Download AP CHEMISTRY - An Incomplete List of Topics

AP CHEMISTRY - An Incomplete List of Topics
KINETICS (chap 12)
Rate
rate law
rate constant
Rate determining step
half life
mechanism
Integrated rate laws
overall order
Intermediates
Determining order
-graphically
-from rate data (with and without showing all the work the long way
-from the mechanism (RDS)
Be able to write rate laws and be able to solve for any possible variable.
Know the units of the rate constant.
Be able to calculate the rate of appearance/disappearance of a molecule in a rxn based on another’s rate via mole ratios.
Be able to relate the inte
Know the relationship between activation energy, temperature and the rate constant - Arhenius equation.
EQUILIBRIUM (chap 13, 16)
Write equilibrium expressions Kc, Kp, Ka, Kb, Ksp
Know when prods or reacts are favored.
Apply le Chatelier's principle – particularly it’s impact on K or the conc of a molecule after an add/loss of another
molecule or a temp change.
Be able to use H in le Chatelier's principle and K.
Solve I.C.E. problems. Also know how to do ICE if your given amounts for two molecules and not just one.
Interconvert between Kp and Kc
Know solubility equilibria, Ksp. Realizing that the smaller the Ksp the less soluble the ions. The [ion] is the value of the
solubility. Write an ICE table for the dissociation of the ionic cpd because the ratios for “+x” or “+2x” etc change based on the
ionic cpd.
Apply the common ion effect
Know that pH effects solubility via le Chatelier’s principle.
Buffers (a good buffer is one that is used at a pH that is less than 1 pH of the pKa of the acid, and is a weak acid with its
conjugate base)
Be able to solve for anything in the Henderson-Hasselbach equation
ACID - BASE (chap 14, 15)
Types of acid/base - Bronsted-Lowry, Arrhenius, Lewis: (Remember Lewis is the opposite of Bronsted-Lowry)
Know formulas of strong acids and bases
Be able to calculate the Ka for either a strong or weak acid/base dissociation
Know the chemical equation for either a strong or weak acid/base dissociation
Be able to calculate the percent dissociation (this is the same as percent ionization)
Know the math to calculate the pH, pOH, [H+] and [OH-]
Be able to predict which molecule is a stronger acid based on Lewis Structure (see the course guide)
Be able to fully apply the Henderson-Hasselbach eq.
Be able to identify what kind of acid/base (weak/strong) is being used in a pH titration curve.
Furthermore, be able to identify the parts of a pH titration curve (see fig 15.1-5)
Be able to use ICE, acid/base dissociation eq’s, mole ratios and molarities to calc the amount of a molecule in solution
after a volume of acid or base has been added. (The text has good ex’s in sec 15.4) [I don’t expect you to do the whole
titration, just after one addition.]
ELECTROCHEMISTRY (chap 18)
Identify the components of a half cell or a voltaic cell, see sec 17.2. (also from the line equations)
oxid at anode; red at cathode (Redcat)
Be able to calc oxid numbers and identify reduction and oxidation
Be able to add half-rxns and E’s to get the overall rxn and E
Nernst Equation
Calculate the cell potential, E
Use standard reduction potentials to write redox rxns from 2 half rxns both written for a reduction and predict the
course of a rxn
Determine spontaneity from G and in turn from E and K.
AP CHEMISTRY - An Incomplete List of Topics
Balance redox rxns
Be able to calculate/interconvert between G, E and K (remember n = # of moles of e's)
Know what Q is and be able to use it.
Apply Faraday’s law.
THERMOCHEMISTRY (chap6, 17)
units for energy
def's of system and surroundings, and state function
be able to apply Hess' Law to solve for heat of rxn or indiv. heats of formation
know that P and T at standard conditions (not the same standard Temp as gases)
def's of free energy, entropy and enthalpy the what their signs (+ or -) mean for the rxn
calculate free energy, entropy and enthalpy from the free energy, entropy and enthalpy of formation
use G = H - TS, also be able to predict the sign ofG from the sign of H and S
determine spontaneity from G
determine exothermic/endothermic from H
Determine spontaneity from G and in turn from E and K.
Be able to calculate/interconvert between G, E and K (remember n = # of moles of e's)
GASES (chap 5)
Laws - Ideal vs van der Waal's
Laws - combined, Boyles, Charles, Avogadro's and Gay-Lusac's
Laws - Grahams Law and urms
Laws- know the dRT/P to get molar mass and density
Know the assumptions etc. that make a gas an ideal gas vs a real gas
Laws - partial pressure
Gas stoichiometry – esp using mole ratios or pressure ratios or volume ratios
LIQUIDS/SOLIDS AND SOLUTIONS (chap 10, 11)
IMF's - know the def and be able to identify the predominant IMF in a molecule
ionic; network covalent; H-bond; dipole-dipole; LD, London Dispersion;
IMF's and relation to bpt and m.pt
Types of phase change
Be able to label and read a heating/cooling curve and a phase diagram
Know the def’s of boiling pt and melting pt.
Concentration - mass percent; mass fraction ; molarity and molality
Colligative properties - bpt elevation and mpt depression by the presence of a solute
Raolt's Law
Know all solubility rules and that ionic cpds dissociate when dissolved.
Apply mole ratios to dissociation eq’s and calc the concentration of an ion in solution prior to and following a precipitation
STOICHIOMETRY and TYPES OF RXNS (chap 3 and 4)
Molar mass
mole fraction
empirical formula
Percent composition
mole ratios
theoretical yield
Limiting reactant
experimental yield
Avogadro's number
Solubility and solubility rules precipitation rxns
net ionic eq's
Redox - oxidation, reduction, oxidizing agent, reducing agent, oxidation number
Electrolytes - weak strong and nonelectrolytes
Solute and solvent
molarity
NUCLEAR CHEMISTRY (chap 21)
