compressibility
4. considered a fluid
5. expands as temperature rise
Pressure:
measured in: Pa, kPa, torr, atm, mm Hg, bar
conversions: 1 atm = 760 torr = 760 mm Hg = 101.3 kPa
Standard pressure = 1atm
Kelvin and Celsius:
To find Kelvin, K= o C + 273
Standard temperature= 0 C or 273 K
Laws of Boyles, Charles, Gay-Lussac and Avogadro:
Volume of gas is inversely proportional to pressure at constant temp (Kelvin) and moles. V 1P1=V2P2
Volume of a gas is directly proportional to temperature at constant Pressure and moles. V 1 / T1 = V2 / T2
Pressure of a Gas is directly proportional to Temperature (Kelvin). P1 / T1 = P 2 / T2
Number of gas molecules are directly proportional to Volume. V 1 / n1 = V2 / n2
Law of the Gas:
COMBINED GAS LAW: when all variables are changing but moles stay the same. V 1 P1 / T1 = V2 P2 / T2
IDEAL GAS LAW: describes ideal gas now, not when variables change. PV = nRT R=.08206 L * atm/mol*K
To find Density: D= MP/RT
To find Molar Mass: M=mRT/PV
Properties of Ideal Gases and a Real Gas
Real gases occupy space and exert attractive forces as Ideal Gases do not.
Real gases behave like ideal gases at ordinary temperatures. When heated their behavior departs from ideal.
Dalton's Law of Partial Pressures:
Each gas exerts a specific pressure, the sum of these pressures with add up to the total pressure of the gas mixture.
PT = P 1 + P 2 + P 3 …
Water Vapor will mix with gas and must be accounted for using PT = Pgas + Pwater
Total moles of a gax in a mixture can be found using mole fraction (X)
X = n 1 / nT
moles and pressure are proportional. P1 / PT = X
PRACTICE QUESTION
1. How much total Pressure is there if there is 700 mm Hg of Carbon, 2.4 Atm of Hydrogen and 200kPa of Oxygen?
a. 9.024 Atm
b. 5.30 Atm
c. 3.36 Atm
d. 6.42 Atm
STP:
STP is the abbreviation for Standard Temperature and Pressure. 0C and 1Atm
At standard conditions 1 mole of gas has a molar volume of 22.4 L.
Effusion and Diffusion:
Effusion = Escape of gas via small gaps
Diffusion = the spread of a substance due to high to low concentration.
Graham's Law: the lighter the molecules the faster they will move, or the heavier the particles the slow they move.
PRACTICE QUESTION
2. Which gas will effuse the quickest?
a. Oxygen
b. Chlorine
c. Helium
d. Carbon
Root Mean Square Speed:
 The average speed of a molecule possessing avererage Kinetic Energy.
 As the temperature increases, the KE and Speed increases.
◦
Urms = √ 3RT/M
M= molar mass (Kg/mol)
T= Kelvin
Kinetic Molecular Theory





