Chemistry 30 Review
UNIT A – THERMOCHEMICAL CHANGES










Enthalpy change is also referred to at the net amount of energy
All energy comes from the sun
Know the photosynthesis and cellular respiration reactions
Know the phase of water in an open and closed system
Calorimetry is used to calculate the enthalpy/energy stored in chemicals
o Q=mc∆t (in units of joules)
o Sort your information into chemical information and calorimetry information
Calorimeter
Chemical System
-Q
=
∆H
-(mc∆t)
=
n∆rH
o calorimeter calculations are based on an isolated system, but in reality,
they are a closed system
In an experiment designed to compare the molar enthalpy of combustion for two
or more different fuels
o manipulating variables: different fuels
o responding variables: temperature change
o controlled variables: type of calorimeter
o amount of fuel doesn’t necessarily need to be controlled because you are
trying to determine the molar enthalpy change
Kinetic energy is related to a temperature change vs. potential energy that is
related to chemical reactions/bond energy/enthalpy
o Know the energy changes that occur during a reaction (conversion of
kinetic into potential and back to kinetic) and apply this to an energy
potential diagram
o Bond breaking requires energy and bond making releases energy, but all
reactions involved bond breaking and bond making. An exothermic
reaction is defined when the energy required for bond breaking is less
than the energy released from bonds forming. An endothermic reaction is
defined when the energy required for bond breaking is greater than the
energy released from bonds forming.
If reaction is exothermic, ∆H (enthalpy) is negative and is written on the product
side
If reaction is endothermic, ∆H (enthalpy) is positive and written on the reactant
side
Enthalpy change (∆H) is very specific for the moles in the particular balanced
reaction
o For example, if the amount of moles is doubled, the enthalpy change (∆H)
should double as well


Molar enthalpy change (∆rH) is the enthalpy for 1 mol of that particular chemical
in a specific reaction
For specific amounts problems that deal with enthalpy change, create a chart of
specific information and balanced reaction information
Specific
Balanced Reaction
n=
n=
∆H =
∆H =
∆H = n∆rH


Potential energy diagrams
o Step up = endothermic
o Step down = exothermic
Use enthalpies of formation to calculate the enthalpy for a balanced reaction
o This method is used when all you know is the balanced reaction
o Remember that phases of the chemicals can change the enthalpies!!!
H  n   r H   n  rH 
products



reactants
Use Hess’s Law to calculate the enthalpy for a reaction when a list of
intermediate reactions are given
o Adding up intermediate reactions to get the net reaction
Define activation energy as the energy barrier that must be overcome for a
chemical reaction to occur
Draw potential energy diagrams that show the difference between a catalyzed
and un-catalyzed reaction
o Explain that catalysts increase reaction rates by providing alternate
pathways from reactants to products, but have no effect on the enthalpy
change for a reaction
o Should be able to determine/calculate the activation energy (Ea) and the
enthalpy change (∆H)for the forward and reverse reactions
o Activation energy is always the difference between the reactants (recall
the products are really the reactants when the reaction is reversed) and
the top of the “hill”
o Activated complex is at the top of the hill and has the most potential
energy
UNIT B – ELECTROCHEMICAL CHANGES









Metals lose electrons to form positive ions and non-metals gain electrons to form
negative ions
Know the difference between oxidation, reduction, oxidizing agent, and reducing
agent
o Oxidation: Theoretically, the loss of electrons. Historically, a reaction with
oxygen
o Reduction: Theoretically, the gaining of electrons. Historically, the
purification of metal ore into pure metal which caused a reduction in
volume
Know the difference between a regular balanced reaction, a ionic balanced
reaction (all aqueous chemicals that are ionic and soluble in water will break into
their ions), and a net ionic reactions (ionic reaction with the spectator ions
removed)
Need to be able to determine if a reaction is redox or not → use oxidation
numbers!
The easiest way to determine if an atom (which is usually part of a molecule) is
undergoing reduction or oxidation is to find the change in its oxidation number
o If the oxidation number becomes more positive, it is losing
electrons→oxidation
o If the oxidation number becomes more negative, it is gaining
electrons→reduction
o Know the rules for identify/calculating oxidation numbers
Disporportionation is when the same atoms undergo both oxidation and reduction
in the same reaction
Need to be able to balance half-reactions in acidic conditions and then use the
balanced half-reactions to add together to get the balanced redox reaction (need
to ensure you have the same amount of electrons gained and lost before adding
the half-reactions together)
o Remember that the half-reactions on pg. 7 in the data book can be read
forward (ie. left/reactants to right/products) to be a reduction reaction. And
the half-reactions can be read backwards/reverse (ie. right/reactants to
left/products) to be an oxidation reaction
We also balance redox reactions using oxidation numbers
o Used for questions that ask you to determine the number of electrons
transferred per atom or per molecule
Compare the relative strengths of oxidizing and reducing agents based on data
o Recall the strongest oxidizing agent (OA) is found on the top left hand side
of the table (pg. 7 of data book) and the strongest reducing agent (RA) is
found on the bottom right hand side of the table. This means a strong OA
has a more positive electric potential and a strong RA has a more
negative electric potential
o Strong agents are the most reactive and cause the most reactions




