Unit 2 CHEMISTRY - 100 facts
Acids and Bases
1.
2.
3.
4.
5.
Examples of acids include hydrochloric acid (brick cleaning), sulfuric acid (car battery), nitric
acid (manufacture of dyes, fertilisers, explosives), acetic acid (vinegar), carbonic acid
(carbonated drinks), phosphoric acid (fertilisers), citric acid (citrus fruits), and lactic acid
(milk products, muscles during exercise).
Indicators are plant extracts that change colour with acids and bases. Eg. litmus,
phenolphthalein, methyl orange, universal indicator.
Properties of acids include sour taste, corrode metals, change colours of indicators (turn blue
litmus red), neutralise bases.
Some examples of bases are sodium hydroxide (drain and oven cleaners), ammonia
(household cleaning products), calcium hydroxide (plaster and cement) and magnesium
hydroxide (antacid preparations).
Properties of bases include bitter taste, feel soapy, change colour of indicators (turn red litmus
blue) and neutralise acids.
6. Acid reactions:
a. acid + metal hydroxide  salt + water
HCl(aq) + NaOH(aq)
 NaCl(aq)
+ H2O(l)
b) acid + metal oxide  salt + water
H2SO4(aq) + CuO (s)  CuSO4(aq) + H2O(l)
c) acid + metals (most)  salt + hydrogen
2HCl(aq) + Zn (s)  ZnCl2(aq) + H2(g)
d) acid + carbonate  salt + carbon dioxide + water
2HCl (aq) + CaCO3(s)  CaCl2(aq) + CO2(g) + H2O(l)
e) acid + metal hydrogen carbonate  salt + carbon dioxide + water
HNO3(aq) + NaHCO3(aq)  NaNO3(aq) + CO2(g) + H2O(l)
f) acid + sulfide  salt + hydrogen sulfide
H2SO4(aq) + FeS (s)  FeSO4(aq) + H2S(g)
g) acid + sulfite  salt + sulfur dioxide + water
H2SO4(aq) + Na2SO3 (aq)  Na2SO4(aq) + SO2 (g) + H2O (l)
7. Arrhenius acid/base concept: Acids dissociate in water to produce hydrogen H+ ions; bases
dissociate in water to produce hydroxide OH- ions.
8. Lowry/ Bronsted acid/base concept: An acid is a proton donor; a base is a proton acceptor.
HCl (aq) + H2O (l)  H3O+ (aq) +
Cl-(aq)
Acid
base acid
base
9. HCl / Cl- and H3O+ / H2O are conjugate acid / base pairs. They differ in their formula by a single
proton (H+) only.
10. H+ ion (or proton) is hydrated in solution to form the hydronium ion, H3O+(aq).
11. Amphiprotic substances are able to act as either an acid or a base. Eg. Water, hydrogen
carbonate (HCO3-) ion.
12. A polyprotic acid is able to donate more than one proton. E.g. sulfuric acid (diprotic H2SO4),
phosphoric acid (triprotic H3PO4).
13. Strong acids / bases readily donate / accept protons. E.g. HCl, HNO3, H2SO4 are strong
acids. They ionise completely. NaOH is a strong base.
14. Weak acids / bases are only partially ionised in water. E.g. acetic acid (CH3COOH) and
ammonia (NH3).
15. A concentrated acid / base contains a large quantity of the acid / base dissolved in water. A
dilute acid / base contains a small quantity of the acid / base dissolved in water.
16. pH is a measure of the level of acidity. Acid solutions have a pH  7; neutral solutions have
a pH = 7; basic solutions have a pH  7. pH = - log10 [ H3O+]
17. The self ionisation constant for water, Kw = [ H3O+] x [ OH-] = 10 -14
18. Volumetric analysis involves a titration to determine the exact concentration of a solution.
19. A primary standard is a solution whose concentration is known exactly.
Molarity and Amount of Substance
20. The mole is the amount of substance that contains the same number of particles as there are in
12 g of C-12.
21. Avogadro's number (NA) = 6.02 x 10 23; is the number of atoms in 12 g of C-12, is the
number of particles in one mole of a substance.
22. Number of particles = amount of substance (mol) x number of particles in 1 mole (6.02 x
10 23)
=
n x NA
23. The molar mass of a substance (M)is the mass of one mole of that substance measured in units
of g mol-1. The molar mass is the relative atomic mass (Ar)or relative molecular mass (Mr)
measured in units of g mol-1.