Gas particles are in constant random motion.
They have no volume
No forces of attraction or repulsion
Elastic collisions between molecules. CONSERVATION OF ENERGY
Average KE is directly proportional to Kelvin Temp.
PRACTICE PROBLEM ANSWERS
1. B) 5.30 Atm. First you convert all the values to Atm and then add them up using the sum of partial pressures.
2. C) Helium because it is the lightest gas out of the 4 choices.
Luke Massey
Unit 5: Gases
Important things to know:101.3 kPa = 760 torr = 760 mm Hg = 1atmmolar volume = 22.4 litersgas constant = .08206P is
pressure, V is volume, n is the number of moles, R is the universal gas constant, and T is temperature (K)
Gases – Fill container (volume of gas is = volume of container), form homogenous mixtures, are far apart due to weak
attractive forces, and are easily compressed
Solid/Liquids – Do not fill containers, are close together due to strong attractive forces, and are not easily compressed
Boyle's law -
Charles's law -
Gay-Lussac's law -
Avogadro's law ideal gas law -
combined gas law Dalton's law Gas Collection over H2O - PT = Pgas +Pwater
Kinetic Molecular Theory
1) Particles are small compared to the distance between them – Volume is assumed to be zero
2) Particles are in constant (random) motion and collide with the walls of the container – Pressure changes
3) Particles exert no force on each other – Do not repel/attract
4) Average kinetic energy is directly proportional to the absolute temperature (K) – At a given temperature, the molecules
have the same average kinetic energy
5) Energy can be transferred between molecular collisions, but the average kinetic energy of the molecules does not change
as long as the temperature remains constant – Collisions are elastic
Root mean square -
Graham's law Questions
A mixture of nitrogen and neon gases contains equal moles of each gas and has a total mass of 10.0 g. What is the density
of this gas mixture at 500 K and 15.0 atm? Assume ideal gas behavior.
1) Convert 30.0 mm of H2O to equivalent mm of mercury:
(30.0 mm) (1.00 g/mL) = (x) (13.534 g/mL)
x = 2.21664 mm (I will carry some guard digits.)
2) Convert mmHg to kPa:
2.21664 mmHg x (101.325 kPa/760.0 mmHg) = 0.29553 kPa
3) Determine pressure of enclosed wet CH4:
At point A in the above graphic, we know this:
Patmo.press. = Pwet CH4 + Pthe 30.0 mm water column
98.70 kPa = x + 0.29553 kPa
x = 98.4045 kPa
4) Determine pressure of dry CH4:
From Dalton's Law, we know this:
Pwet CH4 = Pdry CH4 + Pwater vapor
(Water's vapor pressure at 30.0 °C is 31.8 mmHg. Convert it to kPa.)
98.4045 kPa = x + 4.23965 kPa
x = 94.1648 kPa
Based on provided data, use three significant figures; so 94.2 kPa.
Three 1.00 L flasks at 25.0 °C and 1013 hPa pressure contain: CH4 (flask A), CO2 (flask B) and NH3 (flask C). Which flask (or
none) contains 0.041 mol of gas?
Three important points:
(a) All three flasks are at equal pressures (by the way, 1013 hPa is not usually seen. It is the same as 101.3 kPa, which is 1
atmosphere).
(b) All three flasks are at the same temperature.
(c) All three flasks have the same volume.
The above satisfies Avogaro's Hypothesis: equal volumes of gases under the same conditions of pressure and temperature,
contain equal number of molecules.
Therefore, all three flasks either all contain 0.041 mol or none does. We will check flask A, using PV = nRT:
(1.00 atm) (1.00 L) = (n) (0.08206 L atm mol¯1 K¯1) (298 K)
n = 0.0409 mol
All three flasks contain 0.041 mol of the different gases.
Hal Norris
Chapter 6- Thermochemistry
Vocab:
System: Part of universe that we focus on
Surroundings: Everything else in universe, you measure the temp of the surroundings
Enthalpy (H): Heat of reaction, ΔH= Hf-Hi, -ΔH=Exothermic Rxn +ΔH= Endothermic Rxn
Heat Capacity: Amount of energy required to raise temp of object 1 degree Celsius
Specific Heat Capacity: Energy required to raise temp of one gram of substance 1 degree Celsius
Molar Heat Capacity: Energy required to raise temp of one mole of substance 1 degree Celsius
Standard State: Precisely defined reference state to properly compare a substance with
Units of Energy:
Joule (J), Kilojoule (kJ)
Calorie (c), 1 calorie= 4.184 joules
State Functions:
State Functions refer to a property of a system that depends only on its present state. These state functions include Internal
Energy (E), Enthalpy (H), Entropy (S), and Gibb’s Free Energy (G). All of the characteristics are specific to the state a substance
is in.
First Law of Thermodynamics:
The First Law of Thermodynamics states that energy can change forms but can be neither created nor destroyed. Energy is
constant and be classified as potential or kinetic energy. Potential energy is due to composition or position of the object.
Kinetic energy is due to the movement of the object.
Endo and Exothermic Reactions:
Endo: Reactions that absorb energy from the surroundings are said to be endothermic. They have a positive ΔH because the
final energy in the system is greater than the initial. A chemical reaction that has energy in the reactants are endothermic
(N2+O2+ Energy 2NO)
Exo: Reactions that result in the evolution of heat are exothermic. They have a negative ΔH because the final energy of
system is less than when it began. A chemical reaction has energy in the products (CH 4+O2CO2 +2H2O+ Energy)
Hess’s Law:
Hess’s Law is a principle that states that the enthalpy of reaction is the same no matter what. So you can do it in one step or
multiple steps and the answer won’t change. Examples are shown on pg 256 of the textbook. The heat of reaction will be
equal the sum of the ΔH of each step. You can also use Hess’s Law with bond energies. An example is on pg 372
Questions:
1. Which statement correctly describes an endothermic chemical reaction?
A. The products have higher potential energy than the reactants, and ∆H is negative.
B. The products have higher potential energy than the reactants, and the ∆H is positive.
C. The products have lower potential energy than the reactants, and the ∆H is negative.
D. The products have lower potential energy than the reactants, and the ∆ H is positive
2. Given the two reactions below, what is the ∆H for the reaction, IF5(g)  IF3(g) + F2(g)?
IF(g) + F2(g)  IF3(g) ∆H = -390 kJ
IF(g) + 2F2(g) IF5(g) ∆H = -745 kJ
A. -1135 kJ B.35 kJ C. 355 kJ D. 1135 kJ
Answers:
1.
2.
B, the products have a higher potential energy because they are absorbing energy. ΔH is positive because the final
energy is greater than the initial.
C, IF5 IF+2F2 ΔH=745 kJ IF+F2IF3 ΔH= -390, 745+-390=355kJ
Book Pages:
Chapter begins on 241,Endo+Exo explained on pgs 244+245, Calorimetrypg 251, Hess’s Law pg 257, more practice questions
pg 282
Emily Elliot