Nuclear transformation equations
Know the types of radiation, their symbol and the corresponding changes in mass, Y, and atomic number Z. (see table in
course guide)
Know how to predict if a nucleus is stable, p:n ratio
AP CHEMISTRY - An Incomplete List of Topics
BONDING (chap 8 and 9)
Be able to calculate the heat of rxn, H, via bond energies
Know the octet rule and when it is violated
Be able to apply Ionization Energies to successive electron removals realizing that when breaking up stability causes the greatest
jump in IE.
Calc the # of valence e's
Be able to identify a polar vs nonpolar covalent bond, also covalent vs ionic.
Be able to rank polar bonds in increasing polarity (greatest dipole) based on the greatest electroneg. difference
Know how to write/name ionic formulas. Know the names of polyatomic ions.
Be able to draw structures and give the geometry
Understand resonance of e's in p-orbitals in -bonds
Know the bond angles - tetrahedral 109; pyramidal 106; bent 104. planar -120
Be able to determine the hybridization state of a central atom
Know the def of a hybrid orbital
Differentiate between excited and ground states
Know how the energy level diagram leads to stability
Know paramagnetic vs diamagnetic
ORGANIC (chap 22)
meth-1C
eth-2C
prop-3C
but -4C
pent-5C
alkanes and hydrocarbons end with -ane ex propane, C3H8
double bonds end with -ene
propene, C3H6
triple bonds end with -yne
propyne, C3H4
a -yl ending means it is attached to something else
an ether has an O in the middle, ex dimethyl ether is CH3-O-CH3
diethyl ehter is CH3CH2 -O-CH2CH3
Be able to identify the functional groups when given a molecule
alkane
alkene
alkyne
alcohol
ether
aldehyde
ketone
acid
ester
PREDICT PRODUCTS QUESTION #4 ON FREE RESPONSE
Practice and study the various types of rxns and product predictions. Remember you must do 5 of 8
LABORATORY QUESTION ON FREE RESPONSE
Practice and study the lab questions that are in your possession.
Attacking this problem will lead to the greatest success since this problem can be a wide range of topics. Some basic topics to
know would be dissociation eq’s, solubility rules, calcuation of molarity and molality. Also know how to dilute a solution –
M1V1 = M2V2, be aware that the actual am’t of water to be added would be V 2-V1.
Miscellaneous Things to Know
Chem 2 AP: Things to Know for the AP Exam
The 8 common polyatomic ions:
acetate C2H3O21- (also written as CH3COO1-)
ammonium NH41+
chlorate ClO31hydroxide OH1nitrate NO31phosphate PO43-
carbonate CO32sulfate SO42-
Be able to predict properties using the periodic table in terms of number of valence electrons, number of
shells(main energy levels) and net nuclear force(kernel charge).
Metals in Groups IA, IIA, and IIIA will form ions that have charges of 1+, 2+, and 3+ respectively. These
positive ions have only one possible charge and are named by using the name of the element.
Metals(or nonmetals) that have several positive oxidation states must indicate the charge of the ion by
writing a Roman numeral in parentheses following the name of the element.
Negative, monatomic (one atom) ions are named by using the root word of the nonmetal and adding the
suffix
-ide. Nonmetals in Groups VIIA, VIA, VA and IVA will form ions that have charges of 1-, 2-, 3-,
and 4- respectively in binary(two element) compounds.
The patterns used to determine –ate, –ite, and –ide ions’ formulas and charges are:
per- ---------------hypo- ------
-ate ( 1 MORE oxygen atom than the -ate ion)
-ate
-ite ( 1 LESS oxygen atom than the -ate ion)
-ite ( 2 LESS oxygen atom than the -ate ion)
The number of oxygens in the -ate ion can be determined by using the generalizations based on the "Slivka
square" of elements on the periodic chart. The charges of the –ate and –ite ions are usually the same as that
of the –ide ions.
The patterns used to determine the names and formulas of acids are:
Binary aqueous acid compounds (recognized because hydrogen is the first element that is combined with an
ide ion) are named by using the pattern: hydro- (root word of negative element) -ic acid
Three-element acid compounds (hydrogen with a polyatomic ion) are called ternary or oxyacid compounds.
Use these patterns to name the acids:
hydrogen per -----hydrogen
-----hydrogen
-----hydrogen hypo ------
ate
ate
ite
ite
changes to
changes to
changes to
changes to
per ---------------hypo ------
ic acid
ic acid
ous acid
ous acid
There are also 7 elements(HINClBrOF) that usually exist in nature as diatomic molecules, written as H2,
I2,etc.
1 mole contains 6.02 x 1023 particles and has a mass equal to the molar mass (sum of atomic weights)
expressed in grams. At STP, 1 mole of any gas has a volume of 22.4 L.
For conditions other than STP (P = 1 atm and T = 273K), PV = nRT, where R = 0.0821 atm•L / K•mol
Standard pressure can be expressed as 1 atm = 101.3 kPa = 760 torr = 760 mm Hg = 29.9 inches Hg = 14.7 psi
Miscellaneous Things to Know
For solutions, molarity equals the number of moles of dissolved solute per liter of solution.
Miscellaneous Things to Know
Some of the common ions will produce a specific compound when reacting, and the reaction can be reversed to
re-form the ion.
SO42- ↔ SO3
PO43- ↔ P2O5
NO31- ↔ N2O5
MnO41- ↔ Mn2+
SO32- ↔ SO2
CO32- ↔ CO2
NO21- ↔ N2O3
C2O42- ↔ CO2
NH41+ ↔ NH3
OH1- ↔ H2O
ClO31- ↔ Cl1- + O2
Cr2O72- ↔ Cr3+
Example: When Na2CO3(s) is heated, it will decompose to form CO2(g). Solid Na2O will also be formed as
the remainder of the original compound. The Na2O(s) will slowly re-absorb CO2(g) from the air and convert
back into Na2CO3(s). Other carbonates follow this same pattern, as seen in the reaction of carbonic acid;
H2CO3(aq)
CO2(g) + H2O.
Other generalizations include:

There are 7 diatomic elements (HINClBrOF) that exist as molecules in the elemental state.

