Download An element is a fundamental substance that cannot be chemically

CHEMISTRY: about 250 years old as a science
Ancient roots
Ancient roots II.
medical roots
common trades:
mining and metalworking
physiological effects of certain substances
finding or making medicines
glassmaking
food processing (wine making, vinegar production)
alchemy
‫ال‬
soap making
origin of word 'chemistry' (.... chimie, Chemie, химия)
Arabic 'al' ~ 'the'
Middle Ages: transformation of ordinary metals
into gold
Chema: ancient province in Egypt famous for its
Sir Isaac Newton (1643-1727)
was also an (unsuccessful)
alchemist
soap production or
kēme (chem) = „earth” in Egyptian
Greek science
Greek science
Democritus (460370 BC)
Aristotle (384322 BC)
philosopher
philosopher and experimental
scientist
atomic theory
'' = indivisble
discrete theory of matter
continuous matter
fundamental elements:
earth, air, water, fire
fundamental properties:
cold, hot, dry, wet
Definition of a chemical element
Robert Boyle (1627-1691)
Examples of elements and non-elements
Iron: element
water: not an element, compound
gold: element
air: not an element, mixture of elements and compounds
An element is a fundamental substance
that cannot be chemically changed or
broken down into anything simpler.
STILL CORRECT AS MODERN DEFINITION
brass or bronze: mixture of elements
rows = periods
columns = groups
noble gases
alkali earth metals
Chemistry as a quantitative science
halogens
end of 18th century: Lavoisier and Lomonosov
Antoine Laurent Lavoisier
(1743-1794)
Mikhail Vasilyevich Lomonosov
(1711-1765)
alkali metals
Chemistry as a quantitative science
Physical properties
end of 18th century: Lavoisier and Lomonosov
extensive 
intensive
accurate measurements (mass, volume, ...)
related to size
e.g. mass
unrelated to size
e. g. temperature
discovery of gases
Quantitative measurements
fundamental

not a combination
of other properties
e.g. length
V = 330 ml
derived
a combination of
other properties
e.g. area
1960: Systéme Internationale d'Unites (SI)
physical property number
unit
(volume)
(milliliter)
(330)
earlier systems: e.g. CGS, MKS
Fundamental properties and units in SI
property
Prefixes in SI
symbol
name of unit
symbol
of unit
m
kilogramm
kg
mass
length
meter
l
m
temperature
T
kelvin
K
amount of substance
n
mole
mol
time
t
second
s
electric current
I
ampere
A
luminous intensity
IV
candela
cd
prefix
symbol
factor
prefix
symbol
factor
giga
G
109
deci
d
101
mega
M
106
centi
c
102
kilo
k
103
milli
m
103
hecto
h
102
micro

106
deka
da
101
nano
n
109
e.g. 1 nm = 109 meter
MOST IMPORTANT DERIVED PROPERTIES
Volume:
m3
CONVERSION OF UNITS
Volume:
dm3
= L (or l)
cm3
= mL (or ml) l (or L)
Density:
kg/m3
1 kmol = 103 mol
m3
1 dm3 = 0.001 m3 = 1000 cm3 = 106 mm3
Density:
g/cm3
kg/m3
Pressure:
Pa (pascal) = kgm-1s-2
Pressure:
atm, bar, torr
Energy:
J (joule) = kgm2s-2
1 g/cm3 = 0.001 kg/ 106 m3 = 1000 kg/m3
Pa (pascal) = kgm-1s-2
1 hPa = 100 Pa 1 MPa = 106 Pa
Energy:
cal
J (joule) = kgm2s-2
1 mJ = 0.001 J
1 nJ = 109 J
Significant figures (s.f.):
the total number of digits in a measurement
Accuracy and Precision:
Correct use:
Accuracy: shows how close to the true value a given
measurement is
Precision: shows how well a number of independent
measurements agree with one another
all digits but the last are certain, the last is a best
guess, usually having an error ± 1
4.803 cm
4
0.00661 g
3
55.220 K
5
34200 m
3-5
SOME NUMBERS ARE EXACT!!