Unit 7—Atomic Structure and Periodicity
Wave-Particle Duality for Light
Light is viewed in two ways in the scientific world
As a particle known as a photon. This states that energy may be released in packets of light or quanta which relate to
the equation (E=hv) where E is the energy released, h is Planck’s constant (6.63x10^-34 J-s), and v is frequency.
As a wave. This relates to the electromagnetic spectrum which says that energy is released in waves that all travel at
the speed of light (c=3.00x10^8 m/s). The amount of energy contained in these waves is determined by the frequency
and wavelength. If there is a high frequency and a small wavelength, there will be a lot of energy in the wave.
Wavelength and frequency may be determined by the equation (c= λv), where λ is wavelength and v is frequency.
The two equations may also work together to determine the energy of a wave. If c=λv, then v=c/λ. Therefore if one plugs in the
c/λ for v in the E=hv equation, an equation to find the total energy released in one photon may be found.
E=h(c/λ)
Line Spectra
The electromagnetic spectrum is a tool that displays all forms of radiant energy in order of increasing wavelength and
decreasing frequency.
gammaXray
UV
visible
IR
microwaves
Radio
ROYGBIV
ROYGBIV represents the colors of the visible light spectrum in order of increasing wavelength, however to flow with the
direction of this spectrum they would actually be read as – violet, indigo, blue, green, yellow, orange, red.
Gases emit different colors at specific wavelengths. Hydrogen has 4 spectral lines. The Rydberg equation may be used to
determine the wavelengths of these lines.
1/λ=(Rh)(1/(n1)^2-1/(n2)^2)
Rh is the Rydberg constant, n1 and n2 represent different energy evels.
Quantum Mechanical Model of the Atom
One cannot give the exact location, but a predicted location of an electron, however one can locate an orbital—and area where
there is a high probability of finding an electron. This model is based on complex mathematical equations that allow us to
easily find areas where electrons can most likely be found.
Bohr: Created a model for the atom that assumed electrons moved in circular orbits around the nucleus of the atom in areas
called energy levels. Also said that a fixed amount of energy is released/absorbed as an electron moves in or out of an excited
state. Could not explain spectra of other atoms, electrons do not circle the nucleus.
De Broglie: Made the assumption that if light waves can particulate behavior, then electrons can have wave-like properties.
Basically, that an electron changing energy levels can emit a decent amount of energy for such a small particle.
Pauli: Created the Pauli exclusion principle which states that no two electrons in a single atom can have the same 4 quantum
numbers.
Heisenburg: Created the uncertainty principle that tells us there is a limit to how well we can know the position and
momentum of an electron at any certain time.
Schrodinger: Had a system of equations whose solutions are described by quantum numbers
Quantum Numbers
Principle quantum number (n): Indicates the main energy level of an electron. Values of n=1-7. As n increases, electron energy
increases and so does distance from the nucleus. Max # of electrons in the level 2n^2
Angular momentum quantum number (l): indicates the shape of the orbital (subshell). Values from 0-(n-1).Spherical,
dumbbell, four-lobed. Nth energy level= n sublevels
Magnetic quantum number (m1): indicates orientation of an orbital around the nucleus. Values from -1 to +1. S=1, p=3, d=5,
f=7.
Electron spin quantum number (ms): shows the spin of the electron. Can be +1/2 or-1/2
Electron Configurations and Orbital Notations
Aufbau Principle: electrons fill orbitals from lowest to highest energy level 1s,2s,2p,3s,3p,4s,3d,4p…. Exceptions occur in
the d sublevel when and electron may be taken from the s sublevel to give the d orbital a more stable configuration (filled or
half-filled)
This periodic table shows the respective blocks for orbitals on the periodic table. For example using the different blocks on the
table one can see that the electron configuration for F is 1s2 2s2 2p5 based on simply counting the blocks and sublevels
Periodic Trends
Atomic Radii- Period, decreases because there are more protons and electrons in the same energy level, meaning nuclear
charge is greater, pulling electrons in more closely. Group, increases because electron in higher energy levels cause shielding
and make it more difficult for the nucleus to pull outer electrons in.
Ionization energy- Energy required to remove an electron. Period, increases because there is a greater nuclear charge as an
atom gains protons and electrons in the same energy level. Group, decreases because shielding occurs making it more difficult
for the radius to hold onto electrons in the higher energy levels.
Electron Affinity- Energy to gain an electron. Period, increases because the nucleus has a greater charge when adding more
protons and electrons in the same energy level so it holds on to electrons more tightly. Group, decreases because of shielding.
The atom has many electrons in high energy levels that do not have a great pull on them so the atom does not pick them up as
easily and usually requires energy to gain them.
NUCLEAR CHARGE IS THE PULL OF THE NUCLEUS ON THE SURROUNDING ELECTRONS.
Properties
Metals: Shiny luster, conduct heat and electricity, malleable,
ductile, solid at room temperature, mostly silver in color.
Nonmetals: Not lustrous, poor conductors of heat and electricity,
lower melting points than metals, brittle and soft solids, gases, and
one liquid at room temperature.
Metalloids: Contain properties of both metals and nonmetals.
Alkali Metals: Soft, silvery, most chemically reactive metals, react with nonmetals to
form ionic compounds, reaction with water is extremely exothermic and volatile, stored in
kerosene oil
Alkaline Earth Metals: Harder than alkali, silver, high thermal/electrical conductivity, not
as reactive as alkali with water, for ionic compounds
Equations
1.
𝐸 = ℎ𝑣
2.
λ=
3.
4.
𝑐 = λv
𝑝 = 𝑚𝑣
Halogens: Extremely reactive, gas at room temp except for bromine, mostly diatomic
Noble Gases: inert, monatomic, full octet, very unreactive, found in lightening.
MULTIPLE CHOICE
1.
2.
A laser has a wavelength of 7.80x10^2 nm, calculate the frequency of the light.
a. 3.84x10^14Hz
b. 2.45x10^14 Hz
c. 1.00x10^14 Hz
d. 4.63 x10^14 Hz
Give the correct electron configuration for Rb
a. 1s2 2s2 2p6 3s2 3p6 4s2 4d10 4p6 5s1
b. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
c. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1
d. 1s2 2s2 2p6 3s2 3p6 4s2 4p6 5s1
Nathan Larkin