There are 7 strong acids that ionize completely in water: HCl, HBr, HI, HClO4, HClO3, HNO3 and H2SO4
Write them as the separate ions in reactions, H1+ and the negative anion. All other acids are weak acids,
which are written as molecules in net ionic equations.

All metal hydroxides are strong bases, but only the Group IA hydroxides are very soluble. Write them as
the separate ions in reactions, OH1- and the positive cation.

Metal oxides react with water to form basic solutions. Reactive IA and IIA metals combine with water to
form metal hydroxide solutions and hydrogen gas. Nonmetal oxides react with water to form acidic
solutions.

The conjugate base of a weak acid reacts with water (hydrolysis) to reform the acid. Likewise, the
conjugate acid of a weak base reacts with water to reform the base.

During double replacement (metathesis) reactions, there are no changes in the oxidation states of the
elements involved; only a re-combination of the existing ions occurs. One of the new compounds formed
must be separable from the original mixture, such as a precipitate, a gas, or a nonelectrolyte (water).

When working with water solutions, it is helpful to have a few rules concerning which substances are
soluble, and which will form precipitates. The more common solubility rules are listed below:
1. All common salts of the Group IA(Li, Na, K, etc) elements and the ammonium ion are soluble.
2. All common acetates, nitrates, chlorates and hydrogen carbonates are soluble.
3. All binary compounds of Group VIIA elements(other than F) with metals are soluble, except those of
silver, mercury (I), and lead.
4. All sulfates are soluble except those of barium, strontium, lead, calcium, silver, and mercury (I).
5. Except for those in Rule #1, everything else is insoluble.

During complete combustion reactions of hydrocarbons, CxHy , with oxygen, the products formed will be
CO2 and H2O.
Miscellaneous Things to Know

Splint tests – identify common lab gases by inserting a glowing wooden splint into test tube with unknown
gas. Hydrogen pops (explodes), oxygen makes flame burn more brightly, CO2 extinguishes the flame.

CO2(g) can also be identified by bubbling into limewater, which contains saturated Ca(OH)2(aq) to form a
cloudy precipitate of CaCO3(s).

Single replacement reactions occur only if the free element is more reactive than the element it would
replace in the compound. (check the activity series or Standard Reduction Potential, SRP, chart)

If an element has several possible positive oxidation states, assume complete oxidation occurs during single
replacement reactions and use the highest charge, but during composition reactions use the most stable
oxidation state.

Ionic compounds are formed when a metal and a nonmetal react. Covalent (molecular) compounds are
formed when two nonmetals react. Metals lose electrons and form positively charged ions. Nonmetals try
to gain electrons to form negatively charged ions.

When two nonmetals react, the element with the higher electronegativity (closer to the upper right corner of
the periodic table) will have the negative oxidation number. The percent ionic character (dipole moment) of
the covalent bond increases as the atoms become more reactive metals and nonmetals.

Electronegativities are the relative measure of the attraction of a bonded atom for the shared electrons.

Transition metals typically form colored solutions and colored precipitates.

Transition metal ions act as Lewis acids by accepting an electron pair from anions and/or polar molecules to
form complex ions. The number of attached groups (ligands) is usually twice the charge of the metallic ion.

The total loss of electrons by the oxidation reaction(s) must equal the total gain of electrons by the reduction
reaction(s). First, adjust the coefficients of the reactants and products that contain the oxidized and reduced
elements, then balance any other elements by inspection.

Losing Electrons Oxidation, Gain Electrons Reduction. (LEO the lion goes GER). Substances that are
easily oxidized, like metals, are called strong reducing agents. Strong oxidizng agents, like MnO41-,
C2O42-, and Cr2O72-, are easily reduced.

In electrochemical cells, oxidation always occurs at the anode, and reduction always occurs at the cathode.

When only net ionic equations are used, spectator ions or molecules not involved in the reaction are
excluded.

Net ionic redox reactions are balanced by the following steps:
Step 1:
Write the equation in ionic form.
Step 2:
Write the separate oxidation and reduction half-reactions.
Step 3:
Adjust the coefficients to get an equal number of atoms of all elements EXCEPT oxygen and
hydrogen.
Step 4:
Balance the hydrogen and oxygen atoms by adding H1+ and H2O.
Step 5:
Balance the electric charge on both sides of each half-reaction by adding electrons.
Step 6:
Multiply each half-reaction by an appropriate number to make the electron changes equal.
Step 7:
Combine the two half-reactions and simplify(subtract substances that appear on both the reactant
and product sides of the equation). Check to see if the numbers of atoms of each element and
total charges of reactants equals the atoms and charges of the products.
Miscellaneous Things to Know
Step 8:
If the solution is acidic, you are done. If the solution is basic, add one OH1- for each H1+
present to form water, then add an equivalent number of OH1- to the other side of the equation.
Simplify the number of H2O present.

