Download AP CHEMISTRY NAME____________________ WRITING

AP CHEMISTRY
NAME____________________
WRITING EQUATIONS
Word and formula equations
1. Two types of equations are written by chemists:



word equations: describe the substances that react in a chemical reaction (termed reactants ), and
the products that are formed, along with their states
formula equations are a shorthand method used to describe the same reactions
These are of two types:
o Skeleton equations: (unbalanced) which lists the correct formula of each reacting
substance and product substances, and their states.
o Balanced equations: which list the correct formulas, states and balances the equation for
the number of atoms present. That is it takes into account the Law of Conservation of
Mass, and makes sure there is the same number and type of atom in the reactant and
product.
2. We balance equations by changing the coefficients or numbers in front of the substance.
WE NEVER CHANGE THE FORMULAS OF SUBSTANCES IN ORDER TO BALANCE.
TYPES OF REACTIONS:
1. SYNTHESIS REACTION: (also called composition or combination)
In a synthesis reaction two or more simple substances combine to form a more complex substance. Two or
more reactants yielding one product is another way to identify a synthesis reaction
1. two or more substances combine to form a more complex substance
2. General form: A + X  AX
a. Element + Element  Compound
b. Two compounds  larger compound
metallic oxide + water  base
nonmetallic oxide + water  acid
2. DECOMPOSITION REACTION(also called analysis)
1. substance breaks down to form two or more simpler substances, usually caused by heating or
electric current
2. General form: AX  A + X
Compound  Elements and/or compounds
In a decomposition reaction a more complex substance breaks down into its more simple parts. One reactant
yields 2 or more products. Basically, synthesis and decomposition reactions are opposites.
SIX TYPES OF DECOMPOSTION REACTIONS:

1. Heat a metallic carbonate 

 metallic oxide + CO2
CaCO3 → CaO + CO2
(MCO3 → MO + CO2)

2. Heat a metallic chlorate 

 metallic chloride + O2
2KClO3 → 2KCl + 3O2
(MClO3
→ MCl + O2)
3. Decomposition by electricity
2H2O2 electricity

 2H2O + O2

4. Heating certain oxyacids 

 nonmetallic oxide + H2O
H2SO3 --> H2O + SO2
(nonmetal retains charge in product)

5. Heating a base 

 metallic oxide + H2O

Ca(OH)2 

 CaO + H2O

6. Heating a metallic oxide 

 metal + O2

2 HgO 

 2 Hg + O2
7. Decomposition of ammonium carbonate:

(NH4)2CO3 

 NH3 + CO2 + H2O
Activity
Series
cations
Li
Rb
K
Cs
Ba
Sr
Ca
Na
Mg
Al
Mn
Zn
Cr
Fe
Ni
Sn
Pb
H
Cu
Hg
Ag
Pt
Au
anions
F
Cl
Br
I
3. SINGLE REPLACEMENT REACTIONS
1. one substance is replaced in its compound by another substance
2. General form: A + BX  AX + B
or Y + BX  BY + X
Element + Compound  Element + Compound
In a single replacement (or displacement) reaction one element replaces another element in
a compound. They follow one of the following two patterns:
 Halogen replacement (single anion replaces anion in compound) or
 Replace another metal or hydrogen (single cation replaces cation in compound)
 elements that form cations can only replace cations and elements that form anions
can only replace anions
1. Replacement of a metal by a more active metal:
2 Al(s) + Fe2O3(aq) --> Al2O3(aq) + 2 Fe(s)
2. Replacement of hydrogen in water by an active metal:
2Na(s) + 2HOH(l) --> 2NaOH(aq) + H2(g)
3. Replacement of hydrogen in an acid by a metal:
Zn(s) + 2 HCl(aq)  ZnCl2(aq) + H2(g)
4. Replacement of a halogen by another halogen:
Cl2(g) + 2 KBr(aq) --> 2 KCl(aq) + Br2(l)
5. Notice that the metals lower than hydrogen on the Activity Series will not replace hydrogen in an acid; these
are the metals used for jewelry, coins, eating and cooking utensils. If you see one of these metals with an acid,
the reaction is most likely redox.
4. DOUBLE REPLACEMENT REACTIONS (ION-EXCHANGE):
 Involve no transfer of electrons
 General form: AX + BY
 AY + BX
+
A (aq) + B (aq)  AB(s)
 Precipitation Reactions – know solubility table
[to be a reaction, the product must be a solid, liquid, or gas; aqueous ions remain in solution as
spectator ions]
 Acid base reactions
 neutralization: acid + base  salt + water
 Be alert for Bronsted-Lowry and Lewis reactions
 Hydrides, boron compounds, amides and liquid ammonia indicate unusual acid-base reactions
(unshared pairs of electrons)
 A very broad topic - a lot of reactions can be classified:
 bicarbonate ion,
 complex ion reactions;
 water is both an acid and a base
 Know the conjugate acid-base table.