Chapter 7 Atomic Structure & Periodicity
Explain line spectra
o Inverse relationship between wave length and frequency
o Gamma Rays are the strongest Radio Waves are the weakest
o Visible range is very small, ranges from 400 nm to 700 nm.
 Ultra = Strongest
 Red = Weakest
Calculate wavelength, frequency, and energy of a photon
o Use equation #1.
 E = energy



h = plank’s constant (
2
6.626068 𝑥10−34 𝑚 𝑘𝑔
𝑠
)
 v = frequency
Calculate wavelength for line spectra using the Ryberg Equation
o Use the last equation
 v = frequency

is the Rydberg constant, approximately 1.097 * 107 m-1

and
are integers greater than or equal to 1 such that
Describe the quantum mechanical model of the atom
and .
ℎ
𝑚𝑣
o
o

The first Principal quantum number identifies which energy level an electron is in.
There are 7 possible energy levels in an atom in the ground state (stable) the azimuthal quantum number
identifies the sub-level with in the energy level where the electron is most likely to be found.
o There are 4 types of sub-levels, the magnetic quantum number indicates a region within a sub-level called
an orbital where two electrons reside.
o The number of orbitals in a sub-level depends on the type of sub-level. To distinguish between the two
electrons in an orbital, the electron-spin quantum number is used. There is two possible electron-spin
quantum numbers, 1/2+ and 1/2Describe the contributions of Bohr, De Broglie, Pauli, Heisenberg, and Schrodinger
o Bohr
 Bohr took the quantum theory and used it to predict that electrons orbit the nucleus at specific,
fixed radii, like planets orbiting the Sun. Bohr model worked for atoms and ions with one electron
but not more complex atoms.
o Pauli Exclusion Principle
 States that within an atom, no two electrons can have the same set of quantum numbers. So, each
electron in any atom has its own distinct set of four quantum numbers.
o Heisenberg Uncertainty Principle
 It is impossible to know both the position and the momentum of an electron at a particular
instant.
o De Broglie Hypothesis
 All matter has wave characteristics

λ=
h
𝑚𝑣
o


Schrodinger
 Current Quantum Mechanical Model of the Atom
Explain the limitations of the Bohr Model
o Only works for the Hydrogen atom, becomes much less reliable with more complex atoms.
Explain the meaning of the 4 quantum numbers
o Shells
 n = 1, 2, 3 …. 7
 The shell of an electron determines its average distance from the nucleus as well as its energy. So
electrons in shells with higher values are farther away from the nucleus on average.
 The lower the number, more stable
o Subshells
 l = 0, 1, 2 ….
 First shell (n = 1)(l = (s)0)
 Second Shell (n = 2) has two subshells (l = (s)0, (p)1)
 Third Shell (n = 3) has three subshells (l = (s)0, (p)1, (d)2)
 S = Spherical
 P = Dumbbell Shaped
o Orbitals
 m1 = -1, 0, +1
 The magnetic quantum number, or orbital, describes the orientation of the orbital in space.
Roughly, that means it describes whether the path of the electron lies mostly on the x, y, or z axis
of a three-dimensional grid.
 S subshell = 0
 P subshell = -1, 0, +1
 D subshell = -2, -1, 0, +1, +2
o Spin
 Each orbital contains two electrons; one with a positive spin and one with a negative spin.


1
1
2
2
+ ,-
Write electron configurations and orbital notation for elements
o Electron Configurations (see figure on last page)





abc
 a = level
 b = subshell
 c = number from left of subshell beginning
o Orbital Notation
 Fill all shells before adding additional electrons to a single electron placement.
Explain exceptions to the Aufbau principle
o States that when building up the electron configuration of an atom, electrons are placed in orbitals,
subshells, and shells in order of increasing energy.
Define effective nuclear charge, ionization energy, electron affinity
o Effective nuclear charge
 The amount of pull on the electrons due to the positive charge of the nucleus,
 The stronger effective nuclear charge, the smaller atomic radius.
o Ionization energy
 The energy required to remove electrons from gaseous atoms or ions.
o Electron affinity
 Electron affinity reflects the ability of an atom to accept an electron.
Calculate effective nuclear charge and atomic size given ionization energy
o Each subshell adds a large amount to the radius when added.
o The stronger the nucleus is when compared to it’s subshell, the smaller the atomic radius; as well as
ionization energy increases.
o So the radius shrinks across rows and increases greatly down columns.
Recall properties of metals, nonmetals, and metalloids
o
Metal

They are solid (with the exception of mercury, Hg, a liquid).


They are shiny, good conductors of electricity and heat.
hey are ductile (they can be drawn into thin wires).