Spontaneity Rule
o Spontaneous reaction occurs when the OA is found ABOVE the RA
o Therefore it is important to be able to identify the OA and RA. Can use
oxidation numbers or the short cut that says the OA is usually the more
positive form and the RA is usually the more negative form for a particular
chemical
o Should be able to create a table of oxidizing/reducing agents based on a
series of reactions that are either spontaneous or not
o Should be able to determine the strongest and weakest OA/RA if a chart is
already given showing OAs reacting with different RAs and indicating if
reactions were spontaneous or not
Be able to describe corrosion as a spontaneous redox reactions
o Know methods to prevent corrosion (ie. physical coatings and cathodic
protection)
Stoichometric Titrations
o Need the balanced net ionic reaction
o You can create the half reactions if under acidic and then add the two half
reactions together
o Remember the redox reaction occurs between the strongest RA and OA.
So if no reaction is given, list all entities that are available to react and find
the strongest OA and RA on the table in the data book (pg.7). Simply
copy these two balanced half- reactions and use them to determine the
overall balanced redox reaction by adding the two half-reactions together
o Remember, the electrons lost has to equal the electrons and gained in the
half reactions before the two half-reactions can be added together to get
the redox reaction
o Once you have obtained the balanced redox reaction, you need to 1)
calculate the moles of the substance that you already have the most
information for (ie what you have), 2) create the “want over have” ratio to
calculate the moles of the chemical you are trying to find (ie. what you
want), and 3) use the moles you calculated for the chemical you want to
finish the problem (ie. convert the moles to a mass or concentration).
Voltaic cells vs. electrolytic cells (be able to compare the similarities and
differences between both cells)
o Voltaic cells are a spontaneous redox reaction that produce electrical
energy for chemical energy (ie voltaic cells are batteries)
o Electrolytic cells require a battery to cause a non-spontaneous redox
reaction (ie. converts electrical energy to chemical energy)
o Reduction occurs at the cathode (OA found at cathode) and oxidation
occurs at the anode (RA found at anode) for both cells
o Be able to identify electron flow through external circuit (electrons flow out
of anode and into cathode)



o Be able to identify the ion flow (through salt bridge or porous cup) in a cell
(cations move to cathode and anions move to anode)
o Remember that a voltatic cell is spontaneous. Therefore, strongest OA is
at the cathode and strongest RA is at anode. Or, the more positive
electric potential is found at the cathode and the more negative electric
potential is found at the anode.
o Remember that electrolytic cells are non-spontaneous. Therefore,
weakest OA is at the cathode and weakest RA is at anode. Or, the more
negative electric potential is found at the cathode and the more positive
electric potential is found at the anode.
o Be able to describe the physical observations taking place at the anode or
cathode in any cell based off the half-reactions taking placed at each halfcell (ie. anode decompose, cathode becomes larger, does the
concentration of a certain chemical increase or decrease?, at which
electrode are gas bubbles produced?)
Applications of electrolytic Cells
o Electrolysis is the process of breaking apart molecules into their individual
atoms (ie. metal refining processes and the electrolysis of water or molten
compounds)
o Be careful for the electrolysis of molten compounds (ie NaCl(l)) vs. the
electrolysis of aqueous ionic compounds (NaCl(aq)). In aqueous solutions,
you need to consider that the water can break down too. Therefore, you
need to identify all OA’s and RA’s and find the two agents that have the
smallest difference between their electric potentials
o Electroplating: the object you want to plate/coat has to be on the cathode
and the coating metal must be the anode
Standard Reduction Potential
o The half-reaction for hydrogen was assigned a potential of 0.00V and is
known as the reference half-reaction/half-cell
o Any half-reaction could have been used. If the standard reduction
potential changed from hydrogen, the cell potentials would remain the
same for any cell, but the individual half-potential values would have
shifted all by the same constant
o Should be able to calculate the standard potential of any half-reaction
when the standard reduction potential is changed to another atom other
than hydrogen
Cell Potentials
o Eocell = Eocathode - Eoanode
o ***DO NOT reverse the signs of the standard half-reaction potentials even
if the reaction was reversed/read backwards. Also, NEVER multiply/divide
the half-reaction reduction potentials by a constant/factor to use in the
above equation, even if you had to multiple/divide the half-reactions by
some constant/factor to make the electron transfer equal***