All elements in the same horizontal row on the periodic table, called a period or series, have the same
number of occupied electron energy shells. For example, all the atoms of the elements from K to Kr have 4
shells.

All elements in the same vertical column on the periodic table, called a group or family, have the same
number of valence (outer-shell) electrons AND the same kernel charge (net nuclear force). The kernel
charge can be calculated by subtracting the number of inner shell electrons from the number of protons.

For the Group IA to Group VIIIA elements, the number of valence electrons and the kernel charge (net
nuclear force) increases from 1 to 8, respectively. The greater the kernel charge, the stronger the pull of the
nucleus on the valence electrons. Therefore, ionization energy, electron affinity, and electronegative tend to
increase with the increasing kernel charge.

Filled and half-filled electron sublevels have extra stability. Elements undergo chemical reactions by
gaining, losing or sharing electrons to obtain an octet (8 electrons) in their valence shell, like the Group
VIIIA inert gas elements.

The valence electrons for Groups IA and IIA elements fill the s sublevel. The valence electrons for Group
IIIA and VIIIA elements fill the p sublevel. The valence electrons for Group B transition elements fill the s
sublevel, and then the d sublevel. The valence electrons for the rare earth elements, also called the inner
transition elements, fill the s sublevel, and then the f sublevel.

Atoms increase in size going down a Group, because the extra electron shells have more energy and occupy
a larger space. (There is more electron-electron repulsion.)

Atoms decrease in size going left to right across a period, because the greater kernel charge has a stronger
pull on the same number of electron shells.

When atoms lose electrons to form positively charged cations, the ionic radii decreases. Typically, an
electron shell is lost in this process. When atoms gain electrons to form negatively charged anions, the
electrons fill an already occupied valence shell. This causes the anion to increase in size.

Ionization energy is the energy required to remove an electron from an atom in the gaseous state.
Nonmetals with higher kernel charges and fewer electron shells have the highest ionization energies. Filled
and half-filled sublevels have extra stability that require higher than usual ionization energy.

Electron affinity is the energy change that occurs when an atom gains an electron. If the anion is more
stable, energy is released during an exothermic change, and the electron affinity is a negative value. If the
anion is less stable, then energy must be absorbed during an endothermic change, and EA value is positive.

Equilibrium conditions can occur when the rate of a forward reaction becomes equal to the rate of the
reverse reaction. If a stress is added to an equilibrium system, it will shift to alleviate the stress and reestablish the equilibrium. (LeChatlier’s Principle). Adding a reactant or product shifts the equilibrium to the
opposite side to use up the excess. Removing a reactant or product shifts the equilibrium to replace the loss.

The equilibrium constant is the ratio of the product concentrations raised to a power of the coefficients from
the balanced equation divided by the reactant concentrations raised to a power of the coefficients from the
balanced equation. NEVER include pure solids or liquids in the equilibrium constant expression.
For the reaction, A 2 (g) + 3B2 (g)
2
K eq =
[AB3 (g)]
[A 2 (g)][B2 (g)]3
2AB3 (g)
Miscellaneous Things to Know

Specific designations have been assigned to various types of equilibrium conditions, but they are all based
on the product to reactant ratio. These include Kc, Kp, Ka, Kb, Kw, Ksp, and Kf .

The pH scale is based on the equilibrium for the ionization of water, Kw = [H1+ ][OH1- ] = 1.00 x 10-14.
pH = -log[H1+]
pOH = -log[OH1-]
pH + pOH =14
Ka • Kb = Kw

Buffers consist of a weak acid/conjugate base pair with a pH <7 or a weak base/conjugate acid pair with a
pH >7.

When a strong acid is added to a buffer, the basic form reacts to neutralize the acid. Then equilibrium is reestablished with the remaining base and the extra acidic form that is produced.

When a strong base is added to a buffer, the acidic form reacts to neutralize the base. Then equilibrium is
re-established with the remaining acid and the extra basic form that is produced.

For kinetics, the rate law is based only upon the concentrations of the reactants, never the products. The
concentration of each reactant is raised to a power, called the rate order, based upon how changing its
concentration effects the rate of the reaction. The rate order can ONLY be determined experimentally and
does NOT depend upon the coefficients from the chemical equation.
A + 3B  AB 3
reaction rate k [A] m[B] n

If changing the concentration of the reactant has NO effect on the rate, then it is zero order. A plot of the
reactant concentration versus time will be linear. If the overall reaction is zero order, the units for k will be
M/time or M•time-1.

If changing the concentration of the reactant has the same effect on the rate, then it is first order. For
example, if the concentration of A is doubled, then rate will double. A plot of the natural log of reactant
concentration (ln[A]) versus time will be linear. If the overall reaction is first order, the units for k will be
1/time or time-1. The half-life rate of decay plots for radioactive isotopes are all first order reactions.

If changing the concentration of the reactant has a squared effect on the rate, then it is second order. For
example, if the concentration of A is tripled, then rate will be nine times faster, or 32. A plot of the inverse
of the reactant concentration (1/[A]) versus time will be linear. If the overall reaction is second order, the
units for k will be 1/M•time or M-1•time-1

The slow step for a reaction mechanism is the rate-determining step. The rate order based upon the slow
step must match those in the actual rate law for the overall reaction. Intermediates are substances formed
during a reaction that are consumed in subsequent steps. Therefore, they do not appear in the net equation
for the reaction.

