Chemical Reactions & Equations
The major event in chemistry – chemical reaction,
i.e. chemical transformation of a substance.
The ability (or inability) to undergo chemical reaction
is chemical property.
Examples: Ca metal is fast oxidized by air O2 producing calcium oxide, CaO
- this is a chem. property of calcium. Therefore, Ca is not found in the
nature as a free metal, but only in its compounds.
Gold metal, Au, is not oxidized by O2, & most of other known oxidants, but
reacts with extremely aggressive “Aqua Regia” - a mixture of hydrochloric HCl &
nitric HNO3 acids. These are chem. properties of gold.
In the language of chemistry, chemical reactions are presented as
chemical equations:
“Sentences of Chemistry”
They consist of two parts:
Reactants
2Ca(s) + O2(g)  2CaO(s)
 Products
2Au(s)+11HCl(l)+3HNO3(l)2HAuCl4(aq)+2NOCl(g)+6H2O(l)
Chemical equation shows:
- The nature of reactants & products
- The molar ratios between all the participant substances
These are expressed through STOICHIOMETRIC COEFFICIENTS in front of the formula
- Additional information on the physical state of all substances involved &
conditions of the reaction (as elevated temperature, pressure, etc.)
Types of Chemical Reactions
Synthesis (combination): 2 or more substances combine into one substance,
A+B  P
2Cu(s)+O2(g)2CuO(s)
SO3(s) + H2O(l)  H2SO4(l)
Decomposition: 1 substance transforms into 2 or more substances
A  B+C+…
(NH4)2Cr2O7(s)  N2 + Cr2O3(s) + 4H2O(g)
Single displacement: AX + Y  AY + X
2KI(s) + Cl2(g)(s) 2KCl(s) + I2(g)
CH4(g) + Cl2(g)  CH3Cl(g) + HCl(g)
Zn(s) + HCl(aq)  ZnCl2(aq) + H2(g)
Double displacement: AX + BY  AY + BX (never AB as a product!)
NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq)
CH3I + KOH  CH3OH + KI
Typical problems: identify the type, predict the product of a rxn, &
terminate a chemical equation:
Mg + O2  MgO synthesis
2HgO  2Hg + O2 decompopsition
Fe + 2HCl(aq)  FeCl2 + 2H2 single displacement
BaCl2(aq) + Na2SO4(aq)  BaSO4(s) + 2NaCl(aq)
AgNO3(aq) + NaBr(aq)  AgBr(s) + NaNO3(aq)
Synthesis, Decomposition & Single displacement
with free elements
may be an oxidation-reduction rxn, i.e. a rxn
in which oxidation numbers of elements change:
Synthesis: Sn + 2Cl2  SnCl4
LEFT SIDE:
RIGHT SIDE:
Sn(0) Cl(0)
Sn(+4) Cl(-1)
Oxidation number of a free element is always ZERO.
Sn(0) - 4e-  Sn(+4) oxidation of tin, Sn
Cl2 (0) +2e-  2Cl(-1) reduction of chlorine, Cl2
Decomposition: at heating,
HgO(s)  Hg(l) + O2(g)
Hg(+2) O(-2)
Hg(0)
O(0)
Mercury was reduced: Hg(+2) +2e  Hg(0)
Oxygen was oxidized: O(-2) -2e- O(0)
There are many non-redox Synthesis & Decomposition rxns with
compounds rather than elements:
SO3 + H2O  H2SO4 is a synthesis rxn, but not a redox rxn (no element
changes its oxidation number)
CaCO3  CaO + CO2
Formation of molecules by combination of atoms is also a synthesis rxn, but
not a redox one: 2I  I2 or 2O3  3O2
Single Displacement:
2KBr(s) + Cl2(g)  2KCl(s) + Br2(l)
Oxidation numbers of elements:
LEFT SIDE:
K(+1) Br(-1) Cl(0)
RIGHT SIDE:
K(+1) Cl(-1) Br(0)
In the course of the rxn:
2Br(-1) -2e-  Br2 (0)
Cl2 (0) +2e-  2Cl(-1)
oxidation of bromide ion, Br-
reduction of chlorine, Cl2
More active non-metal displaces less active non-metal from its compounds
Most active non-metals are in the upper right part of the Table; Cl2 more active than Br2 or I2,
therefore it displaces them as in the above rxn.