o Negative cell potential means electrolytic cell/non-spontaneous reaction
and a positive cell potential indicates that it is a voltaic cell/spontaneous
reaction
o Remember, the more positive electric potential is the cathode for a voltaic
cell and the more positive electric potential is the anode for an electrolytic
cell
Faraday’s Law
o Need to be able to calculate the mass, amounts, current, and time in a
voltaic cell or electrolytic cell by applying Faraday’s law and stoichiometry
o Need a balanced half-reaction (never the redox reaction) to use Faraday’s
law
o Remember that the moles of electrons in the balanced reaction is very
important and contains information that fits into the following formula
It
ne  
F
o Use Faraday’s law in a regular stoichiometry questions. 1) calculate the
moles of the substance that you already have the most information for (ie
what you have), 2) create the “want over have” ratio to calculate the moles
of the chemical you are trying to find (ie. what you want), and 3) use the
moles you calculated for the chemical you want to finish the problem (ie.
convert the moles to a mass or concentration, or find current or time of the
operating cell).
UNIT C- ORGANIC CHEMISTRY








Need to memorized the carbon molecules that are NOT considered organic (ie.
carbonates, cyanides, carbides, and oxides of carbon)
Look over the common organic molecules that we use in everyday lives
The IUPAC nomenclature/naming is a big part of this unit
o Should be able to name alkanes, alkenes, alkynes, aromatic compounds,
cyclic compounds, alcohols, carboxylic acids, alkyl halides, and esters
Be able to identify the types of compounds from the hydroxyl, carboxyl, ester
linkage and halogen functional groups given a structural formula
Also need to know the properties of each functional group (ie. boiling points,
solubility, smells)
Know the difference between aliphatic (contains no benzene rings) and aromatic
(contains at least one benzene ring)
Know the difference between saturated (all single bonds between carbon-carbon
atoms) and unsaturated (a compound that has at least one double or triple bonds
between carbon-carbon atoms)
Structural isomers have the same number of elements, but different structures
(ie. can be drawn different)









Know the industrial processes (ie. fractional distillation, cracking, reforming, and
solvent extraction)
Addition reaction: start with an alkene/alkynes and a small molecule and end up
with only one product than has less multiple bonds
o Be able to interpret the results of a test to distinguish between a saturated
and an un-saturated aliphatic using the addition of aqueous bromine or
potassium permanganate solutions. Formation of two layers means no
reaction took place, therefore aliphatic was saturated. If a color change
takes place or a precipitated forms, a reaction occurred, which means the
aliphatic was un-saturated
Elimination reaction: start with one organic molecule and end up with a product
that contains a double or triple bond and a small molecule
o The opposite of an addition reaction
o We only looked at elimination of alcohols and alkyl halides
o Know the required conditions for each
Substitution reaction: the functional groups switch places
o We only looked at the substitution of alcohols and alkyl halides
o Alkanes and aromatics can also undergo substitution, but require energy
or a catalyst
Esterification reaction: reaction between an alcohol and a carboxylic acid to
produce an ester and water
o This reaction is also known as a condensation reaction
Combustion reactions: hydrocarbon and oxygen react to produced carbon
dioxide and water and thermal energy/heat
o Complete versus incomplete
Be able to provide examples of polymers in living and non-living systems
Condensation polymerization and addition polymerization reactions to form
polymers
Polymers are named after the monomer that was used in the reaction to create
the polymer or names after the functional group formed in the polymer
UNIT D – CHEMICAL EQUILIBRIUM: FOCUSSING ON ACID-BASE REACTIONS

Equilibrium is defined when a system has its forward and reverse reaction
happening at the same rate
o There are no physical observations or measurements that can be
observed on a macroscopic level when a system is at equilibrium
o At a microscopic level, both the forward and reverse reaction occur at
same rate
o Equilibriums can only be established in a closed system
o Use the double headed arrow to show forward and reverse
reactions/equilibrium

Equilibrium expression/law can be written as follows for any system at
equilibrium
aA + bB ↔ cC + dD
Kc