The activation energy is the minimum energy needed to overcome the energy barrier between the reactants
and products and cause the reaction to occur. A catalyst speeds up a reaction rate by lowering the activation
energy

Reactions tend to form products that have lower enthalpy (energy) or higher entropy (disorder).
Miscellaneous Things to Know

Solids have the lowest enthalpy, with well-defined crystal structures. The solid particles have only
vibrational motion. As heat is added, the solid melts and the liquid particles gain rotational motion. Gases
have the greatest disorder and highest enthalpy.

All exothermic reactions (∆Hº = -) that have an increase in entropy (∆Sº = +) are spontaneous. The value of
Gibbs’ Free Energy, ∆Gº, is negative and Eºcell is a positive value. ∆Gº = ∆Hº – T∆Sº

The normal boiling point of a liquid is the temperature at which the equilibrium vapor pressure of the liquid
is equal to standard pressure (1 atm). In a vacuum, water boils at a lower temperature.

The normal melting point of a solid is the temperature at which the equilibrium vapor pressure of the solid is
equal to standard pressure.

Adding a nonvolatile solute always lowers the vapor pressure of a liquid solvent. This raises the
temperature needed to boil the solution. The freezing point is lowered for the solution, as well as the
osmotic pressure.

Nonpolar molecules are held together by weak dispersion interaction. This occurs when a temporary,
induced dipole is created by the repulsion of the electrons in one molecule by the electrons in another. The
strength of the dispersion forces increases with increasing molar mass. (more protons and electrons)

Hydrogen bonding occurs when hydrogen atoms bonded to N, O, F or Cl in one molecule, attempt to form a
coordinate covalent bond to an unshared pair of electrons on a neighboring molecule.

According to the VSEPR Theory, molecules with 2 areas of electron charge are linear (180º angles), 3 areas
of charge are trigonal planar (120º angles), and 4 areas of charge are tetrahedral (109.5ºC). Also, sp
hybridization forms linear molecules, sp2 hybridization forms trigonal planar molecules, and sp3
hybridization forms tetrahedral molecules. The arrangement of 5 areas of charge is trigonal bipyramidal and
6 areas is octahedral.

Single bonds are formed by the overlap of the electron orbitals directly between two nuclei, called sigma ()
bonds. Multiple bonds are formed by the lateral, side-to-side overlap of two orbitals, called pi () bonds.

Organic compounds consist of mostly C and H atoms held together by nonpolar bonds. The names are
derived from the number of carbon atoms in the main structure.

Methane CH4, ethane CH3CH3, propane CH3CH2CH3, butane CH3CH2CH2CH3, then the prefix
indicates the number of carbon atoms. Pentane C5H12, hexane C6H14, etc.

Organic compounds with carbon-carbon double bonds end in –ene, carbon-carbon triple bonds end in –yne.

Alcohols have an –OH group on the carbon chain and end in –ol. A number is used to locate the position of
attachment on the carbon chain. For example, 2-propanol has three carbon atoms held together by single
bonds with an –OH group attached on the middle carbon atom.

Organic acids have the carboxyl group, -COOH, attached at the end of the carbon chain. The doubledbonded oxygen increases the polarity of the O-H bond and causes some of the H1+ ions to form.

The half-reaction with the more positive Eºred value on the Standard Reduction Potential (SRP) chart will
be reduced and the other half-reaction is reversed and written as the oxidation reaction. The sign of the
Eºred value is also reversed for the Eºox value. The potentials are intensive properties and do not depend on
the number of moles of reactants used. Eºcell = Eºred + Eºox.
Miscellaneous Things to Know

The standard conditions for the Eºred values are at 25ºC, 1 atm of pressure, and 1 M solution concentrations.
The experimental voltage, Ecell, can be predicted by using the Nernst equation. At higher temperatures or if
the value of the reaction quotient, Q, is >1, then the experimental voltage will be lower.