Significant figures (s.f.):
Significant figures and calculations
all digits other than 0 significant
1. Zeros in the middle of a number are always
significant
2. Zeros at the beginning of a number are never
significant
3. Zeros at the end of the number and after the decimal
point are always significant
4. Zeros at the end of a number and before the decimal
point may or may not be significant
1. Multiplication or division: the answer cannot have
more significant numbers than either of the original
numbers
11.78945 g / 11.9 cm3 = 0.991 g/cm3
2. Addition or subtraction: the answer cannot have
more digits to the right of the decimal point than either
of the original numbers
3.18 g + 0.01315 g = 3.19 g
Rounding numbers
5.664525
1. If the first digit you remove is less than 5, round by
dropping it and all following digits.
(3 s.f.)  5.66
2. If the first digit you remove is 6 or greater, round by
adding 1 to the digit on the left
(2 s.f.)  5.7
3. If the first digit you remove is 5 and there are more
nonzero digits following, round up.
(4 s.f.)  5.665
4. If the digit you remove is 5 with nothing following,
round down.
(6 s.f.)  5.66452
LAW OF MASS CONSERVATION
Mass is neither created nor destroyed in chemical
reactions.
TODAY: Not exactly true
Einstein's mass-energy equivalence equation
DE = Dmc2
~100 kJ/mol
~10-12 kg/mol = 1 ng/mol
LAW OF DEFINITE PROPORTIONS
1799 Joseph Louis Proust (1754-1826)
LAW OF MULTIPLE PROPORTIONS
John Dalton (1766-1844)
1803
Different samples of a pure chemical substance always
contain the same proportion of elements by mass.
1.000 g Cu
400 °C
1.252 g CuO
400 °C
HNO3
Cu(NO3)2
Elements can combine in different ways to form
different substances, whose mass ratios are small
whole-number multiples of each other.
Example: nitrogen oxides
Atomic theory:
compound
%N
%O
mO/mN
ratio
N2 O
63.7
36.3
0.571
1
NO
46.7
53.3
1.142
2
N2 O 3
36.9
63.1
1.713
3
NO2
30.4
69.4
2.284
4
N2 O 5
25.9
74.1
2.855
5
John Dalton, 1808
• Elements are made of tiny particles called atoms.
• Each element is characterized by the mass of its
atoms. Atoms of the same element have the same
mass, but atoms of different elements have different
masses.
• Chemical combination of elements to make different
substances occurs when atoms join together in small
whole-number ratios.
• Chemical reactions only rearrange the way that atoms
are combined; the atoms themselves are unchanged.
LAW OF COMBINING GAS VOLUMES
Most elements: atoms and bonds between atoms in a
well defined spatial arrangement
1808
Joseph Louis Gay-Lussac (1778-1850)
Some elements: small independent units containing a
few identical atoms  MOLECULES
H2, N2, O2, F2, P4, S8, Cl2, Br2, I2
AVOGADRO'S LAW
Gases at constant temperature and
pressure react or are produced in a
chemical reaction in volume ratios of
the small whole numbers.
HOW SMALL IS AN ATOM?
WHAT IS THE MASS OF AN ATOM?
Amedeo Avogadro (1776-1856)
1811
Identical volumes of different gases
contain the same number of
molecules at the same temperature
and pressure.
first: relative atomic masses
atomic mass unit (amu): exactly one twelfth of the mass
of a carbon-12 atom (1.660539  1027 kg)
later: numbers of atoms or molecules are conveniently
measured by the mole
1 mol: the number of atoms in exactly 12 g carbon-12
Avogadro's number NA = 6.022  1023 mol-1
ATOMIC NUMBER (Z) =
STRUCTURE OF ATOMS
Number of protons in atom's nucleus =
Number of electrons around atom's nucleus
electron cloud:
electrons
MASS NUMBER (A) =
protons
nucleus
Number of protons (Z) + Number of neutrons (N)
neutrons
Hydrogen:
mass (amu)
electric charge
electron
0.00055
1e
proton
1.00728
+1e
neutron
1.00866
0
mass number
e = 1.602  1019 C
C
1 proton + 1 neutron (deuterium)
1 proton + 2 neutrons (tritium)
Isotopes = atoms with identical atomic numbers but
different mass numbers
Many elements: mixtures of isotopes.
symbol
12
6
1 proton + O neutron (protium)
carbon - 12
Chlorine:
atomic number
75.77% 35Cl
34.969 amu
37Cl
36.966 amu
24.23%
31
15
P
phosphorus - 31
Average atomic mass of chlorine:
235
92
268
109
U
Mt
uranium - 235
meitnerium - 268
0.7577 34.969 amu + 0.2423  36.966 amu = 35.46 amu
The isotopic composition of an element is constant in
nature.