They are malleable (they can be easily hammered into very thin sheets).
N
onmet
als
They are weak
They are dull, poor conductors of
electricity and heat.
They are not ductile (they can be
drawn into thin wires).
They are not malleable (they can be
easily hammered into very thin
sheets).
Metalloids
They are solid, sometimes soft
They are shiny, good conductors of
electricity and heat.
They are ductile (they can be drawn into thin wires).
They are malleable (they can be easily hammered into very thin sheets).
Robert Girouard
Chapters 8-9 Bonding
Ionic Bonding- An ionic bond is a type of chemical bond formed between two oppositely charged ions. Ionic bonds
are formed between a cation, which is usually a metal, and an anion, which is usually a nonmetal. 1 atom loses
electrons easily, the other has a high electron affinity. Ex. NaCl
Coulombs law used to calculate energy of bond. E=2.31*10-19 j*nm (Q1Q2/r)
Covalent Bonding- A chemical bond that involves the sharing of pairs of electrons between atoms, usually
nonmetals. For many molecules, the sharing of electrons allows each atom to attain the equivalent of a full outer
shell, corresponding to a stable electronic configuration. Low mp/bp, poor conductors. Ex. H2O
Octet Rule- Atoms tend to gain, lose or share electrons so as to have eight electrons in their outer electron shell
(Ex: C, N, O, F,) There are many bonding situations where it does not apply.Exceptions:
 Electron Deficient atoms, Not enough to fill orbitals Ex:(B, Be)
 Elements can (sometimes) accommodate more than 8 electrons, (Si, P, S, Cl, Br)
Lattice Energy -The lattice energy of an ionicsolid is a measure of the strength of bonds in that ionic compound.
Lattice energy may also be defined as the energy required to take gaseous ions and pack them together to form an
ionic solid. Lattice Energy is a state Function. Larger ions= Larger Energy
1.
In which of the following species
does the central atom not form
sp2 hybrid orbitals?
a) SO2
b) BF3
c) NO3
d) SO3
e) PCl5
How to draw Lewis Structures
Valence electrons- The electrons in the last shell or energy level of an atom.
They show a periodic pattern. The valence electrons increase in number as you go across a period.
They are the electrons that get most involved in chemical reactions
Electronegativity- The ability of an atom in a molecule to attract shared electrons to it. Electronegativity increases
left to right across periods and decreases down a group. Highest 4.0 Fluorine, lowest 0.7 Cesium.
Electronegativity in relation to Bond Polarity: The Larger the Difference, The more Polar
Difference in Electronegativity: Zero = Covalent Bond, Moderate = Polar Covalent, Large = Ionic Bond
Example- Bonds in order of Polarity
H-H < S-H < Cl-H < O-H < F-H
(EN. values)
(2.1, 2.1) (2.5, 2.1) (3.0, 2.1) (3.5, 2.1) (4.0, 2.1)
Dipole Moment- When a molecule has a center of positive charge and negative charge. This is illustrated by an
arrow pointing to the negative charge center with the tail indicating the positive center.
Formal Charge- The difference between the number of valence electrons on the free atom and the number of
valence electrons assigned to the atom in the molecule.
To determine Formal charge:
1. The number of valence electrons on the free neutral atom (which has zero net charge)
2. The number of valence electrons belonging to the atom in a molecule.
3. Then Compare the numbers, if same number formal charge is zero, if -1 charge is -1, etc.
Formal Charge = (number of valence electrons on the free neutral atom)-(number of valence electrons belonging to
the atom in a molecule)
-The sum of formal charges must equal the overall charge.
The best Lewis diagram is the one with the lowest formal charges and all octets satisfied
Resonance Structures- When one molecule can have multiple valid Lewis structures, the structure is then given
by the average of all valid structures.
Enthalpy of reaction from bond energies2. Which of the following compounds
The bond energy (kJ) for H2, F2, and HF are 436, 158 and 568 kJ
contains both ionic and covalent bonds?
respectively, calculate the enthalpy (energy) of the reaction,
H2(g) + F2(g) = 2 HF
a) H2O2
Solution
b) CH3Cl
Based on the bond energies given, we have
c) C2H3OH
H2 > 2H
D = 436 kJ/mol
d) NaNO3
F2 > 2F
D = 158 kJ/mol
2H + 2F > 2HF H = -568*2 kJ/mol
Adding all three equations and energies leads to
H2(g) + F2(g) = 2 HF
DH = -542 kJ/equation
VSEPR-(Valence shell electron pair repulsion)- the structure around a given atom is determined principally by
minimizing electron pair repulsions. Bonding and non-bonding pairs are positioned as far apart as possible. BeCl2
:Cl--Be--Cl: (pretend there’s more dots) the shape that allows the electron pairs to be farthest apart is linear (180*)
BF3, The arrangement that minimizes the repulsions is at 120*, so trigonal planar.
Hybrid Orbitals- Hybrid orbitals are the result of a model which combines atomic orbitals on a single atom in ways
that lead to a new set of orbitals that have geometries predicted by the VSEPR model, Hybridizations: 2 electron
pairs= sp, linear 3 pairs=sp2, 4 pairs=sp3, 5 pairs=dsp3, 6 pairs= d2sp3
Sigma Bonds- a single bond between two elements.
Pi Bonds- bonds that are formed alongside sigma bonds to create double( 1,1) or triple bonds(1,2)
Electron Configuration- P-1s2 2s2 2p6 3s2 3p3, [Ne] 3s2 3p3. F-, 1s2, 2s2, 2p6
Sr2+, 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
Q1: E PCl5 forms sp3 orbitals instead of sp2
Q2: DNaNO3 contains an Na+ ion and an NO3- polyatomic ion, which is held together by covalent bonds
Tyson O'Ham
Chapter 8-9, Bonding.
- Define Ionic And Covalent bonds
Covalent bonds are created when atoms share a pair or several pairs of electrons to fill their valence electron shells. Ionic
bonds form when the electronegativity difference between two atoms is such that the most electronegative takes an electron
from the lesser. This process forms charged ions that are strongly attracted to one another. Covalent bonding forms discrete
molecules, whereas ionic bonds form a crystal matrix of closely packed ions. Ionic bonds are stronger than covalent bonds, and
form substances that conduct heat and electricity poory except in molten state. Covalent molecules conduct poorly as well.
-Define lattice energy and discuss those aspects which affect lattice energy
Lattice energy represents the force of attraction between ions in an ionic crystal matrix. The principle factors that determine
lattice energy are distance between the centers of each ion and the ionic charge. lattice energy can be measured with the
formula E ion pair = 2.31 x 10^-19 J · nm (Q1·Q2)/r, where r is the distance between ions, Q1 is the charge of the first ion, Q2
the second. The charges have more effect on the overall lattice energy than the distance between ions does, as they scale
multiplicitavely with each increase.
-Draw lewis structures
Lewis Structures are representations of atoms, either alone, in a molecule, or in an ionic compound. They consist of dots that