There are several types of graphical representations that you should be able to interpret and explain:
Heating Curve for Water at S tandard Pressure(101.3 kPa)
120
boiling-condensation
Gas
80
Liquid
Temp. ºC
40
melting-freezing
0
S olid
- 40
0
2
8
6
Heating Time in M inutes
4
10
12
Equilibrium Vapor Pressure for Liquids A, B & C
100
A
B
C
75
Pressure
kPa
50
25
0
0
20
40
60
Temp ¼C
80
100
120
Miscellaneous Things to Know
Solubility Curves
Energy Changes During Reaction
Molecular Speed of Gas Molecules
Equilibrium Reactions Ex: PCl3 + Cl2  PCl5
Effect of Catalyst on Activation Energy
uncatalyzed
reaction
P Cl 3
concentration
P Cl
P Cl 3
concentration
5
potential energy
Cl 2
Cl 2
P Cl
time
time
catalyzed
reaction
5
reaction co-ordinate
Miscellaneous Things to Know
Supercooling of Solid During Freezing
Freezing Point Depression of a Solution
Time
Titration of Diprotic Acid with NaOH
Titration of Ca(OH)2 with HCl
pH
0
Titration of Weak Acid with NaOH
EMR Absorbance of Colored Solution
12
A
b
s
o
r
b
a
n
c
X
10
X
8
pH
X
X
6
X
X
mL HCl
X
4X
e
2
Concentr ation, Molarity
0
0
5
10
15
20
millilitres of NaOH
25
30
Miscellaneous Things to Know
Phase Diagrams
Half-life of Radioactive Isotope
Descriptive Colors in Chemistry
Flame Test Colors
Li+ Deep red (crimson)
Sr2+ Red
Na+ Yellow
Ba2+ Green
K+ Violet
Cu2+ Blue-green
Ca2+ Orange-red
Aqueous Ion Colors
Cu+ Green
Cu2+ Blue
Cu(NH3)42+ Dark Blue; produced when ammonia added to Cu2+(aq)
Co2+ Pink CoCl42- Blue (Co2+ with HCl will form a blue CoCl42- complex ) Ni2+ Green
Fe2+ yellow-green (depending on anion)
Fe3+ orange-red
FeSCN2+ Wine-red to dark orange
Mn2+ Pink
MnO41- Purple (Mn w/ +7 oxidation state is purple)
Cr3+ Violet (Cr(NO3)3 to Green (CrCl3)
CrO42- Yellow
Cr2O72- Orange
• Al, K, Li, Mg, Na, Ca, Ba, Sr, Zn are colorless aqueous ions and most of their solid salts are white.
• Transition element ions with partially filled d orbitals tend to release colored light.
Assorted Compounds
F2 Pale-yellow gas
Cl2 Green-yellow gas
Br2 Red-brown liquid
I2 Dark-violet vapor & dark metallic looking solid
S8 Yellow, odorous solid
NO Colorless gas; from reactions of metals and dilute HNO3
PbI2 Bright yellow precipitate
NO2 Brown gas; from reactions of metals and concentrated HNO3
Fe2O3 Reddish brown (rust)
Metallic sulfides Sulfides of transition metals tend to be black Litmus (paper) red in acid, blue in base
Phenolphthalein Colorless (pH<7) to Pink (pH>8 ; when OH- is present)
Hints for Reaction Prediction
Types of equations to know:
1.
decomposition
(DECOMP)
AB → A + B
2.
synthesis
(SYN)
A + B → AB
3.
single replacement
(redox)
A + BC → AC + B or A + BC → BA + C
4.
combustion of a hydrocarbon
(comb)
CH(O) + O2 → CO2 + H2O
5.
double replacement precipitate
MUST WRITE NET IONIC EQUATION!!!
(DR-met)
AB + CD → AD + BC
5.
double replacement neutralization (
MUST WRITE NET IONIC EQUATION!!!
MUST WRITE NET IONIC EQUATION!!!
DR-neut)
acid + base H+ + OH - → H2O
strong acids just H+, strong bases just OHweak acids or bases – write the molecule
5.
double replacement hydrolysis (lysis = cut)
MUST WRITE NET IONIC EQUATION!!!
(DR-hydro)
AB + H2O → metal hydroxide + acid
8.
redox
9.
Coordination compound/ligand type reaction Metalcation + molecule/ion → Mt(molec/ion)2xcharge
total charge
subscript
metal oxides + H2O → basic hydroxide compnounds
nonmetal oxides + H2O → acidic compounds
the nonmetal’s oxidation number is constant
Both oxide reactions are reversible when heat is added to the products
10. Oxides (e;g KO, N2O5)
Hints for predicting the Assigned Reactions:
1.
AB, CD
2.
Burned = Combusted
3.
two reactants with a common “anion” so it cannot be SR or DR
4.
metal oxide and water
5.
acid and base
6.
single “atom” plus single atom
7.
two reactants: molecule plus a metal atom
8.
acid plus base
9.
two reactants: molecule plus molecule {note that mercury (I) is a polyatomic, Hg 22+}
10. Heating an oxysalt produces a metal oxide plus a nonmetal oxide or a metal salt plus oxygen,
More info for reactions……….
Common Chemical Reactions
1.
When two elements react, a combination reaction occurs (think: could any other type of reaction occur?),
producing a binary compound (that is, one consisting of only two types of atoms). If a metal and a nonmetal
react, the product is ionic with a formula determined by the charges on the ions the elements form. If two
nonmetals react, the product is a molecule with polar covalent bonds, with a formula consistent with the normal
valences of the atoms involved. Some pairs of elements may react only slowly and require heating for
significant reaction to occur.
Examples:
K + S8
K2S (ionic)
Ca + O2
CaO (ionic)
Al + I2
AlI3 (ionic)
H2 + O 2
H2O (covalent)
I2 + Cl2
ICl, ICl3, or ICl5 (covalent)
(exact product depends on relative amounts of I2 and Cl2)
(NOTE: The above reactions are not balanced, nor were they intended to be. They, like the others in this
handout, are meant only to show the correct formulae for the reactants and products. You may wish to
balance the reactions in the handout as an exercise.)
2.
Reaction of a metal oxide with water produces a metal hydroxide; that is, a strong base. Reaction of a nonmetal
oxide with water produces an oxyacid in which the nonmetal is in the same oxidation state as in the oxide you
