1. Definitions

2. Energy level
is the fixed energy value that an electron in an atom may have.
Exam Q (Hons)
‘08/Q10c
‘07/Q4

3. Ground State lowest energy state ( in 1s orbital)
Excited state = higher energy state
Exam Q (Hons)
08/Q10c

4. An orbital
is a region in space within which there is a high probability of finding an electron.
Exam Q (Hons)
‘06/Q5

5. An element
is a substance that cannot be split up into simpler substances by chemical means.

6. A triad
is a group of three elements with similar chemical properties in which the atomic weight of the middle element is approximately equal to the average of the other two.
(Dobereiner)

7. Newlands’ Octaves
are groups of elements arranged in order of increasing atomic weight, in which the first and the eighth element of each group have similar properties.

8. Mendeleev’s Periodic Law
When elements are arranged in order of increasing atomic weight (relative atomic mass), the properties of the elements vary periodically.

9. The atomic number(Z)
is the number of protons in the nucleus of that atom.

10. Periodic Table
is an arrangement of elements in order of increasing atomic number.

11. Elements are arranged
in order of increasing atomic number, the properties of the elements vary periodically.

12. Mass number (A)
is the
sum of the number of protons and neutrons in the nucleus of an atom of that element.

13. Isotopes
Exam Q (Hons)
’06/Q10a
are atoms of the same element ( i.e. they have the same atomic number) that have different mass numbers due to the different number of neutrons in the nucleus.

14. Relative Atomic Mass
is the average of the mass numbers of the isotopes of the element
as they occur naturally
taking their abundances into account
relative to 1/12th mass of carbon 12 atom
(expressed on a scale in which the atoms of carbon 12 isotope have a mass of exactly 12 units).
Exam Q (Hons)
’06/Q10a

15. Mass Spec. V I A S D Vaporisation how? why? Ionisation Acceleration
Separation
Victor
A+
Detection

16. Aufbau Principle
that when building up the electronic configuration of an atom in its ground state, the electrons occupy the lowest available energy level.

17. Hund’s Rule of Maximum Multiplicity
states that
when two or more orbitals of equal energy are available,
the electrons occupy them singly first before filling them in pairs.

18. Pauli Exclusion Principle
that no more than two electrons may occupy an orbital and they must have opposite spins.

19. Compound
is a substance that is made up of two or more different elements combined together chemically.

20. Octet Rule
that when bonding occurs, atoms tend to reach an electron arrangement with eight electrons in the outermost shell.

21. An Ion
is a charged atom or group of atoms.

22. An Ionic bond
is the force of attraction between oppositely charged ions in a compound.

23. A transition metal
is one that forms at least one ion with a partially filled d sublevel.

24. Molecule
is a group of atoms joined together. It is the smallest particle of an element or compound that can exist independently.

25. Valency of an element
is defined as the number of atoms of hydrogen or any other monovalent element with which each atom of the element combines.

26. Electronegativity
is a measure of the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond.
Exam Q (Hons)
’06/Q5

27. Electronegativity
difference > 1.7 indicates ionic bonding in a compound.
An electronegativity difference ≤ 1.7 indicates covalent bonding in a compound.

28. The value of electronegativity
decrease down the groups in the Periodic Table for two reasons:
increasing atomic radius
screening effect of inner electrons

29. The values of electronegativity
increase across the periods in the Periodic Table for two reasons:
increasing nuclear charge
decreasing atomic radius
F= most electronegative element.
Halogens –decrease in reducing power down the group due to drop in electroneg. values.

30. Intermolecular Forces
attractive (repulsive) forces
between molecules
Intramolecular forces are
attractive (repulsive) forces within a molecule

31. Vans der Waals Forces
are weak attractive forces between molecules resulting from the formation of temporary dipoles.

32. Dipole-dipole
Dipole – dipole forces are forces of attraction between the negative pole of one molecule and the positive pole of another.

33. Hydrogen bonds
are particular types of dipole-dipole attractions between molecules in which hydrogen atoms are bonded to nitrogen, oxygen or fluorine.
The hydrogen atom carries a partial positive charge and is attracted to the electronegative atom in another molecule. Thus, H acts as a bridge between two electronegative atoms.

34. The Law of Conservation of Mass
the total mass of the products of a chemical reaction is the same as the total mass of the reactants.

35. The Law of Conservation of Matter
that in any chemical reaction, matter is neither created nor destroyed but merely changes from one form into another.

36. Tests for Anions Chloride Sulfate/sulfite carbonate/hydrogen carbonate
nitrate
phosphate
(NB know confirmatory test too!)

37. Chloride Add AgNO3 Get white ppt Confirm = ppt dissolves in
dilute ammonia
Equation needed

38. Sulfate/sulfite Add BaCl2 Get white ppt Distinguish add
dil HCl to white ppt
ppt remains = sulfate
ppt dissolves = sulfite
Equation needed !!