24. Amount of substance = mass / molar mass
n (mol) =
m (g)/ M (g mol-1)
25. The molarity of a solution is the number of mole of solute in 1 litre of solution
24. Amount of substance = concentration x volume of solution in L
n (mol) = c (mol L -1) x V (L)
26. When working with diluted solutions:
c1V1 = c2V2 (as the amount of substance, mole, remains constant)
Quantities in Chemistry
27. The relative isotopic mass of an isotope is the mass of an atom of the isotope relative to the
mass of an atom of C-12, taken as 12 units exactly. eg. for Cl - 35, RIM is 34.969.
28. The relative atomic mass of an element (Ar) is the weighted average of the relative masses of
the isotopes (RIM) of the element on the C - 12 scale. As Ar is a relative measure, it has no
units.
Eg. Ar(Cl) = (% abundance of Cl-35 x isotopic mass Cl -35) + (% abundance of Cl-37 x
isotopic mass Cl -37)
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100
29. A mass spectrometer can be used to determine the relative masses of isotopes of an element.
30. The relative molecular mass (Mr) is the sum of the relative atomic masses (Ar) of the
elements in the molecular formula. For non-molecular (eg. ionic) compounds the term
relative formula mass (Mr) is used.
31. The molar mass (M) of an element or compound is the relative molecular mass of the element
or compound expressed in grams. Its units are g mol-1.
32. Percentage composition:
% by mass of an element =
mass of element in 1 mole of the compound x 100
--------------------------------------------mass of 1 mole of compound
=
Ar
x number in formula x 100
----------------------Mr
33. The empirical formula of a substance is the simplest whole number ratio of the atoms present.
To determine the empirical formula of a compound, an experimentally determined ratio by
mass must be converted to a ratio by numbers of atoms (using the amount (mol) of each
element).
34. The molecular formula shows the actual composition of the compound. A molecular formula
is always a whole number multiple of the empirical formula. It can be obtained from the
empirical formula if the molar mass of a compound is known.
35. Stoichiometry is the calculation of the quantity of one of the reactants or products involved in
a reaction given the quantity of one of the other reactants or products. Quantities may be
given in amount of substance (mol), mass (g, kg) or volume of solution of a known
concentration.
36. Steps for solving a stoichiometric problem:
1.
Write a full balanced equation for the reaction. Write the relevant information from the
question either above or below the equation.
2. Calculate the amount (mol) of the known substance
Amount = mass / molar mass
or
Amount = concentration x volume
3.
Use the equation to determine the ratio of the amount of required substance / amount of
known substance
4. Calculate the required quantity (mass, volume, concentration)
Mass = amount x molar mass or
Volume = amount / concentration
Gases of the Atmosphere
37. The small fraction of the earth's environment that supports life is known as the biosphere.
Water is essential for the support of life (unit 1) as are the gases of the atmosphere.
38. Photosynthesis is the process in which green plants convert carbon dioxide and water to
glucose and oxygen. Chlorophyll (a green pigment) and sunlight are conditions necessary for
this. This process removes CO2 from the atmosphere and adds O2 to it.
6 CO2(g) + 6 H2O(l)  C6H12O6(aq) +
6 O2(g)
39. Glucose (C6H12O6) formed by plants during photosynthesis undergoes condensation
polymerisation to form starches and cellulose.
n C6H12O6 (aq)  -(- C6H12O6-)- n +
n H2O(l)
starches and cellulose
40. In respiration, glucose, which is produced in photosynthesis, reacts with oxygen to produce
carbon dioxide and water.
C6H12O6(aq) +
6 O2(g)  6 CO2(g) + 6 H2O(l)
41. Nitrogen fixation is any chemical process in which atmospheric nitrogen is converted into a
form that can eventually be used by plants to manufacture organic compounds. Nitrogen
fixing bacteria in the soil, lightning, meteor trails, car engines and industrial combustion all
produce high enough temperatures to cause nitrogen fixation.
42. Combustion is a process where a substance rapidly reacts with oxygen to produce heat and
light. This commonly adds carbon dioxide and often water and sulfur dioxide to the
atmosphere.
C(s) +
O2(g)  CO2(g)
43. Atmospheric pollution includes the destruction of the ozone layer by CFS's, global warming
caused by an increase in "greenhouse gases" (CO2, nitrogen oxides, CH4, CFC's), acid rain
(caused by nitrogen oxides, sulfur oxides and CO2) and photochemical smog (caused by
nitrogen oxides, unburnt hydrocarbons, ozone and PAN).
Oxygen, Carbon Dioxide, Carbon Monoxide
44. The atmosphere contains approximately 21% oxygen, 78% nitrogen, 1% argon and <1% of
other gases.