started with. Both of these are combination reactions, and both can be reversed by heating the products. Metal
hydroxides decompose on heating to give the metal oxide and water, and oxyacids decompose on heating to
give water and the nonmetal oxide in the appropriate oxidation state.
Examples:
Na2O + H2O
MgO + H2O
SO2 + H2O
Cl2O5 + H2O
HNO3
Fe(OH)3
3.
NaOH
Mg(OH)2
H2SO3
HClO3
N2O5 + H2O
Fe2O3 + H2O
Reaction of a metal oxide with a nonmetal oxide gives an oxysalt; reaction of a metal hydroxide with a
nonmetal oxide produces a "hydrogen" oxysalt. This is essentially a reaction of the O2- or OH- in the metal
compound with the molecular nonmetal oxide. This combination reaction occurs only if no water is present; in
the presence of water, the nonmetal and metal oxides react with the water to produce acid and hydroxide,
respectively (as shown in (2) above), then these react as in (4) below.
Examples:
CaO(s) + SO3(g)
NaOH(s) + CO2(g)
4.
CaSO4(s)
NaHCO3(s)
Reaction of an acid with a base gives a salt plus water. The cation in the salt comes from the base; the anion
comes from the acid. The base may be a metal hydroxide, a metal oxide, or a weak base such as NH 3. The acid
and/or base may be pure solids, liquids, or gases, or in aqueous solution. The oxidation states of the anion of the
acid and cation of the base normally remain unchanged.
Examples:
Common Chemical Reactions
HCl(aq) + Ca(OH)2(aq)
H2SO4(aq) + Fe(OH)3(s)
NH3(g) + HC2H3O2(l)
Al2O3(s) + HClO4(aq)
5.
CaCl2(aq) + H2O(l)
Fe2(SO4)3(aq) + H2O(l)
NH4C2H3O2(s)
Al(ClO4)3(aq) + H2O(l)
Ammonium salts react with metal hydroxides and oxides in an acid-base reaction to produce ammonia. This is
essentially the reverse of one of the reaction types mentioned in (4) above. Either or both of the reactants may
be a pure material or in aqueous solution.
Examples:
NH4Cl(aq) + KOH(aq)
NH4NO3(s) + CaO(s)
6.
NH3(g) + H2O(l) + KCl(aq)
NH3(g) + H2O(l) + Ca(NO3)2(s)
Reaction of the salt of a weak acid (that is, a compound containing the anion of a weak acid) with a strong acid
produces the weak acid and a salt. This is another example of an acid-base reaction, in addition to the ones
given in (4) and (5) above. The original salt of the weak acid may be either a pure solid or in aqueous solution.
The cation in the salt formed as the product comes from the weak acid salt; the anion in the product salt comes
from the strong acid. In many cases, the weak acid produced is unstable and decomposes to give the oxide of a
nonmetal and water (see (2) above). This is especially true if the nonmetal oxide is a compound of limited
solubility in water such as SO2, CO2, or the nitrogen oxides. The best-known examples of this type of reaction
involve carbonates, bicarbonates, sulfides, and sulfites, but many other examples are known as well. Normally,
these reactions do not involve oxidation or reduction.
Examples:
BaCO3(s) + HBr(aq)
BaBr2(aq) + H2O(l) + CO2(g)
NaHCO3(aq) + H2SO4(aq)
Na2SO4(aq) + CO2(g) + H2O(l)
MgS(s) + HCl(aq)
H2S(g) + MgCl2(aq)
K2SO3(aq) + HNO3(aq)
KNO3(aq) + SO2(g) + H2O(l)
Ca3(PO4)2(s) + HCl(aq)
CaCl2(aq) + H3PO4(aq)
Zn(C2H3O2)2(aq) + HBr(aq)
ZnBr2(aq) + HC2H3O2(aq)
7.
Reaction of solutions of two soluble salts with one another can give a precipitate of an insoluble salt formed by
a double replacement reaction (also called a metathesis). Whether or not a precipitate forms depends on the
exact combination of salts used. To make a prediction as to whether a reaction will take place or not, you must
know the solubility rules for common salts (Ebbing 4/e, page 104; lab manual, Appendix 7). Some
combinations of salts may give oxidation-reduction reactions (see (11) below), but most do not.
Examples:
CaCl2(aq) + K2CO3(aq)
AgNO3(aq) + FeCl3(aq)
but: NiSO4(aq) + MgI2(aq)
CaCO3(s) + KCl(aq)
AgCl(s) + Fe(NO3)3(aq)
no reaction
(NiI2 and MgSO4 are both soluble)
Al(NO3)3(aq) + Pb(C2H3O2)2(aq)
no reaction
(Al(C2H3O2)3 and Pb(NO3)2 are both soluble)
8. Heating an oxysalt produces a metal oxide plus a nonmetal oxide or a metal salt plus oxygen, or some
combination of these two decomposition reactions.
Common Chemical Reactions
Examples:
KClO3(s)
KCl(s) + O2(g)
CaCO3(s)
CaO(s) + CO2(g)
Pb(NO3)2(s)
9.
PbO(s) + NO(g) + NO2(g) + O2(g)
Heating a hydrated material initially causes a decomposition reaction to produce the anhydrous compound and
water. Further heating may yield further decomposition, depending on the material. (See (2) and (8) above.)
Most binary compounds are stable to heat.
Examples:
H2C2O4. 2H2O(s)
H2O(g) + H2C2O4(s); followed by
H2C2O4(s)
CaCl2.
H2O(g) + CO(g) + CO2(g)
6H2O(s)
CaCl2(s)
CuSO4.
H2O(g) + CaCl2(s); followed by
no reaction
5H2O(s)
CuSO4(s)
H2O(g) + CuSO4(s); followed by
CuO(s) + SO3(g) (requires strong heating)
10. Reaction of an element with a compound often gives a single replacement reaction in which a nonmetallic
element can replace a combined nonmetal, and a metallic element can replace a combined metal, or hydrogen
from an acid. As a general rule, a more active (reactive) element will replace a less active (reactive) element
from its compounds. In general (but with many exceptions), the most reactive nonmetals are found to the upper
right in the periodic table, and the most reactive metals are found to the lower left. The order of reactivity of the