2HCl(aq) + Zn(s)  ZnCl2(aq) + H2(g)
2H(+1) + 2e-  H2(0)
Zn(0) - 2e-  Zn(+2)
hydrated H+ ion was reduced
zinc metal was oxidized
Cl(-1) does not change
CuCl2(aq) + Zn(s)  ZnCl2(aq) + Cu(s)
More active metals displace less active metals from their compounds
Cu(+2)  Cu(0) Zn(0)  Zn(+2)
Li>K>Ba>Ca>Na>Mg>Al>Mn>Zn>Fe>Cd>Co>Ni>Sn>Pb>H2>Cu>Ag>Pt>Au
ACTIVITY SERIES OF METALS
contains metals arranged according to their
decreasing
activity.
Those to the left displace those to the right from their
compounds:
Metals to the left of hydrogen displace H2 from acids, i.e. aqueous solutions
of HA compounds releasing H+ ions into solution, while metals to the right
from H2 cannot do that. M + HA  MA + H2↑ (i.e. hydrogen gas)
Most
Activ
e
displace H2 from liquid water: Ca + H2O  Ca(OH) 2 + H2
displace H2 from acids: Zn + HCl  ZnCl2 + H2
& from hot steam: Mg + H2O  MgO + H2
displace H2 from acids: Sn + HCl  SnCl2 + H2
do not displace hydrogen
Least
activ
e
Typical problem:
Determine that some single displacement rxn can, or cannot occur, & if it
can, predict its products (terminate the equation):
Cu + AgNO3  Cu(NO3)2 + Ag↓
Cu + HCl(aq)  (no rxn)
Fe + MgSO4 
Al + HCl(aq) AlCl3 + H2↑
Activity series predicts rxns in aqueous solutions.
Most active metals (IA Group, also Ca, Ba) displace hydrogen not only
from acids, but from liquid water producing H2 gas & metal hydroxide
(alkali):
2Na +2H2O  2NaOH + H2
Ca+2H2O  Ca(OH) 2+ H2
At high temperature, many other metals displace hydrogen from water
vapor (steam):
Mg + H2O(g)  MgO(s) + H2
Fe + H2O(g)  FeO(s) + H2
Other reducing agents:
At heating, H2 & C displace most of metals from their compounds,
especially from oxides – most common form of metal ores.
symbol for heating
C(s) + Fe2O3(s)  Fe(s) + CO(g)

H2(g) + Fe2O3(s)  Fe + H2O(g)
Those are also oxidation-reduction rxns.
Single displacement is not necessarily a redox rxn:
SiO2 + CaCO3  CaSiO3 + CO2 no change in oxidation numbers.
A
+ BX  AX + B
COMBUSTION is a particular case of oxidation-reduction with O2 as
oxidant & resulting in strong heat release:
Gasoline:
C8H18(g) + O2(g)  CO2(g) + H2O(g)
Oxygen is reduced, carbon is oxidized.
BALANCING CHEMICAL EQUATIONS
Since atoms do not disappear in chemical reactions, a proper
chemical equation must be balanced for each element involved.
3C + Fe2O3  3CO + 2Fe
C8H18 + O2  CO2 + H2O
Never start with a free element. Ignore it first. There are 8C to the left – need 8CO2 to the right:
C8H18 + O2  8CO2 + H2O Now count H: there are 18H to the left, put 9H2O to the
right:
C8H18 + O2  8CO2 + 9H2O Now count oxygen to the right: there are 16+9=25
oxygen atoms.
To the left, there is 2 oxygens – consider at first, that this is not dioxygen, but O:
C8H18 + 25 O  8CO2 + 9H2O
Since it should be O2, take it all twice:
2C8H18 + 25O216CO2+18H2O
2Al + 3FeO  Al2O3 + 3Fe
FeS2 + O2  Fe2O3 + SO2
Check if balancing for each element.
balance Fe first:
2FeS2 + O2  Fe2O3 + SO2 next balance S
2FeS2 + O2  Fe2O3+ 4SO2 count oxygen to the right: 11 O, take 11 O to the left, ignore O2
2FeS2 + 11O2  Fe2O3 + 4SO2 take it all twice:
4FeS2 + 11O2  2Fe2O3 + 8SO2
check the balance for each element.
Reactions of Ions in Aqueous
Solutions
Ionic compounds dissolve in water by dissociating into
hydrated ions:
Dissolution of NaCl:
H2O
NaCl(s)  Na+(aq) + Cl-(aq)
Dissolution of AgNO3:
H2O
AgNO3(s)  Ag+(aq) + NO3-(aq)
Substances that form ions at dissolution are
ELECTROLYTES.
Their solutions conduct electrical current.
Ions in a solution travel & react independently from
each other.
If a combination of ions in a solution forms an insoluble compound, that
compound precipitates, meaning that the reaction occurs & may go to
completion.