represent lone or paired electrons, filled in so that each side of the symbol has a single electron before one has a pair, the
Atomic symbol for the element in the center, lines representing covalent bonds, multiple lines in a bond represent pi bonds.
brackets with a charge represent an ion. Ex:
-Define and identify valence electrons
Valence electrons ere those in the outermost energy level of the atom. They are the ones that pair up or are taken/given away to
form bonds. Ex: sodium has one valence e- because its outermost energy level is 3 and its only e- in n=3 is 3s^1.
-Explain the octet rule
The octet rule states that a given atom will attempt to achieve a complete valence shell of electrons by bonding if it does not
already posess a full octet. A full valence shell is usually four pair of electrons, or eight electrons, hence the moniker octet rule.
-List the Elements that work well with the octet rule
All elements in the s and p block with valence energy levels below n=4 excluding H, He, B, and Be work well with the octet
rule
-Write the electron configurations for ions
Ion electron configurations differ from their base atom. This is because an ion is an atom that has gained or lost electrons. for
atoms in the s and p block, you simply add the amount of electrons an anion has gained or subtract the total a cation has lost
from the usual number and write the electron configuration with tha number of electrons. however, for d block elements, you
always take electrons from the outermost shell and not the highest energy level first, except in cases where it makes the ion
more stable. EX:
Na: 1s2 2s2 2p6 3s1 | [Na]+: 1s2 2s2 2p6 || O: 1s2 2s2 2p4 | [O]2-: 1s2 2s2 2p6 || Sc: [Ar] 4s2 3d1 | [Sc] 2+: [Ar] 3d1 ||
Fe: 1s2 2s2 2p63s2 3p6 4s2 3d6 | [Fe]2+: 1s2 2s2 2p6 3s2 3p6 4s1 3d5
-Define and predict bond polarity using electonegativity
Bond an molecular polarity are functions of differences in electronegativity, and they show partial charge in a bond or
molecule. If a bond is polar then the electons being shared between the bonded atoms are closer to the more electronegative
atom. Nonpolar bonds share the electrons equally. bonds with an EN difference less than .4 are considered nonpolar and bonds
with an electronegativity difference greater than .4 but less than 1.7 are polar. 1.7, the bonds become ionic. EX: A N-F bond (N
has EN =3.0 and F =4.0) would be polar because the EN difference is one. An N=N bond (EN values are the same for the same
atom) would equal 0, and therefore be nonpolar. All diatomic elements are nonpolar.
-Define Dipole-Dipole and formal charge
Dipole-dipole moment is the intermolecular force that results from polar bonds in a molecule forming in partial charges on
different ends of said molecule. This results in attraction between the polar molecule and other charged molecules. drawn with
arrow pointing towards negative side. It is important to note that not all molecules with polar bonds are polar molecules too.
the vectors of the dpoles can geometrically cancel out and result in a nonpolar molecule. (EX: CH4.) Formal charge is used to
determine the most valid form of a molecule or polyatomic ion with multiple configurations. Formal Charge is calculated as
follows: FC = valence number of electrons –(total # of lone electrons around the atom + # of bonds connected to that atom).
The criteria for form validity are that a) formal charges are as close to 0 as possible for all atoms, and b) the central atom is
most negative. Additionally, the sum of the formal charges must equal the charge of the molecule in polyatomic ions.
-Write formal charge and use to predict lewis structures
I will now predict the valid lewis structure of HCN using formal charge.
In (1) The FC of H is 0 (1 e- minus 1 bond), C is 0 (4 e- minus 4 bonds) and N is 0
(five e- minus three bonds plus two paired electrons). In (2), H is 0, c is -1, and n is -1. (1) is the more valid configuration
because it has FC's all equal to zero.
-Draw resonance structures
Resonance structures are formed when the average amount of bonds across a molecule is not a whole number because there
aren't enough electrons to spread evenly around homogenous terminal atoms. these are represented by drawing the multiple
configurations possible in the resonance structure. EX: SO3
-Explain octet rule exceptions
Atoms with valence energy levels greater than n=3 can exceed the amount of valence electrons they coodrinate by filling the d
shell out, having a maximum possible number of bonded electons up to 9. EX: S can form SF6 and P can form PF5, giving
them 12 and 10 valence electrons respectively.
-Explain the basis of VSEPR
VSEPR is used to determine the shape of molecules based on the position of bonds, lone pairs, and number of terminal atoms.
It is useful in determining bond angles and orbital hybridisation. essentially moleculat structures are broken up into categories
based on electron pairs off the central atom, then subdivided further by the number of bonding pairs(multiple bonds still count
as one) and lone pairs there are. the basic forms are linear for one and two electron pairs, trigonal planar for three, tetrahedral
for four, trigonal bipyramidal for five, and octahedral for 6. for every lone pair of the pairs off the central atom you eliminate
one of the members of the geometry, creating a different and often unique shape.
-Use VSEPR model to predict bond angles and geometry
EX: N2 has one electon and one bonding pair, so it will be straight, meaning a bond angle of 180 degrees. H2O has four
electron and two bonding pairs, so its base geometry is tetrahedral, but the actual molecule is bent, with a bond angle of about
109.5 degrees. (the normal tetrahedral bond angle is changed by the lone pair orbitals, which expand further than atoms.)
Methane has 4 pair as well, but all are bonding pairs. this means that its geometry is tetrahedral.
-define and explain hybrid orbitals
Hybrid orbitals are bonding orbitals in covalent molecules that do not adhere to the usual orbital shapes (spherical, barbell,
etc.) They instead use electrons from different sublevels to make new geometric configurations for bonding. the number of
bonding sites and lone pairs off an atom determine the hybridization. for each of these count a letter to name the hybridization:
s, p1-3, d1-2. EX: carbon in CH4 would have sp3 hybridization because of the 4 bonding pairs.
- Define and explain pi and sigma bonding
Sigma bonds represent single bods, and the portion of multible bonds that make physical contact with the other atom's electron
cloud. pi bonds are bonds in multiple bonds that do not physically contact each other. EX: a N-H bond would have one sigma
bond while a N=N bond would have one sigma and two pi bonds.
Questions:
-What is the orbital Hybridization of P in PI5?
a. sp3 b. dsp3 c. d2sp3 d. d2sp2 e. none of these
-Which bond is most polar
a. H-F b. H-Cl c.H-Br d. H-I e H-O
Tony Yun
Chapter 10 IMF, Solids & Liquids
Objective: Define Intermolecular Forces