Hydrolysis – adding water to a salt (that is not a neutral salt)
a. really an acid-base reaction in reverse
b. usually H+ or OH is one of the products
c. Remember the coordination number for metals

Amphiprotic (or amphoteric) reactions  may act as acid OR base
a. Look for “excess NaOH or KOH
b. Usually zinc and aluminum ions are involved, as well as Cr3+, Sn2+, and Pb2+
c. Water and ammonia are amphiprotic.
In a double replacement reaction parts of two compounds switch places to form two new compounds. Two
reactants yield two products.
Remember that one or more of the following must be formed for a reaction to occur; if the same four ions are
present in the products, no reaction has taken place.

When the two solutions are added together, the silver ions and chloride ions find each other and
become a solid precipitate. (They ‘rain’ or drop out of the solution, this time as a solid.) Since silver
chloride is insoluble in water, the ions take each other out of the solution.
Ag+ (aq)+ NO3- (aq)+ K+(aq) +Cl- (aq)
AgCl(s) + K+(aq)+ NO3-(aq)

Formation of water (such as in an acid-base reaction): The hydrogen and hydroxide ions take each
other out of the solution by making a covalent compound (water).
HCl + NaOH
HOH + NaCl or
H+((aq) + Cl- (aq)+ Na+ (aq)+ OH-(aq)

One more way for the ions to be taken out of the water is for some of the ions to escape as a gas.
CaCO3 + 2 HCl
CaCl2 + H2O + CO2
Ca2+ + CO32- + 2 H+ + 2 Cl




HOH(l) + Na+(aq) + Cl- (aq)
Ca2+ + 2 Cl- + H2O(l) + CO2 (g)
One way to consider double replacement reactions is as follows: Two solutions of ionic compounds are
really just sets of dissolved ions, each solution with a positive and a negative ion material. The two are
added together, forming a mixture of four ions. If two of the ions can form
(1) an insoluble material,
(2) a covalent material such as water, or
(2) a gas that can escape, it qualifies as a reaction
Not all of the ions are really involved in the reaction. Those ions that remain in solution after the
reaction has completed are called spectator ions, that is, they are not involved in the reaction. There is
some question as to whether they can see the action of the other ions, but that is what they are called.
DOUBLE REPLACEMENT REACTIONS FOLLOWED BY DECOMPOSITION:

Any carbonate that reacts with an acid produces CO2 + H2O
Solid potassium carbonate reacts with 0.5M sulfuric acid. (0.5M tells you the solution is dilute)
K2CO3 + H2SO4  K2SO4 + H2CO3
↑
↑
↑
↑
Solid
strong acid
soluble
decomposes to CO2 + H2O
So net ionic equation is:
K2CO3 + H+  K+ + CO2 + H2O
Other cases:
 H2SO3 → H2O + SO2
 NH4OH → NH3 + H2O
5. COMBUSTION REACTIONS:
Combustion reactions always involve molecular oxygen, O2. Anytime anything burns (in the usual sense), it is a
combustion reaction. Combustion reactions are almost always exothermic.
Methanol burns in air.
C3H8O + O2 --> CO2 + H2O
If you have not learned how to balance with 1/2 coefficients, realize that if you double all coefficients when you
have an odd number of O’s that that seems to help balance the odd number of O’s.
2 C3H8O + 9 O2 --> 6 CO2 + 8 H2O
Of course, not all combustion reactions release CO2 and water, e.g., the combustion of magnesium metal:
6. BALANCING REDOX REACTIONS:
A. Look for reactions that
 are single replacement
 involve a transition element (several oxidation states)
 have polyatomic ions with an element in a high oxidation state
conc. HNO3 (N = +5), H2SO4 (S = +6), KMnO4 (Mn = +7), K2CrO7 (Cr = +6)
 include the words are “oxidized by” or “reduced by”
B. Look for strong oxidizing agents, such as conc. HNO3, HCl, K2CrO4, KMnO4