c
d

C  D 

Aa B b
or more simply K c 
products
reac tan ts
o Remember that a chemical that is liquid or solid does not appear in the
equilibrium expression
o Remember that the coefficients in the balanced equation are the
exponents in the equilibrium expression
Kc is the equilibrium constant
o Small Kc means the reactants are favoured
o Large Kc means the products are favoured
Le Chatelier’s Principle
o The four stress we looked at were temperature change, volume/pressure
change, remove/add a chemical, or add a catalyst
o Temperature is the only stress that will change the equilibrium constant
(Kc). If the shift is toward the products, the Kc value would increase.
o Catalyst will have no effect on the equilibrium (will not shift it to the left or
the right), it will only speed up the time it takes the system to reach
equilibrium
o If you apply a stress, the system will always do the opposite to try and
counter the stress that you applied
o For a pressure change, you only need to consider the reactants and
products that are gaseous
o For temperature change, you need to know the enthalpy change (∆H) and
write it into the equation (ie. make a thermochemical equation).
Remember to treat a change in temperature like a concentration and you
can increase/decrease the “concentration” of the energy by changing the
temperature
o If a precipitate forms, that means the concentration of the aqueous ions
involved in the precipitate will decrease
o If an acid is added, that means the concentration of H3O+/H+ is increasing
and the concentration of OH- is decreasing
o If an base is added, that means the concentration of H3O+/H+ is
decreasing and the concentration of OH- is increasing
Need to use ICE tables to find equilibrium concentrations when initial
concentrations are given
o Remember ICE table values need to be a concentration
o Use ICE tables with the equilibrium expression
o Know when and how to use the approximation rule (when Kc is very small)









o For ICE tables to work properly you will always need to know the
equilibrium constant or an equilibrium concentration
Recall the Bronsted- Lowry definitions for acid and bases as proton acceptors
and donors
Acids will have a low pH (high pOH), which means there are more H3O+/H+ ions
than OH- ions
Bases will have a high pH (low pOH), which means there are more OH- ions
than H3O+/H+ ions
Need to know how to predict the direction of an acid-base reaction
o Identify the acid or base on both the reactant side and product side
o The reaction favours the side with the weaker acid or base
Be able to identify conjugate acid-base pairs (two substances that only differ by
the absence or presence of one hydrogen ion/proton
Know the difference between amphiprotic and polyprotic substances
Strong acids and bases ionize/dissociation complete (ie. reactions go to
completion) and cannot form an equilibrium
o Only six strong acids (located on the top of table on pg. 8-9 of data book)
o Strong bases contain a OH- ion
Only weak acids and bases can establish an equilibrium
The equilibirium expression for a acid is
HA(aq) + H2O(l)
Ka 
A

↔
H O  = x
HA 
HA 
( aq )

3
A-(aq) + H3O+(aq)
2
( aq )
( aq )
where x= [H3O+(aq)] = [A-(aq)]
( aq ) initial
using the approximation rule

The equilibrium expression for a base is
B(aq) + H2O(l)
Kb 
HB

OH
B 
( aq )
( aq )


( aq )
=
x2
B 
↔
HB+(aq) + OH-(aq)
where x= [OH-(aq)] = [HB+(aq)]
( aq ) initial
using the approximation rule
Remember that there is a lot of information in the data book regarding acids and
their conjugate base
o Ka values for many common acids and indicators are in data book (pg. 89, 10)
o Names of acids are also found in that table
o If the Ka is known for an acid, then the Kb for the conjugate base can be
found out as well
where Kw = 1.0x10-14 and is found on pg. 3 of
K w  K a  Kb

data book
o Remember top six strong acids are found at the top left of the table in pg.
8-9 of data book and decrease in strength as you go down
o Strong bases are ionic compounds that include OH- ions (ie. NaOH(aq) or
KOH(aq)). The weakest bases are found at the top right of the table (pg. 89) and increase in strength as you go down
Acid/base equations from chemistry 20 that are important for this unit are as
follows:
pH   log H 3O 
H 3 O   10  pH


pOH   log OH 

 
OH   10

 pOH
14 = pOH + pH



Buffers are systems that resist a change in pH
o A buffer system needs to be made up of a weak acid and its conjugate
base or vice versa. The conjugate acid/base pair should be in
approximately equal concentrations
o Know the three different ways to create a buffer. 1.) Add equal amounts
of a conjugate acid/base pair together. 2.) Add a small amount of a strong
acid to an excess amount of weak base. 3.) Add a small amount of a
strong base to an excess amount of weak acid.
We looked at a few different titration curves
o Know the difference between a titration of a strong base with a strong
acid, a weak acid titrated with a strong base, and a weak base titrated with
a strong acid
o The differences include where the equivalence pH is located, where the
curves start (indicates what sample is), and where the curves end
(indicates what titrant is)
o A polyprotic substances has more than one equivalence point
o Be able to identify the area on a titration curve that represents a buffer (at
the beginning of the titration curve when a small amount of strong titrant is
added to a weak sample)
Indicators are listed on pg. 10 of data booklet
o You want the pH range of the indicator to include the pH of the
equivalence point/endpoint when doing a titration
o Indicators are weak acids that can donate hydrogen ions. The conjugate
acid/base pair for an indicator are two different colors. An indicator can be
used in acid/base calculations/equilibriums