Intermolecular forces are forces between the molecules. The strengths may vary but they are weaker than ionic or
covalent bonding. Intermolecular forces can be used to determine many physical properties like boiling/melting
points.
Objective: Distinguish between the three intermolecular forces

Ion- Dipole Forces
 Between an ion and polar molecule
 Magnitude of attraction increases as either the charge of the ion or magnitude of dipole moment
increases.






Positive ions are attracted to negative end of polar molecule and negative ends are attracted to
positive ends of polar molecule.
vander Waals Forces
o Dipole-Dipole Forces
 Molecules with dipole moments attracting each other electrostatically by lining up so that the
positive and negative ends are close to each other.
 Between two polar molecules
 Weaker than ion-dipole forces
o London Dispersion Forces
 Forces that exist among noble gas atoms and nonpolar molecules.
 Forces are created by instantaneous dipole moment which then causes induced dipole on another
molecule and thus the attraction between two molecules or atoms.
 Polarizability
 Shape of molecule affects LDF strength (more surface area, more contact = larger forces)
Hydrogen bonding
 Unusually strong dipole-dipole forces among molecules in which hydrogen is bound to a highly
electronegative atom (i.e. oxygen, nitrogen, or fluorine)
 Strongest of the van der Waals Forces
 Objective: Interpret a heating curve
Heating curve is a graph detailing how much energy is needed to change state when the pressure is held constant.
Plateaus- The plateaus in the graph are at which the change takes place (melting, freezing, condensing, boiling)
Rises- The rises in the graph shows the increasing temperature and also the state (solid, liquid, gas)
*(A picture of a heating curve for water can be found in the textbook on page 490, figure 10.42)*
Objective: Calculate change in heat of vaporization


Can be calculated if two vapor pressures and two temperatures are measured.
Claussius-Clayperon Equation :
R = 8.314 J/K mol
ln P = natural log of vapor pressure
T = temperature in Kelvin
C = Constant specific to substance
Objective: Interpret phase diagram




A phase diagram represents the phases of a substance as a function of temperature and pressure.
The lines represent temperature and pressure at which two phases are at equilibrium
The triple point is where all three lines intersect, where a solid, a liquid and a gas are all in equilibrium.
The critical point is the temperature and pressure at which a gas no longer liquefies even though the pressure is
forcing molecules closer together.
*(A picture of a phase diagram for water can be found in the textbook on page 493, figure 10.47)*
Multiple Choice Questions: Refer questions 1 and 2 to the following phase diagram.
h
f
d
b
c
a
g
e
g
t
1.) The point at which the solid, liquid, and gas phases exist simultaneously is point
A.) a B.) b C.) c D.) d E.) e(answer: (a), the triple point, where 3 lines intersect)
2.) Vaporization could take place
A.) Between points (a) and (b) B.) Between points (e) and (a) C.) Between points (a) and (d)
D.) Between points (a) and (h) E.) Above point (d)(answer: (c), the vaporization curve)
Kristina Nguyen
Unit 10: Intermolecular Forces, Solids, & Liquids
Solid
Definite shape & volume
Virtually incompressible
Does not flow
Diffusion occurs slowly
Particles are very close
Strong attractive forces
Possess a crystalline structure
with vibrational motion
Condensed phase
Liquid
Indefinite shape but definite
volume
Virtually incompressible
Flows readily
Diffusion occurs slowly
Particles are not as close
Attractive forces are not as strong
Gas
Indefinite shape & volume
Is compressible
Flows readily
Diffusion occurs rapidly
Particles are very far apart
Condensed phase
Intermolecular Force: a force between molecules
Intramolecular Force: a force that holds together atoms to make a compound or molecule
Dipole – Dipole
Between two polar molecules
London Dispersion
Between 2 non-polar molecules or
atoms
Middle (Strong/Weak)
Weakest
Vocabulary:
Hydrogen
Between 2 non-polar molecules or
atoms where hydrogen atom
connects to F, O, or N
Strongest
1.
2.
3.
4.
5.
Phase Change: change of a substance from one phase to another by heating or cooling, changing average KE OR
changing pressure on a substance
Heat of Vaporization: amount of energy needed to break 1 mol of liquid to a gas; takes place at boiling point
Heat of Fusion: amount of energy needed to break 1 mol of solid particles into liquid ; takes place at melting point
Critical Temperature: the temperature above which a gas cannot be liquefied, regardless of the pressure applied
Critical Pressure: the pressure above which a gas cannot be liquefied, regardless of the temperature applied
Heating Curve:
Cooling Curve:
Vocabulary:
1.
2.
Vapor Pressure: the pressure exerted by a vapor in equilibrium with its solid or liquid phase
Boiling Point: temperature at which a liquid boils( converts from a liquid to a gas)
Phase Diagrams:
1.
2.
Triple Point: conditions at which a solids, liquid, and a gas are all in equilibrium
Critical Point: temperature and pressure at which a gas can no longer liquefy
Amorphous Solids
 Lack a well defined shape, no orderly structure
 Mixtures of molecules that do not stack well or very complex molecules
Crystalline Solids
 Atoms, ions, or molecules are in an ordered well-defined arrangement (have faces at definite angles with each other)
 Orderly stacks of particles
Atomic Solids
Particles at
lattice points
Metallic
Metal atoms
Bonding type
Delocalized
covalent
Properties
Soft to hard,
low to high
melting
points,
excellent
thermal and
electrical
conductors,
malleable and
ductile
Covalent Network
Nonmetal
atoms
connected in a
chain of
molecules
Directional
covalent leads to giant
molecules
Very
hard,brittle,
high melting
points, often
poor
electrical and
thermal
conductors
Ionic Solids
Ions
Ionic
(electrostatic
attractions)
Hard,
brittle,
high
melting
points,
poor
thermal
and
electrica
l
conduct
ors
★Chapter 16 Spontaneity★
Haolin Cai
Defining Entropy in terms of randomness or disorder and state the 2nd Law of Thermodynamics.
 Entropy is a thermodynamic property that measures how close a system is to equilibrium and also
measures the chaos/disorder in the system.
 2nd Law of Thermodynamics- There is a natural tendency to become more disordered (lower to higher
entropy)
 When ∆G is Negative, rxn is spontaneous (∆S)
 When ∆G is Positive, rxn is not spontaneous, in the other direction (-∆S)
 When ∆G is Zero, rxn is at equilibrium
Predicting the sign of entropy of a given process and state the 3rd Law of Thermodynamics.
 2 H2(g) + O2(g) → 2 H2O(g)
There are 3 moles on the reactant side and only 2 on the product side. The change in entropy will be
negative, because the in the reaction the # of moles is reduced.
C6H12O6(s) + 6 O2(g) → 6 CO2(g) + 6 H2O(g)
On the reactant side there is a solid and a gas, on the product side both are gases. The change in entropy
is positive, because there is more disorder in two gases than in a solid and a gas.
 3rd Law of Thermodynamics- Entropy of a pure crystal at 0k is zero, all the other > zero
 Standard enthalpies of elements in standard states are NOT zero
Describing the Effects of Temperature and State changes on Entropy.
 Entropy in the different states is ranked from lowest Entropy to highest entropy
Solid<Liquid<Gas
 Solids have the lowest Entropy because there are the least amount of ways to arrange the
molecules, so less disorder.
 Larger
substances have more Entropy (moles)
−∆𝐻
 ∆SSurrounding=
; The sign of ∆SSurrounding depends on the direction and magnitude of the Heat transfer
𝑇
Calculating ∆ S for a reaction using a table of absolute Entropies.
7
 C2H6 (g) + O2(g) → 2 CO2 (g) + 3 H20 (l)
2
Solve for the ∆S of the combustion 1 mole of C2H6
∆S C2H6 (g) =229.5 J/Kmol
Formula: ∑products - ∑reactants
=∆S
7
∆S O2(g)= 205 J/Kmol
[2(214)+3(70)]-[1(229.5)+ (205)] = 309 J/Kmol
2
∆S CO2 (g) =214 J/Kmol