To balance redox reactions, assign oxidation numbers to the reactants and products to determine how
many moles of each species are needed to conserve mass and charge.
First, separate the equation into two half-reactions, the oxidation portion and the reduction portion.
This is called the half-reaction method of balancing redox reactions or the ion-electron method.
Each half-reaction is balanced separately and then the equations are added together to give a balanced
overall reaction. We want the net charge and number of ions to be equal on both sides of the final
balanced equation.
in redox reactions, the number of electrons lost by the reducing agent must be equal to the number of
electrons gained by the oxidizing agent
Look for word clues like “acidified”, “dilute”, “concentrated”, “warmed”
Look for a polyatomic ion that has an element with a high oxidation state, like permanganate, chromate,
dichromate, nitric or sulfuric acid
A special kind of redox is disproportionation – the same element is both oxidized and reduced; the
halogens and peroxides are capable of this. (Also, lead in the storage battery.)
Pb(s) + PbO2 (s) + SO42-(aq)
PbSO4(s)
Chlorine gas is bubbled into a solution of cold, dilute sodium hydroxide.
3Cl2 + 6OH− → 5Cl− + ClO3− + 3H2O

Write the “skeleton equation”, that is, leave out spectators.
Iron sulfate reacts with potassium permanganate in acidic solution
Fe2+(aq) + MnO4-(aq)
Fe3+(aq) + Mn2+(aq) (acidic solution)
STEPS FOR BALANCING REDOX REACTONS:






Balance element being oxidized or reduced
If polyatomic ion is present, satisfy O atoms by adding H2O to opposite side.
If water is added, balance H atoms by adding H+ ions to opposite side.
Balance charges by adding electrons:
o for reduction electrons are gained by adding to the left side of the arrow
o for oxidation electrons are lost by adding to the right side of the
Most but not all are redox reactions. For example: 2H2O → H3O+ + OH- is a disproportionation but is
not a redox reaction.
The reverse of disproportionation is called comproportionation.
7. REACTION OF WATER WITH A HYDRIDE:




Most common hydrides are Group 1 or 2A hydrides.
Hydride is the name given to the negative ion of hydrogen, H−
Hydrides bonds range from very covalent to very ionic as well as multi-centered bonds and metallic
bonding.
Metallic hydride + water  base + hydrogen gas
NaH(s) + H2O (l) → H2 (g) + NaOH(aq)
8. LEWIS ACID-BASE REACTIONS:
A Lewis acid is an electronpair acceptor, and a Lewis base is an electronpair donor. The mixing
of two species capable of forming a coordinate covalent bond results in a synthesis reaction.


Lewis definition defines a base (referred to as a Lewis base) to be a compound that can donate an
electron pair, and an acid (a Lewis acid) to be compound that can receive this electron pair
consider this classical aqueous acid-base reaction:
HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
Boron trifluoride gas is mixed with trimethylamine gas.
Phosphine (phosphorus trihydride) is bubbled into liquid boron trichloride.
The Lewis definition does not regard this reaction as the formation of salt and water or the transfer of H+ from
HCl to OH−. Instead, it regards the acid to be the H+ ion itself, and the base to be the OH− ion, which has an
unshared electron pair.
Therefore, the acid-base reaction here, according to the Lewis definition, is the donation of the electron pair
from OH− to the H+ ion. This forms a covalent bond between H+ and OH−, thus producing water (H2O).
Complex ion formation/destruction results in the attachment/detachment of a ligand (Lewis base) to a transition
metal (Lewis acid) or aluminum. (Expect ligands like ammonia, cyanide, hydroxide, thiocyanate to complex.) Very often,
soluble ion will be taken out of solution as a precipitate, or a solid precipitate will dissolve. Noticeable color changes occur

the Lewis definition can be applied to reactions that do not fall under other definitions of acid-base
reactions, such as complex ions:
Ammonia gas is bubbled into a solution of silver nitrate.
Ag+ + 2 :NH3 → [H3N:Ag:NH3]+
where the : represents the electron pair
1. Aqueous ammonia is rubbed onto a coating of silver chloride.
2. Excess hydrochloric acid is added to diamminesilver(I) nitrate.