Elements
Compound: a pure substance that is formed when
atoms of two or more different elements combine and
create a new material with properties completely unlike
those of its constituent elements.
Chemical compounds
element
Pure Substances
Matter
@ symbol
compound @ chemical formula
Heterogeneous
Mixtures
NaCl, H2O, C12H22O11, ThO2, BaTiO3, Ca(ClO4)2
Chemical equation:
Homogeneous
2 H2 + O2  2 H2O
reactants
products
Mixture: a simple blend of two or more substances
added together in some random proportion without
chemically changing the individual substances
themselves.
Chemical changes: atoms interact primarily through the
electron structure
Covalent bond: atoms share electrons, bonding
electrons belong to each atom forming the bond
Heterogeneous: the mixing is not uniform, the mixture
has regions of different composition
Molecule: the unit of matter that results when two or
more atoms are joined by covalent bonds
example: water + oil
O2, N2, F2, I2, P4, S8
Homogeneous: the mixing is uniform
example: dust-free air
H2O, NH3, HF, C3H8
C2H6O, C2H5NO2
Ionic bond: transfer of one or more electrons from one
atom to another
Metallic bond: sharing of electrons between a large
number of metal atoms
Na + Cl  Na+ + Cl
cation
Metallic solid: metal atoms connected by metallic
bonding and packed together in a regular way
anion
Ionic solid: anions and cations packed together in a
regular way
Na, Ti, Fe, Cu, In, Pb
NaCl, CsCl, MgCl2, CaO, Ce(SO4)2
Elements: Symbols (Ca for Calcium, Sn for Tin, .....)
Compounds: Formulas
NaCl
formula
for sodium chloride or common salt
chemical name
systematic name
common name
trivial name
Formulas: different types for different purposes
Formulas
empirical formula: shows the ratios of atoms in a
compound
NaCl
CH2O
molecular formula: shows the actual number of atoms
in a molecule
----CH2O
(CH2O)6 = C6H12O6
structural formula: shows how the atoms are connected
CHO
within a molecule
H OH
H
OH
HO
H
HO
H
----HO
C O
H
OH
OH
HO
H
H
OH
H
OH
H
H
CH2OH
Example:
NAMING COMPOUNDS
unknown compound 92.26 % C and 7.74% H
100 g compound: 92.26 g C  92.26 / 12.011 = 7.681 mol C
only inorganic compounds, organic compounds will be
covered as part of the Organic Chemistry course
7.74 g H  7.74 / 1.008 = 7.68 mol H
empirical formula: C7.681H7.681

CH
if M = 26 g/mol
C 2H 2
(acetylene)
if M = 78 g/mol
C 6H 6
(e.g. benzene)
NOMENCLATURE = system of names, naming rules
present chemical nomenclature:
discussed and approved by IUPAC
percent composition: (experimental) elemental analysis
International Union of Pure and Applied Chemistry
molecular mass: (experimental) mass spectrometry
BINARY IONIC COMPOUNDS
BINARY IONIC COMPOUNDS
CrCl3: chromium(III) chloride
name = positive ion + negative ion
Cr can form both Cr2+ and Cr3+ ions!