∆S H20 (l)=70 J/Kmol
These are the coefficients
(These Enthalpies are given in a table in the back of the textbook)
Defining Free Energy in terms of Enthalpy and Entropy, Explain the relationship of the sign of ∆ G and spontaneity
of the reaction.
 Gibb's Free Energy- can be known as "free enthalpy" and is the amount of energy available to do work at
a given temperature and pressure (∆G)
 ∆G = ∆H - T∆S
 ∆G = 0 = equilibrium
 ∆G < 0 = spontaneous reaction
 ∆G > 0 = non-spontaneous reaction
Calculating ∆G using the table of ∆G formation for the reactants and products.
7
 C2H6 (g) + O2(g) → 2 CO2 (g) + 3 H20 (l)
2
Solve for the
∆G of the combustion 1 mole of C2H6
∆G C2H6 (g) =-32.9 KJ/mol
Formula: ∑products - ∑reactants
=∆S
7
∆G O2(g)= 0 KJ/mol
[2(-349)+3(-237)]-[1(-32.9)+ (0)] = -1376.1 KJ/mol
2
∆G CO2 (g) =-349 KJ/mol




∆G H20 (l)=-237 KJ/mol
These are the coefficients
(These Enthalpies are given in a table in the back of the textbook)
Describing the conditions of standard state from standard free energy.
 Standard Free Energy is the free energy change that occurs when reactants in their standard states turn
into products in their standard states
Interconvert ∆G with K for a reaction.
 2PbS(s) + 3O2(g)  2PbO(s) + 2SO2(g) ΔH=-844KJ ΔS=-165J/K
ΔG=ΔH-TΔS
=-844000-(298*-165)
=-794.83kJ
Describing the relationship between ∆G and work.
 Free energy describes the type of energy available from a reaction for the performance of work.
 When ∆G is negative, it loses energy to the surroundings and work is being done by the system on the
surroundings. When ∆G is positive, it gains energy from the surroundings and work is being done by the
surroundings on the system.
Calculating Free energy change for a reaction at non standard conditions.

2H2S(g) + SO2(g) = 3S(s) + 2H2O(g)
Calculate ∆G under the following conditions:
T = 298 K ΔG° = -90.0 kJ/mol
Pressure of H2S = 1.54 x 10-3 atm
Pressure of SO2 = 1.82 x 10-1 atm
Pressure of H2O = 2.80 x 10-1 atm
 ΔG = ΔG° + RT lnQ
Q=
P(H2O)^2
P(H2S)^2P(SO2)
= 1.82e3
ΔG = -90.0 kJ/mol + 8.314 J/K/mol*298 K x 1 kJ/1000J x ln(1.82e3)
= -90.0 + 18.6 kJ/mol
= -71.4 kJ/mol
Predicting how ∆G changes with temperature given the signs of ∆H and ∆S.




∆G is spontaneous at all temperatures when ∆H is NEGATIVE and ∆S is POSITIVE
∆G is spontaneous at HIGH temperatures when ∆H is POSITIVE and ∆S is POSITIVE
∆G is spontaneous at LOW temperatures when ∆H is NEGATIVE and ∆S is NEGATIVE
∆G is NEVER spontaneous at any temperatures when ∆H is POSITIVE and ∆S is NEGATIVE
Estimate ∆G at any given temperature given ∆H and ∆S.
 ΔG=ΔH-TΔS
4NO(g)2N2O(g) + O2 ; ∆H=-199.5 kJ/mol; ∆S=-198.2 J/mol at 250°C
ΔG=(-199500)-(523)(-198.2)=-95841.1 J/mol = -95.841 kJ/mol
Multiple Choice Problems
1. Which has the greatest change in Entropy?
a) The phase change from a solid to a liquid. b) The phase change from a liquid to a gas.
c) The phase change from a solid to a gas. d) The phase change from a gas to a liquid.
Answer: C
2. How can you solve for ∆G°?
a) ΔG°=ΔH°-TΔS°. b)Hess' Law and known reactions. c) ∆G°Reaction=∆G°Products-∆G°Reactants using the appendix values
d) All of the answers stated before.
Answer: D