Another theory describes an acid as an oxide ion acceptor and a base as an oxide ion donor.
MgO (base) + CO2 (acid) → MgCO3
CaO (base) + SiO2 (acid) → CaSiO3

Reactions with boron (atom lacks an octet)
9. Buffers:
A buffer is a solution that contains a weak conjugate acidbase pair and can resist drastic changes in pH when
a strong acid or base is added. It can do this because it contains both an acidic species to neutralize OH  ions
and a basic one to neutralize H+ ions, and yet do not consume each other. Consider a buffer made of a weak
acid (HX) and one of its salts (MX) where pH is determined by the ratio of the concentrations of the conjugate
acidbase pairs, [HX]/[X].

If a strong base is added to the buffer, it reacts with the acid component:
OH(aq) + HX(aq)  H2O(l) + X(aq)

where [HX] decreases and [X] increases
If a strong acid is added to the buffer, it reacts with the base component:
H+(aq) + X(aq)  HX(aq) + H2O(l)
where [X] decreases and [HX] increases
As long as the change in the ratio [HX]/[X] is small, the change in pH will be small.
1. A solution of sodium hydroxide is added to a solution of sodium dihydrogen phosphate until the
same number of moles of each compound has been added.
2. A solution of ammonium sulfate is added to a potassium hydroxide solution.
10. Formation of Esters and Organic Salts:
Carboxylic (organic) acids can undergo condensation reactions with alcohols to form esters. The H
atom of the acid is replaced by a hydrocarbon group. Esters are named by using the name of the
alcohol first and then the group that is derived from the acid. Esters usually have aromas that are
pleasant, such as fruits and flowers. Notice that the name is the reverse of the formula.
Acetic acid +
ethanol →
ethyl acetate +
water
CH3COOH
+
O
||
CH3COH +
C2H5OH
→
HOCH2CH3
CH3COOCH2CH3
→
O
||
CH3COCH2CH3 + H2O
Ethanol and formic acid are mixed.
11. Reactions of Alkenes (Markovnikov's rule): Unsaturated hydrocarbons such as alkenes and
alkynes are much more reactive than the parent alkanes.
A. Halogenation: They react rapidly with bromine, for example, to add a Br 2 molecule across the
C=C double bond.
B. The addition of a protic acid such as H-X to an alkene
 the acid hydrogen (H) becomes attached to the carbon atom with the greatest number of
hydrogens on asymmetrical molecules
 the halide (X) group becomes attached to the carbon with the fewest hydrogens
 In other words the negative part of the addendum attaches itself to the carbon atom with the
fewest hydrogen atoms.
 Empirical rule which predicts that the addition on a double bond of the hydrogen of a HX acid
occurs on the most hydrogenated carbon atom.
CH3CH = C(CH3)CH(CH3)2 + HBr → CH3CH2CH2BrCH(CH3)2

On symmetrical molecules, the rule does not apply
C. An alkene reacts with water in an addition reaction to form alcohol.
 The hydroxyl group (OH) bonds to the carbon that has the greater number of carbon-carbon
bonds
 the hydrogen bonds to the carbon on the other end of the double bond, that has more carbonhydrogen bonds
The rule may be summarized as "the rich get richer and the poor get poorer": a carbon rich in substituents will
gain more substituents and the carbon with more hydrogens attached will get the hydrogen in many organic
addition reactions.
D. The oxidation of primary alcohols to aldehydes and carboxylic acids
When a primary alcohol is converted to a carboxylic acid, the terminal carbon atom
increases its oxidation state by four.
Primary alcohols can be oxidized to either aldehydes or carboxylic acids depending on
the reaction conditions. In the case of the formation of carboxylic acids, the alcohol is
first oxidized to an aldehyde which is then oxidized further to the acid. [O] is from the
oxidizing agent.
E. The oxidation of secondary alcohols to ketones
The hydrogen from the hydroxyl group is lost along with the hydrogen
bonded to the second carbon. The remaining oxygen then doublebonds with the carbon. This leaves a ketone, as R1-CO-R2. Ketones
cannot normally be oxidized any further because this would involve breaking a C-C bond, which requires too
much energy.