positive ion = name of the element
PbS:
negative ion = name of the element + suffix -ide
lead(II) sulfide
FeCl2: iron(II) chloride  ferrous chloride
KCl:
potassium chloride
FeCl3: iron(III) chloride  ferric chloride
LiF:
lithium fluoride
SnCl2: tin(II) chloride  stannous chloride
SnCl4: tin(IV) chloride  stannic chloride
BaCl2: barium chloride
Cu2O: copper(I) oxide  cuprous oxide
Ba always forms Ba2+ ions
AlBr3: aluminum bromide
CuO:
copper(II) oxide  cupric oxide
Al always forms Al3+ ions
BINARY MOLECULAR COMPOUNDS
name = more positive part + more negative part
(suffix –ide)
use of Greek numbers in names:
BINARY MOLECULAR COMPOUNDS
N2O: dinitrogen monoxide
HF:
hydrogen fluoride
CO:
carbon monoxide
CO2: carbon dioxide
1 --- (mono)
5 penta
2 di
6 hexa
SeBr4: selenium tetrabromide
3 tri
7 hepta
PCl5: phosphorus pentachloride
4 tetra
8 octa
AsCl3: arsenic trichloride
SF6:
sulfur hexafluoride
IF7:
iodine heptafluoride
NO:
nitrogen monoxide
N2O3: dinitrogen trioxide
NO2: nitrogen dioxide
N2O4: dinitrogen tetroxide
N2O5: dinitrogen pentoxide
SPECIAL CASES
POLYATOMIC IONS
names must be learnt e.g.
CO32- carbonate ion
SO42- sulfate ion
OsO4: osmium tetroxide or osmium(VIII) oxide
NO3-
PO43- phosphate ion
CrO3: chromium trioxide or chromium(VI) oxide
NH4+ ammonium ion
Some metal compounds (often oxides) are molecular
nitrate ion
MnO2: manganese dioxide or manganese(IV) oxide
(NH4)2SO4 ammonium sulfate
Very rare exceptions:
Fe(NO3)2
iron(II) nitrate or
ferrous nitrate
V2O5: vanadium pentoxide or vanadium(V) oxide
Hg(NO3)2 mercury(II) nitrate or mercuric nitrate
P2O5: phosphorus pentoxide
KCN
potassium cyanide (CN- cyanide ion)
Balancing chemical equations
Chemical equations
conservation of mass and conservation of charge
molecular equation:
3CaCl2 + 2Na3PO4  Ca3(PO4)2 + 6NaCl
AgNO3 + NaCl  AgCl + NaNO3
ionic equation (aqueous phase = in water):
Ca:
3
=
Cl:
3×2 = 6
=
2×3 = 6
=
3
6
Ag+ + NO3  + Na+ + Cl  AgCl + Na+ + NO3
Na:
P:
2
=
2
6
Ag+ + Cl  AgCl
O:
2×4 = 8
=
2×4 = 8
charge:
0
=
0
Balancing chemical equations
Balancing chemical equations
Step 1: Write reactants and products
C3H8 + O2
S:
2S2O32 + I2

2×2 = 4
=
O:
I:
2×3 = 6
2
=
=
S4O62 + 2I
4
C3H8 + O2
6
2×1 = 2
2×-2 = -4
=
-2+2×-1 = -4
CO2 + H2O

CO2 + H2O
Step 3: Find coefficients to balance this atom
1C3H8 + O2
charge:

Step 2: Find one atom that occurs only in one
substance on both sides
 3CO2 + H2O
Step 4: Find another unbalanced atom which
occurs in only one substance
1C3H8 + O2
 3CO2 + H2O
Stoichiometry
Step 5: Find coefficient to balance this atom
1C3H8 + O2
 3CO2 + 4H2O
(Greek) στοιχεῖον = element and μέτρον = measure
Step 6: Repeat steps 4-5 until you have
balanced all atoms
1C3H8 + O2
 3CO2 + 4H2O
1C3H8 + 5O2
 3CO2 + 4H2O
quantitative relationship between reacting or produced
substances (mass, amount of substance)
C2H4 (g) + HCl (g)  C2H5Cl (g)
Step 7: Make sure the coefficients are reduced
to the smallest whole number values
1C3H8 + 5O2
1 mol hydrogen chloride to produce
 3CO2 + 4H2O
1 mol ethyl chloride
Step 8: Check the balanced equation
1C3H8 + 5O2
3C, 8H, 10O
gas
1 mol ethylene reacts with
28.02 g ethylene reacts with
 3CO2 + 4H2O
36.46 g hydrogen chloride to produce
3C, 10O, 8H
64.48 g ethyl chloride
Limiting amounts of reactants
C2H4 (g) + HCl (g)  C2H5Cl (g)
How much hydrogen chloride reacts with 15.00 g
ethylene?