halogens is F2>Cl2>Br2>I2. For hydrogen and the more common metals, the order of reactivity (the activity
series) is
Li>K>Ca>Na>Mg>Al>Zn>Cr>Fe>Ni>Sn>Pb>H2>Cu>Hg>Ag>Pt>Au
In these two series, one element can replace another one to its right in the series. Metals to the left of H 2 can
replace H+ from acids. The very reactive metals (Li, K, Na, Ca) can replace H+ from cold water; metals of
intermediate reactivity (Mg, Al) can replace H+ from hot water or steam. Any single replacement reaction can
also be categorized as an oxidation-reduction (redox) reaction.
Examples:
Al(s) + NiSO4(aq)
Al2(SO4)3(aq) + Ni(s)
Fe(s) + HBr(aq)
FeBr3(aq) + H2(g)
Cl2(g) + KI(aq)
KCl(aq) + I2(s)
Na(s) + H2O(l)
NaOH(aq) + H2(g)
Zn(s) + Cu(NO3)2(aq)
Cu(s) + Zn(NO3)2(aq)
but: Ag(s) + HClO4(aq)
Br2(l) + ZnCl2(aq)
Sn(s) + H2O(l)
Pb(s) + CrF3(aq)
no reaction
no reaction
no reaction
no reaction
Common Chemical Reactions
11. Compounds containing one or more atoms in high oxidation states often act as oxidizing agents; compounds
containing atoms in low oxidation states often act as reducing agents. For most elements, the (old) group
number of the atom in the periodic table gives the highest oxidation state possible for that element. For
nonmetals, the lowest oxidation state possible is given by the (old) group number minus eight. Elemental metals
most often act as reducing agents (they are oxidized); nonmetals frequently act as oxidizing agents (they are
reduced).
For the representative elements (i.e., those in the first two and last six columns of the periodic table), oxidation
states most often are two units apart. For example, Sn forms Sn(II) and Sn(IV); Br forms Br1-, Br(I), Br(III),
Br(V), and Br(VII). For the transition elements, (i.e., those in the "center" ten columns of the periodic table),
oxidation states are often one unit apart, but can be in almost any relationship to one another. For the transition
elements, the common oxidation states (charges on their ions) must be memorized. For example, Fe forms Fe 2+
and Fe3+; Cu forms Cu+ and Cu2+, etc. Some of the transition elements form oxyanions as well as cations. For
example, Mn forms Mn2+, Mn3+, MnO42-, and MnO4-; Cr forms Cr2+, Cr3+, CrO42-, and Cr2O72-.
Any atom in its highest possible oxidation state can only act as an oxidizing agent; any atom in its lowest
possible oxidation state can only act as a reducing agent. Atoms in intermediate oxidation states can be either
oxidized or reduced; that is, they can act as either reducing or oxidizing agents. Some of the oxidizing agents
most commonly encountered are MnO4-, CrO42-, Cr2O72-, HNO3, H2O2, and the halogens. Some of the more
common reducing agents are elemental H2, metals, carbon, and I-.
In predicting products of oxidation-reduction reactions, don't forget their name--oxidation and reduction must
occur simultaneously! It is impossible for oxidation to occur without reduction or vice versa.
Examples:
Sn2+(aq) + F2(g)
Sn4+(aq) + F-(aq)
Mn2+(aq) + BiO3-(aq)
Bi3+(aq) + MnO4-(aq)
(note that the Bi is in its highest possible oxidation state in BiO3-)
K(s) + P4O10(s)
K3PO3(s)
(note that P is reduced from P(V) to P(III))
MnO4-(aq) + I-(aq)
Mn2+(aq) + I2(aq)
CuS(s) + HNO3(aq)
Cu(NO3)2(aq) + S8(s) + NO2(g)
(note S2S0 and N(V)
N(IV))
Fe2O3(s) + C(s)
CO2(g) + Fe(s)
Reaction Prediction
PRACTICE:
REACTION PREDICTION – WRITE THE BALANCED EQUATION
1. Powdered magnesium carbonate is heated strongly
2. Solid iron (II) sulfite is heated strongly
3. A strip of zinc is placed in a solution of nickel (II) nitrate
4. Solutions of strontium nitrate and sodium sulfate are mixed
5. Excess concentrated potassium hydroxide solution is added to a solution of nickel (II) chloride
6. A solution of sodium phosphate is added to a solution of aluminum nitrate
7. Solutions of silver nitrate and sodium chloride are mixed
8. Ethanol is burned in oxygen
9. Solutions of tin (II) chloride and iron (III) chloride are mixed
10. Ammonia gas is mixed with hydrogen chloride gas
11. Sulfur dioxide gas is bubbled into distilled water
12. Solid potassium carbonate is added to 1.0 M sulfuric acid
13. Concentrated ammonia solution is added to copper (II) sulfate solution
14. Excess ammonia is added to a solution of zinc nitrate
15. An excess of hydrochloric acid is added to solid zinc sulfide.
16. Hydrogen chloride gas is bubbled through a solution of potassium cyanide
17. Ammonia gas is mixed with hydrogen chloride gas
18. Excess concentrated ammonia solution is added to a solution of nickel (II) sulfate
19. A drop of potassium thiocyanate solution is added to a solution of iron (III) nitrate
20. Solid sodium hydrogen carbonate (sodium bicarbonate) is strongly heated
21. A solution of hydrogen peroxide is exposed to strong sunlight
22. Solid iron (II) sulfite is heated strongly
23. A solution of copper (II) chloride is added to a solution of sodium sulfide
Reaction Prediction
24. Solutions of strontium nitrate and sodium sulfate are mixed
25. A solution of sodium iodide is added to a solution of lead (II) acetate
26. Solid lead (II) carbonate is added to a 0.5 M sulfuric acid solution
27. Sulfur in its standard state is burned in air
28. Ethyne (acetylene) is burned in air
29. Propene is burned in air
30. Excess hydrobromic acid solution is added to a solution of potassium hydrogen carbonate