+
When solutions of NaCl & AgNO3 are mixed, Ag & Cl form
an insoluble salt AgCl
Na+(aq) + Cl-(aq)+Ag+(aq) + NO3-(aq)  AgCl + Na+(aq) + NO3-(aq)
Only Ag+ & Cl- ions participate in the rxn forming a solid,
while Na+ & NO3- remain in the solution. They are
SPECTATOR IONS.
NET IONIC EQUATION ignores spectator ions:
Ag+(aq) + Cl-(aq)  AgCl(s)
Spectator ions may be replaced by other ions with no
change in the net ionic equation.
KCl(aq)+AgC2H3O2(aq)AgCl(s)+KC2H3O2(aq)
Classes of Inorganic Compounds
Metal cations (&NH4+)
Non-metal & polyatomic anions
Na+, Mg2+, Al3+, Pb4+ etc.
Cl-, O2-, S2-, N3-, C4-, OH-, SO42-, CO32-, MnO4-, etc.
Their combinations:
Metal cation & oxide:
Metal oxide
Cu2O, MgO, Fe2O3, PbO2 etc.
Metal cation & OH- hydroxide: Metal hydroxide
NaOH, Mg(OH)2, Al(OH)3 etc.
Metal cation & any some other anion: SALTS
NaCl, CoS, Ca(NO3)2, KMnO4 etc.
Acids: compounds that release H+ in a solution:
HCl(aq)  H+(aq) + Cl-(aq),
Major acids: HCl
hydrochloric acid, (salts: chlorides)
HNO3 nitric acid (salts nitrates)
H2SO4 sulfuric acid (sulfates)
HClO4 perchloric acid (perchlorates)
HNO2 nitrous acid (nitrites)
H2SO3 sulfurous acid (sulfites)
HClO hypochlorous acid
(hypochlorites)
H3PO4 phosphoric acid (phosphates)
H3PO3 phosphorous acid
(phosphites)
Acetic acid CH3COOH  CH3COO- + H+ (acetates)
Bases: compounds accepting H+ in a solution, in particular,
those containing OH-, such as metal hydroxides:
Ca(OH)2  Ca2+ +2OHThese OH- combine with H+ producing water, H-O-H
This is NEUTRALIZATION rxn:
+
-
H + OH  H2O
e.g. NaOH + HCl  NaCl + H2O
The outcome of a double displacement rxn between ionic compounds
may be predicted out of solubility rules.
General solubility rule known to alchemists: like dissolves like.
Non-polar organic substances do not dissolve in water, but dissolve in non-polar
organic liquids, e.g. gasoline (octane, C8H18) does not dissolve in water, but nonpolar organic substances dissolve in octane.
Ionic compounds may dissolve in water, to some extent – in other polar solvents, as
ethanol, C2H5OH, but not in octane. Aqueous solutions are most important for ionic
compounds.
To determine if a precipitate will form in an ionic double
displacement rxn, apply
SOLUBILITY RULES
1. All salts of Na+, K+ (IA Group) & ammonium, NH4+, &
All nitrates (NO3-) & acetates (C2H3O2-) are soluble (no
exception). [e.g. K3PO4, (NH4)2CO3, Pb(NO3)2, Mg(C2H3O2)2]
2. Most chlorides (Cl-) & sulfates (SO42-) are soluble, except:
AgCl, BaSO4, PbSO4 which are insoluble. [FeCl2, MgSO4 ]
3. Most carbonates (CO32-), phosphates (PO43-), sulfides (S2-) &
hydroxides (OH-) are insoluble (except listed above).
e.g. [MgCO3, FePO4, CuS, Pb(OH)2]
4. All salts containing H in the anion (acidic salts) are soluble:
insoluble
soluble
CaCO3
Ca(HCO3)2
Ca3(PO4)2
CaHPO4 Ca(H2PO4)2
Typical problem:
Predict if a rxn will occur when the two given solutions are mixed.
Compile the equations for dissolution of the reagents, then full
ionic & net ionic equations where appropriate:
Solutions of Ba(NO3)2 & Na2SO4
Ba(NO3)2(s)  Ba2+(aq) + 2NO3-(aq)
Na2SO4(s)  2Na+(aq) + SO42-(aq)
Ba2+(aq)+2NO3-(aq)+2Na+(aq)+ SO42-(aq)  BaSO4 + 2Na+(aq) + 2NO3-(aq)
Net ionic equation:
Ba2+(aq)+SO42-(aq)BaSO4
Solutions of KOH & FeCl3
KOH(s)  K+(aq) + OH-(aq)
FeCl3(s)  Fe3+(aq) + 3Cl-(aq)
3K+(aq) + 3OH-(aq) + Fe3+(aq) + 3Cl-(aq)  Fe(OH)3 + 3K+(aq)
+ 3Cl-(aq)
Fe3+(aq) + 3OH-(aq)  Fe(OH)3 
Solutions of (NH4)2CO3 & CaCl2
Complete independently.