15.00 g ethylene = 15.00g/28.02gmol-1 = 0.5353 mol
C2H4 (g) + HCl (g)  C2H5Cl (g)
What happens if we have 5 mol C2H4 and 3 mol HCl at the
beginning?
5 mol C2H4 would react with 5 mol HCl  HCl not enough
0.5353 mol ethylene reacts with
3 mol HCl would react with 3 mol C2H4  C2H4 too much
0.5353 mol hydrogen chloride
After reaction: 3 mol C2H5Cl and 2 mol C2H4 unchanged
0.5353 mol hydrogen chloride =
HCl: limiting reagent, completely consumed
0.5353 mol × 36.46g/mol = 19.51 g
C2H4: excess reagent, some remains unreacted
Yields of chemical reactions
C4H8 (g) + CH4O (l)  C5H12O (l)
Stoichiometry: calculated for an IDEAL case when the
limiting reagent is completely converted to the desired
product
Reality:
1. products other than the desired one may
form (side reactions)
2. some of the limiting reagent may remain
unchanged
YIELD: the amount of desired product actually formed
Actual yield of desired product
PERCENT YIELD =
Theoretical yield of desired product
liquid
From 26.3 g of isobutylene and excess methanol 32.8 g of
methyl tert-butyl ether is prepared. What is the percent
yield?
26.3 g C4H8 (56.12 g/mol)  0.469 mol
Theoretical yield of C5H12O:
0.469 mol × 88.17 g/mol = 41.35 g
PERCENT YIELD = 32.8 g / 41.35 g = 79.3 %
× 100%
A chemical equation
HCl (g)  in gas phase
does show - identity of reactants and products
- stoichiometry
CH4O (l)  in liquid phase
- direction of spontaneous reaction
2H2 + O2  2H2O
SiO2 (s)  in solid phase
may show
NaCl (aq)  in aqueous phase = dissolved in water
or
2H2 + O2 = 2H2O
- phase of reactants and products
- energy change of the reaction
2H2 (g) + O2 (g)  2H2O (g) DE = 484 kJ/mol
does not usually show
CaCO3 (s) + 2CH3COOH (l) 
- time needed for completion
Ca(CH3COO)2 (aq) + H2O (l) + CO2 (g)
- conditions (e.g. temperature)
EXAMPLE:
SOLUTION
Coca-Cola
=
SOLUTE(S)
+
SOLVENT
carbon dioxide
water
sugar
5.0 g glucose dissolved in water to make 100 mL solution.
5.0 g C6H12O6 / 180.18 g/mol = 0.028 mol
molarity: 0.028 mol / 0.100 L = 0.28 M
(many) other substances
AgNO3 + NaCl  AgCl + NaNO3
concentration  various units
Equivalent quantities: reacting exactly with each other
molarity (M): most important in chemistry
1.0 L of 1.0 M AgNO3 and 1.0 L of 1.0 M NaCl
1.0 L of 1.0 M AgNO3 and 2.0 L of 0.5 M NaCl
Moles of solute
MOLARITY =
Volume of solution
unit: molL-1
mol/dm3
M
DILUTION:
4.0 L of 0.25 M AgNO3 and 1.0 L of 1.0 M NaCl
0.40 L of 0.10 M AgNO3 and 0.20 L of 0.20 M NaCl
GROUPS OF CHEMICAL REACTIONS
concentrated solution + solvent  dilute solution
Based on the number of reagents and products
(mainly in organic chemistry)
Example: to 50 mL 0.28 M glucose solution enough water
given to make a final volume of 1000 mL
50 mL 0.28 M solution 
0.050 L × 0.28 molL-1 = 0.014 mol solute
final volume: 1.000 L
final molarity: 0.014 M
initial molarity initial volume
Moles of solute = Molarity × Volume = Mi × Vi = Mf  Vf
final molarity final volume
- isomerization: NH4OCN  H2NCONH2
- addition:
CH2=CH2 + Cl2  CH2ClCH2Cl
special cases: dimerization, trimerization,
polymerization: 3C2H2  C6H6