45. Oxygen is a colourless, odourless gas which is slightly soluble in water (supports aquatic life)
and supports combustion.
46. Oxygen is a part of many chemical processes such as the manufacture of steel, sulfuric acid,
medical industry, sewage treatment and combustion of fossil fuels in power stations and
industry.
47. Oxygen is produced commercially by the fractional distillation of air and in the laboratory by
the decomposition of hydrogen peroxide, H2O2, using manganese dioxide as a catalyst.
2 H2O2(aq)  2 H2O(l) + O2(g)
48. Carbon dioxide is a colourless, odourless gas that is slightly soluble in water (forming
carbonated beverages). It is denser than air (will extinguish a flame) and does not support
combustion. Solid carbon dioxide ("dry ice") does not melt, but sublimes directly into a gas
when heated.
49. Carbon dioxide is used to make soft drinks, as a refrigerant (dry ice), in fire extinguishers and
to preserve food (many harmful organisms cannot survive in a carbon dioxide environment.)
50. Carbon dioxide is produced commercially by the fermentation of sugars to alcohol and heating
calcium carbonate (limestone). In the laboratory, carbon dioxide is produced by the reaction
of dilute hydrochloric acid on marble chips (often using Kipps apparatus).
51. Both oxygen and carbon dioxide are cycled in the biosphere in the carbon-oxygen cycle.
52. Carbon monoxide is produced by the incomplete combustion of fuels. Carbon monoxide
produced in the atmosphere is eventually oxidised to carbon dioxide.
53. Haemoglobin combines with the carbon monoxide in blood more readily than with oxygen.
This causes shortness of breath, dizziness, a ruddy pink complexion and ultimately loss of
consciousness and death due to oxygen starvation of the cells.
Nitrogen
54. Nitrogen is a colourless, odourless gas that is insoluble in water. It is very unstable and
unreactive.
55. Nitrogen is needed by plants (but is only in a useful form as soluble nitrate ions NO3- and
ammonium ions NH4+). It is used by the fertiliser industry, wherever an unreactive oxygen
free environment is needed (electronics industry, glass-making, welding, to remove flammable
vapours from oil pipelines, refuelling of aircraft during flight) and to snap freeze foods (b.p. =
- 196  C), cool superconductors and store frozen eggs, sperm and embryos.
56. Nitrogen is produced commercially by the fractional distillation of air. It is too unreactive to
be produced in the laboratory.
57. Nitrogen forms several oxides. Two of these are nitric oxide (nitrogen (II) oxide) which is
colourless and forms from nitrogen and oxygen mixing at very high temperatures (e.g.
lightning flashes and meteor trails) and nitrogen dioxide (nitrogen (IV) oxide) which is brown,
toxic and forms from the oxidation of nitric oxide. Nitrogen dioxide dissolves in water to
form a mixture of nitric and nitrous acids.
N2(g) + O2(g)  2 NO (g)
2 NO (g) + O2(g)  2 NO2(g)
2 NO2(g) + H2O(l)  HNO3(aq) + NHO2(aq)
58. Oxygen will re-light a glowing splint of wood. Carbon dioxide will extinguish a lighted
match and turn limewater milky. Hydrogen will react explosively (with a "pop") when
exposed to oxygen and an ignition source.
Noble Gases
59. The inert or noble gases are found in group VIII (He, Ne, Ar, Kr, Xe, Rn) of the periodic
table. All of these elements exist as discrete atoms, have low melting and boiling
temperatures and are generally chemically inert. Only xenon and krypton form compounds,
but they do not do so readily.
60. Helium is found in the sun. It is used in balloons, airships and mixed with oxygen for divers,
welding and rocket fuel. Neon is used in neon signs, fluorescent lamps and high-voltage
testers. Argon is used in electric light bulbs, radio tubes, Geiger counters, in arc-welding and
in the production of semi-conductor crystals (e.g. Si and Ge). Krypton is used in high speed
photography, fluorescent lamps and Kr-85 (radioactive) is used to detect flaws in metal
surfaces. Xenon is used in strobe lights, lamps for high-speed photography, in lasers and as an
anaesthetic. Radon is used in radiotherapy to treat cancer.
Kinetic Molecular Theory
61. The kinetic molecular theory describes gas particle behaviour. Gas particles move in constant
straight-line motion. Most of the volume of a gas is empty space. Particles are very far apart
and have no forces acting between them. As the temperature increases, the average kinetic
energy of particles increases and so does their speed. All collisions between particles and
with the walls of the container are elastic (no energy is lost).