Two other reasons, besides precipitation, why a reaction would
occur & go to completion:
Formation of a gas (bubbles observed, one of the products
escapes as a gas):
Example 1.
Solutions of Na2CO3 & HCl mixed:
Na2CO3(s)  2Na+(aq) + CO32-(aq)
HCl  H+(aq) + Cl-(aq)
2Na+(aq) + CO32-(aq)+ 2H+(aq) + 2Cl-(aq)  2Na+(aq) + 2Cl-(aq) + H2O + CO2
Na+ & Cl- are spectator ions
Net ionic equation:
CO32-(aq)+ 2H+(aq)  H2O + CO2
Example 2.
KOH + NH4Cl  NH3 + KCl + H2O
Complete independently.
___________________________________________________________________________________________________________
Formation of a molecular, non-ionized compound, such
as water:
KOH + HCl  H2O + KCl
K+ + OH- + H+ + Cl-  H2O + K+ + ClNet ionic equation
+
-
H + OH  H2O
Neutralization rxn
(highly exothermic!)
Water formation is a powerful driving force for a rxn to occur.
Even water-insoluble metal hydroxides, such as Cu(OH)2,
or metal oxides, as Fe2O3 react with acids to form water:
Cu(OH)2(s) + HCl(aq)  CuCl2(aq) + H2O
Net ionic: Cu(OH)2(s) + 2H+(aq) Cu2+(aq) + 2H2O
Fe2O3(s) + HCl(aq)  FeCl3(aq) + H2O
Net ionic: Fe2O3(s) + 6H+(aq)  2FeCl3(aq) + 3H2O
This is also a neutralization rxn (since acid disappears).
Non-metal oxides react with bases:
SO2(g) + Ca(OH)2(aq)  CaSO3(s) + 2H2O
Metal & non-metal oxides react between themselves:
CaO(s) + CO2(g)  CaCO3(s)
Oxidation-Reduction in an aqueous solution can also be
considered in terms of participating ions
Cu(s)+2AgNO3(aq)  Cu(NO3)2(aq) +2Ag(s)
Cu(s)+2Ag+(aq) +2NO3-(aq)Cu2+(aq) + 2NO3-(aq)+2Ag(s)
Net ionic equation:
Cu(s)+2Ag+(aq)  Cu2+(aq) + 2Ag(s)
Electron transfer half-reactions:
Oxidation Cu(s)  Cu2+(aq) + 2eReduction 2Ag+(aq)+2e-  2Ag(s)
Since e- are not destroyed in chemistry, oxidation is always coupled
to reduction, & the number of electrons released by one element
must be equal to the number of electrons accepted by another
element.
That gives an approach to balance redox rxns. The basis is: the
number of released electrons must be equal to the number of gained
electrons.
Al(s) + CuCl2(aq)  AlCl3(aq) + Cu(s)
2| Al (0) -3e  Al3+
3| Cu2+ +2e  Cu(0)
Cl- does not change its ox.#,
2Al + 3Cu2+  2Al3+ + 3Cu
it is a spectator ion
This is the net ionic equation for the redox rxn
Half-rxns may include H2O, H+ or OH-, giving especially powerful
tool for balancing redox rxns.
KMnO4(aq) + HCl(aq)  MnCl2(aq) + H2O + Cl2↑ + KCl
purple
pinkish
2Cl
2e
 Cl2
5
MnO4- + 8H++5e-  Mn2++ 4H2O (first add water, then H+)
2
10Cl- +2MnO4- + 8H+  5Cl2 + 2Mn2+ + 8H2O
2KMnO4 + 8HCl 2MnCl2+ 8H2O + 5Cl2↑ + 2KCl
KMnO4(aq) + K2SO3(aq)+ KOH K2MnO4(aq) + H2O + K2SO4(aq)
SO32- +2OH- - 2e- SO42- + H2O (first add twice OH- then H2O)
2 MnO4- + e-  MnO422MnO4- + SO32- + 2OH-  2MnO4- + H2O + SO42-
2KMnO4(aq)+K2SO3(aq)+2KOH 2K2MnO4(aq)+H2O+K2SO4(aq)
More examples: P + HNO3 + H2O  NO + H3PO4
Total charge to the left
+0
+3
Total charge to the right
3 P + 4H2O - 5e  H3PO4 + 5H+
5 NO3- +4H+ +3e  NO + 2H2O
3P + 5HNO3 + 2H2O  5NO +3H3PO4
+5
+0
The solution above was ACIDIC solution, i.e. contained an acid, HNO3
Therefore, we operated with H+ & H2O