- elimination/decomposition: NH4Cl  NH3 + HCl
- substitution: C2H5Cl +NaOCH3  C2H5OCH3 + NaCl
GROUPS OF CHEMICAL REACTIONS
Based on the observable change
GROUPS OF CHEMICAL REACTIONS
element, molecules formed by atoms
2Mg (s) + O2 (g)  2MgO (s)
- gas formation: Zn + 2HCl  ZnCl2 + H2
- color change: 2KBr + Cl2  2KCl + Br2
element, atom
+2 ion -2 ion
electron transfer occurs: REDOX REACTION
- precipitation: ZnCl2 + H2S  ZnS + 2HCl
- combustion: CH4 + 2O2  CO2 + 2H2O
NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (l)
+1 ion
-1 ion
+1 ion
-1 ion
no electron transfer occurs: ACID-BASE REACTION
REDOX REACTIONS
reduction – oxidation reaction
reduction: the gain of one ore more electrons by a
substance – element, compound, ion
oxidation: the loss of one ore more electrons by a
substance
4Fe (s) + 3O2 (g)  2Fe2O3 (s)
oxidation of iron
reduction of oxygen
EARLIER DEFINITIONS
still used widely in organic chemistry
reduction: removal of oxygen (O) or addition of
hydrogen (H)
oxidation: addition of oxygen (O) or removal of
hydrogen (H)
also reaction with oxygen element (O2)
2Fe2O3 (s) + 3C (s)  4Fe (s) + 3CO2 (g)
reduction of iron
oxidation of carbon
reduction
Fe2+ + 2e–  Fe
Fe3+ + e–  Fe2+
Oxidation number (oxidation state)
a formal number assigned to every atom in a compound
or element, which indicates whether the atom is neutral,
electron-rich or electron-poor
Fe3+ + 3e–  Fe
– useful tool for classifying compounds and balancing
equations
Ag+ + e–  Ag
– does not necessarily imply actual charges on atoms
Cl2 + 2e–  2Cl–
O2 + 4e–  2O2–
MnO4– + 8H+ + 5e–  Mn2+ + 4H2O
oxidation
ELECTRONEGATIVITY (EN):
a dimensionless number showing the ability of an atom
in a molecule to attract the shared electrons in a
covalent bond
F: 4.0
O: 3.5
Cl: 3.0
H: 2.1
Al: 1.5
K: 0.8
(p. 248)
Rules for assigning oxidation numbers
1. An atom in its elemental state has an oxidation
number of 0.
Na, H2, Br2, F2, O2, Ne, S8, P4
oxidation number 0
2. An atom in a monatomic ion has an oxidation
number identical to its charge.
Na+: +1 Ca2+: +2 Al3+: +3 Cl–: -1 O2–: -2
3. In a heteronuclear covalent bond, the bonding
electron pair is assigned to the more electronegative
atom.
H–Cl
H: +1 Cl: -1
H–O–H
AlBr3 Al: +3 Br: -1
8. Hydrogen can have an oxidation number of either
+1 or -1 depending on the electronegativity of the
other element forming the compound.
CaH2
Ca: +2 H: -1
NH4+ N: -3 H: +1
H2SO3 H: +1 S: +4 O: -2
5. Fluorine always has an oxidation number of -1 in
its compounds.
XeF2 Xe: +2 F: -1
SF6 S: +6 F: -1
6. Alkali metals always have an oxidation number of
+1, alkaline earth metals +2 in their compounds.
LiH
Li: +1 H: -1
BaO2
Ba: +2 O: -1
H: +1 O: -2
7. Aluminum usually has an oxidation number of +3.
LiAlH4 Li: +1 Al: +3 H: -1
4. The sum of the oxidation numbers is 0 for a neutral
compound and is equal to the net charge for a
polyatomic ion.
H2S H: +1 S: -2
9. Oxygen usually has an oxidation number of -2.
XeO2 Xe: +4 O: -2
SO2 S: +4 O: -2
10. Halogens have an oxidation number of -1 in all
compounds not containing oxygen or another halogen.