62. The average kinetic energy of a gas is not dependent on the particular gas. Within each gas
there will be particles with very low kinetic energy and particles with very high energy. The
average kinetic energy is related to the speed of the particles (K.E. = ½ mv2). As the average
kinetic energy of particles in different gas samples is the same at a given temperature, the
lighter gas particles must have a greater average velocity.
63. Diffusion is the term used to describe the way each gas in a mixture of gases spreads itself
evenly to fill the total volume available. The rate at which gases diffuse depends on the
average velocity of the particles and consequently on the molecular mass and the temperature
of the gas.
64. The kinetic molecular theory explains why a gas (odour) will diffuse through a space, gases
expand to fill a container and gases can be compressed.
The Gas Laws
65. The pressure exerted by a gas depends on the number of collisions between the molecules and
the walls of the container. In a mixture of gases, the total pressure is the sum of the partial
pressures exerted by each of the gases alone, as the molecules of each gas move independently
of one another. The pressure will increase if the amount of gas increases, the temperature of
the gas is increased or the volume of the container is decreased.
66. Pressure = force / area. SI units for pressure are newtons per square metre (N m-2) or pascal
(Pa). Pressure can also be measured in units of millimetres of mercury (mm Hg), atmosphere
(atm) and bar. 1 atmosphere is the average atmospheric pressure at sea level. 1 atm = 760
mm Hg
= 101 325 Pa
= 101 325 N m-2
= 101.325 kPa
1 bar = 100 kPa
(hectopascals, used on weather maps) 1 hPa = 100 Pa
67. Boyle's Law: For a fixed mass of gas at constant temperature, the volume of the gas is
inversely proportional to the pressure.
V
 1/p
PV = k
68. Charles' Law: For a fixed mass of gas at constant pressure, the volume of the gas is directly
proportional to the absolute temperature (in degrees Kelvin). V  T
V/T= k
69. The absolute zero of temperature is 0 K or - 273 C. At 0 K all molecules have zero kinetic
energy.
Kelvin temperature = Celsius temperature - 273
70. Avogadro: Equal volumes of gases at the same temperature and pressure contain equal
number of particles.
Vn
V1 / V2 = n1 / n2
71. The molar volume of a gas is the volume occupied by one mole of the gas. 1 mole of any
gases at the same temperature and pressure will occupy a fixed volume, regardless of the
particular temperature and pressure.
72. At standard laboratory conditions (SLC, 25C and 101.3 kPa), the molar volume of any gas is
24.5 L mol-1. At standard temperature and pressure (STP, 0C and 101.3 kPa), the molar
volume of any gas is 22.4 Lmol-1
73. A combination of Boyle's and Charles' Laws leads to: p1V1 / T1 = p2V2 / T2
p and V may be in any units, provided these are consistent. T is measured in Kelvin.
74. A combination of Boyle's, Charles' and Avogadro's Laws lead to the general gas equation:
pV = nRT
R = universal gas constant = 8.31 J K-1 mol-1 (p in Pa, V in m3, T in
Kelvin)
= 8.31 J K-1 mol-1 (p in kPa, V in L, T in Kelvin)
75. A gas that behaves according to the general gas equation is said to be an ideal gas. Most gases
are ideal at low pressures and high temperatures.
76. Stoichiometry can be used to determine the quantities of gases involved in reactions. The
ratios in a balanced equation can be used to calculate volumes of gaseous reactants or
products. If the pressure and temperature are constant over the reaction, then volume ratios of
gases can be used instead of mole ratios. (Hence simplifying many calculations.)
Corrosion and Redox Reactions
77. Corrosion is a process in which water and gases of the atmosphere react with many metals and
cause them to oxidise. Corrosion processes are oxidation and reduction (redox) processes.
78. The corrosion of metals always involves the oxidation of the metal. Steel corrodes to form
rust (iron oxide), aluminium corrodes to form a protective layer of aluminium oxide. The ease
of corrosion depends on the metal, the corrosion product and the environment surrounding the
metal (e.g. presence of water, oxygen and temperature.)
79. Oxidation and reduction always occur simultaneously. While one substance is oxidised in a
reaction, another substance is simultaneously being reduced.
80. The addition of oxygen to a substance is an oxidation reaction. Eg. The combustion of fuels
always requires oxygen and in the formation of metal oxides the metal is oxidised.