FeCl3 Fe: +3 Cl: -1
FeBr2 Fe: +2 Br: -1
Other cases:
NaClO2 Na: +1 Cl: +3 O: -2
BrO2 Br: +4 O: -2
HClO H: +1 Cl: +1 O: -2
BrCl Br: +1 Cl: -1
IF5 I: +5
ClO4– Cl: +7 O: -2
ClF3 Cl: +3 F: -1
F: -1
Notable exceptions:
oxygen fluorides OF2 O: +2 F: -1
peroxides: H2O2 H: +1 O: -1
superoxides: KO2
BaO2 Ba: +2 O: -1
There are cases when more then one assignment can
be used without problems.
Fe3O4  Fe: +8/3 O: -2
 Fe2: +3 Fe1: +2 O: -2
K: +1 O: -½
REDUCING AGENT:
0
0
+3
• causes reduction
-2
4Fe (s) + 3O2 (g)  2Fe2O3 (s)
• loses one or more electrons
iron: 0  +3 undergoes oxidation REDUCING AGENT
• undergoes oxidation
oxygen: 0  -2 undergoes reduction OXIDIZING AGENT
• oxidation number of atom increases
+3
-2
0
0
+4 -2
2Fe2O3 (s) + 3C (s)  4Fe (s) + 3CO2 (g)
OXIDIZING AGENT:
iron: +3  0 undergoes reduction OXIDIZING AGENT
• causes oxidation
carbon: 0  +4 undergoes oxidation REDUCING AGENT
• gains one or more electrons
oxygen: -2  -2 no change
• undergoes reduction
• oxidation number of atom decreases
+1 +7 -2
+1 +4 -2
+1 +6 -2
+1 +7 -2
2 KMnO4 + 5 Na2SO3 + 3 H2SO4 
+2 +6 -2
+1 +6 -2
+1 +4 -2
+1 -2
2 KMnO4 + 3 Na2SO3 + 1 H2O 
+1 +6 -2
+1 -2
+4 -2
2 MnSO4 + 5 Na2SO4 + 1K2SO4 + 3 H2O
+1 +6 -2
+1 -2 +1
2 MnO2 + 3Na2SO4 + 2 KOH
Mn: +7  +2 decrease, oxidizing agent
Mn: +7  +4 decrease, oxidizing agent
S in Na2SO3: +4  +6 increase, reducing agent
S: +4  +6 increase, reducing agent
-5 change
2 mol needed
Mn:
-3 change
2 mol needed
S in Na2SO3: +2 change
Mn:
5 mol needed
S:
+2 change
3 mol needed
+1 +7 -2
+1 +4 -2
+1 -2 +1
2 KMnO4 + 1 Na2SO3 + 2 KOH 
+1 +6 -2
+1 +6 -2
+1 -2
2 K2MnO4 + 1Na2SO4 + 1H2O
Fe (s) + Cu2+ (aq)  Fe2+ (aq) + Cu (s)
Fe: 0  +2 increase, reducing agent
Cu: +2  0 decrease, oxidizing agent
Mn: +7  +6 decrease, oxidizing agent
S: +4  +6 increase, reducing agent
Mg (s) + 2H+ (aq)  Mg2+ (aq) + H2 (g)
Mg: 0  +2 increase, reducing agent
H: +1  0 decrease, oxidizing agent
Mn:
-1 change
2 mol needed
S:
+2 change
1 mol needed
Activity series
of elements:
strongly reducing
more active
weakly reducing
less active
Li
K
Ca
Na
Mg
Al
Zn
Cr
Fe
Co
Sn
H2
Cu
Ag
Hg
Pt
Au
react with H2O to
produce H2
Au (s) + 3Ag+ (aq) 
X Au3+ (aq) + 3Ag (s)
does not occur
strongly reducing
more active
do not react with H2O
react with H+(aq) to
produce H2
do not react with H2O
or H+(aq)
weakly reducing
less active
Li
K
Ca
Na
Mg
Al
Zn
Cr
Fe
Co
Sn
H2
Cu
Ag
Hg
Pt
Au
Fe (s) + Cu2+ (aq) 
Fe2+ (aq) + Cu (s)
Mg (s) + 2H+ (aq) 
Mg2+ (aq) + H2 (g)
Au (s) + 3Ag+ (aq) 
X
Ca (s) + Sn2+ (aq) 
Ca2+ (aq) + Sn (s)
Co (s) + Zn2+ (aq) 
X
SIMPLE APPLICATIONS OF REDOX REACTIONS
more active metal + ion of less active metal 
ion of more active metal + less active metal
less active metal + ion of more active metal 
X