C + O2  CO2 (C is oxidised)
2Mg + O2  2 MgO (Mg is oxidised)
81. In redox reactions electrons are transferred. The substance that loses electrons is oxidised, the
substance that gains electrons is reduced. (OIL RIG)
82. A substance that causes another to be oxidised is an oxidant or oxidising agent. An oxidant is
an electron donor. (It is itself reduced)
83. A substance that causes another to be reduced is a reductant or reducing agent. A reductant is
an electron acceptor. (It is itself oxidised)
84. Redox reactions can be represented by two half equations.
Cu 2+(aq) + Fe(s)  Cu(s) + Fe 2+(aq) (redox reaction)
Cu 2+(aq) + 2 e-  Cu(s) (reduction)
Fe(s)  Fe 2+(aq) + 2 e- (oxidation)
Electrochemical Cells
85. In a galvanic or electrochemical cell a chemical reaction is used to produce electricity.
86. In a galvanic cell electrons flow from the anode to the cathode. Oxidation occurs at the anode
(production of electrons), which is negative. Reduction occurs at the cathode (absorption of
electrons), which is positive. (AN OIL RIG CAT)
87. A half-cell must contain an electrode (either cathode or anode) and an electrolyte (containing
ions that are free to move).
88. An electrochemical cell contains two half cells which cannot mix, a wire joining the
electrodes of the half cells (which allows flow of electrons) and a salt bridge which acts as an
electrolytic conductor and allows ion movement to balance the positive and negative charges
in each cell. The oxidant is separated from the reductant by two half cells.
89. The electrochemical series is an ordered list of chemical reactivity. It shows the relative
reactivities of a number of metals and their corresponding cations.
90. The electrochemical series can be used to make predictions about redox reactions. Any
oxidant on the left will react with any reductant on the right which is below it.
91. The most reactive metals are those on the far left of the periodic table in groups I and II. The
transition metals and those in group IV (Sn, Pb) are less readily oxidised.
Corrosion of Iron
92. Corrosion of iron is generally accelerated by water, impurities in the water (salt, acidic
pollutants) and impurities in the iron. Corrosion is generally inhibited by alloying the iron,
using a protective coating and changing the surface chemistry.
93. Corrosion of iron: Fe(s)  Fe 2+(aq) + 2 eO2(aq) +
2 H2O(l) + 4e-  4 OH-(aq)
Overall: 2 Fe(s) + O2(aq)
Further reaction: Fe 2+(aq)
Fe(OH)2(s) + O2(aq)
+
+
+
2 H2O(l)  2 Fe 2+(aq) +
4 OH-(aq)
2 OH-(aq)  Fe(OH)2(s) (green precipitate)
2 H2O(l)  4 Fe(OH)3(s) (red-brown precipitate, rust)
Protection against Corrosion
94. Corrosion can be prevented passively by covering the surface with paint, plastic, another
metal (e.g.chromium, tin).
95. Corrosion can be prevented by sacrificial protection. A more active metal is attached to that
which needs to be protected. The more active metal is corroded in preference to the protected
metal. E.g. galvanised iron, zinc corrodes before iron; aluminium or zinc blocks on oil
platform legs, ship hulls corrode in preference to the iron. These sacrificial anodes are
consumed and so must be replaced periodically.
96. Cathodic protection of a steel structure can be achieved by using a low-volatage (DC) power
supply. The steel being protected has a negative charge imposed on it that inhibits oxidation /
corrosion of the steel as oxidation would produce electrons. This is used in steel pipelines for
natural gas.
97. Oxidation numbers are assigned to substances to help to determine whether a redox reaction
has occurred. They are assigned using a series of rules and have no physical significance. A
substance that undergoes oxidation experiences an increase in oxidation number, a substance
that undergoes reduction experiences a decrease in oxidation number.
98. The chemical industry is essential in today's society. It includes areas such as paints, plastics,
rubbers, fertilisers, pesticides, industrial gases, soaps, detergents and explosives. Chemicals
have modified most of the everyday things that we eat, use, make or wear.
99. The treatment of waste is of great importance. The treatment depends on such areas as the
availability of the technology to treat the waste, the relative cost of the treatment process, the
affect of the chemicals on the environment, government legislation and the scarcity of
materials.
100.
Wastes can be classed as intractable (almost impossible to make safe) and tractable.
Intractable wastes can be incinerated (toxic chemicals e.g. heavy metal compounds,
organochlorine compounds - DDT) or stored (radioactive materials, PCBs ). Tractable wastes
are either biodegradable or non-biodegradable. Biodegradable wastes can be recycled (paper,
oil) or dumped (paper, oil, garbage). Non-biodegradable wastes can be reused (glass),
recycled (aluminium, iron, glass), treated (nitrogen oxides, carbon monoxide) or dumped
(iron, glass, nitrogen oxides, carbon momoxide).