more reducing metal + ion of less reducing metal 
ion of more reducing metal + less reducing metal
Combustion: burning of fuel by oxidation with oxygen in
air. Fuel: gasoline, natural gas, fuel oil, wood, carbon
CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (l)
Bleaching: redox reactions to decolorize or lighten
colored materials (hair, wood pulp, clothes)
Bleaching agent = oxidizing agent
less reducing metal + ion of more reducing metal X

Batteries: providing electrical current using spontaneous
chemical reactions. Redox reactions are needed based
on common and inexpensive materials
Zn (s) + 2MnO2 (s) + 2NH4Cl (s) 
usually H2O2, NaOCl, Cl2
Corrosion: deterioration of a metal by oxidation, rusting,
usually oxygen from air and moisture is needed
4Fe (s) + 3O2 (g) + 2H2O (l)  2Fe2O3H2O (s)
iron rust
ZnCl2 (aq) + Mn2O3 (s) + 2NH3 (aq) + H2O (l)
Metallurgy: the science of extracting and purifying
metals from their ores, always based on redox reactions
ZnO (s) + C(s)  Zn (s) + CO (g)
Respiration: the process of breathing and using oxygen
for the many biological redox reactions that provide the
energy needed by living organisms.
C6H12O6 + 6O2  6CO2 + 6H2O + biological energy
glucose (a sugar)
Simple acid-base reactions
- precipitation
- gas formation
- neutralization
Electrolytic dissociation
C12H22O11 (s)
H2O
C12H22O11 (aq)
sucrose molecules
no ions in water: NONELECTROLYTE
does not conduct electricity
- complex formation
very important property: electrolytic dissociation
NaCl (s)
H2O
Na+ (aq) + Cl (aq)
common salt: ionic solid
ions in water: ELECTROLYTE
conducts electricity
Svante August
Arrhenius (1859-1927)
HCl (aq)
H+ (aq) + Cl (aq)
0.1%
Strong electrolytes
'complete dissociation'
HCl, HBr, HI
STRONG ELECTROLYTE
CH3COOH (aq)
Weak electrolytes
Nonelectrolytes
99.9%
CH3COO (aq) + H+ (aq)
99%
1%
CH3OH
HNO3
HCOOH
C2H5OH
H2SO4
H3PO4
NaCl
WEAK ELECTROLYTE
NaOH, KOH
precipitation
neutralization
AgNO3 + NaCl  AgCl + NaNO3
HCl + NaOH  NaCl + H2O
precipitate
FeCl3 + 3NaOH  Fe(OH)3 + 3NaCl
gas formation
CH3COOH + NH3  NH4CH3COO
2CsOH + H2SO4  Cs2SO4 + 2H2O
complex formation
FeS + 2HCl  FeCl2 + H2S (g)
AgCl + 2NH3  [Ag(NH3)2]Cl
Na2CO3 + H2SO4  Na2SO4 + H2O + CO2 (g)
HgO + 4KI + H2O  K2[HgI4] + 2KOH
NH4Cl + NaOH  NaCl + H2O + NH3 (g)
Zn2+ + 4OH  [Zn(OH)4]2
SOLUBILITY:
No simple rules for most of the compounds.
A compound is probably soluble in water if it contains
one of the following cations:
 Li+, Na+, K+, Rb+, Cs+
 NH4+
A compound is probably soluble in water if it contains
one of the following anions:
 Cl, Br, I except that of Ag+, Hg22+, and Pb2+
 NO3, ClO4, CH3COO
 SO42 except BaSO4, Hg2SO4, PbSO4, CaSO4
H2O
CH3COOH
KBr
'partial dissociation'
Pb(NO3)2 + H2S  PbS + 2HNO3
HF
HClO4
C12H22O11