39. CO32- /HCO3- Add dil. HCl (or any acid) Get
bubbles of CO2 (limewater milky)
Distinguish
add
MgSO4 to fresh solution
white ppt. immediately = carbonate
white ppt on heating = hydrogen carbonate
Equation needed !!

40. Nitrate Brown Ring Test Add fresh FeSO4
At slant add conc. H2SO4 drop wise
Get
brown ring at junction of 2 layers
No equation needed

41. Phosphate Add ammonium molybdate
5 drops of conc. nitric acid (warm the solution)
Get yellow ppt
No equation needed
Confirm: Goes colourless when add dilute NH3

42. The atomic radius of an atom
is defined as half the distance
between the nuclei of two atoms
of the same element that are
joined together by a single covalent bond.
Exam Q (Hons)
07/Q4

43. The values of atomic radius
increase down any one group in the Periodic Table for two reasons:
extra shell
screening effect of inner electrons

44. The values of atomic radius
decrease from left to right across a Periodic Table for two reasons:
increasing nuclear charge
no increase in screening effect

45. The first ionisation energy of an atom
is the minimum energy required
to completely remove
the most loosely bound electron
from one mole of neutral gaseous atom
in the ground state.******
2004 =9 marks (2.25%)
2002 = 8marks(2%)

46. The values of ionisation energy
decrease down the groups in the Periodic Table for two reasons:
increasing atomic radius
screening effect of inner electrons

47. The values of ionisation energy
increase across the Periodic Table for two reasons:
increasing nuclear charge
decreasing atomic radius

48. More on ionisation energy
First Ionisation Energy
M – e- M+
Second Ionisation energy
M+ – e- M2+
Major jump in I.E. values – significance

49. The value of electronegativity
decrease down the groups in the Periodic Table for two reasons:
increasing atomic radius
screening effect of inner electrons

50. The values of electronegativity
increase across the periods in the Periodic Table for two reasons:
increasing nuclear charge
decreasing atomic radius
F= most electronegative element.
Halogens –decrease in reducing power down the group due to drop in electroneg. values.

51. A gas
is a substance that has no well-defined boundaries but diffuses rapidly to fill any container in which it is placed.

52. Radioactivity is the spontaneous breaking up of unstable nuclei
with the emission of one or more types of radiation.

53. Alpha particles loss of He nucleus (2p + 2n) mass number down by 4
atomic number down by 2
element changes to element two places back

54. Beta particle neutron changes to proton and electron electron emitted
mass number stays same
atomic number drops by one
element changes into element one place back

55. Gamma radiation no new atoms formed (no transmutation)
only energy lost

56. Half Life of an element is the time taken for half the nuclei
in any given sample to decay.

57. Mole
is the amount of a substance which contains 6 X 1023 particles of that substance
(avogadro’s number or constant =L)

58. a few numbers Kelvin = Celsius + 273 standard temp = 273 K
standard pressure = 1X105 Pa
(100kPa)
m3 = litres X10 -3
m3 = cm3 X10 -6
(1 litre = 1000cm3)

59. Mole contains 6 X 1023 particles has mass equal to Ar or Mr in grams
occupies 22.4 litres at s.t.p (if gas)

60. Boyle’s Law states that: at constant temperature,
the volume of a fixed mass of gas is
inversely proportional to its pressure.

61. Charles’ Law states that: at constant pressure,
the volume of a fixed mass of a gas
is directly proportional to its temperature measured on the Kelvin scale.

62. General Gas Law P1 X V1 = P2 X V2 T1 T2 Temp in Kelvin
Units for volume same each side
Units for pressure same each side

63. Gay Lussac’s law of Combining Volumes
the volumes of the reacting gases and the volumes of any gaseous products are
in the ratio of small whole numbers
provided the volumes are measured at the same temp and pressure

64. Avogadro’s Law states that equal volumes of gases contain
equal numbers of molecules under the same conditions of temp. and pressure
Exam Q (Hons)
‘07/Q10b

65. Molar Volume At s.t.p one mole of any gas occupies 22.4 litres
Remember to watch out for r.t.p in questions
room temp. and press = as given in Q
(often 24 litres)

66. Ideal Gas is one which perfectly obeys
Exam Q (Hons)
’06/Q11a
is one which perfectly obeys
all the gas laws and all the assumptions* of the kinetic theory of gases
under all conditions of temperature and pressure.
(Know the assumptions)

67. Real v. ideal gas
Real gases differ from ideal gases at high pressure and low temp. because
there are forces of attraction/repulsion between the molecules*
the volume of the molecules is not negligible compared to the distances between them
(*know examples of real gases and the forces involved)

68. Empirical Formula
gives the simplest whole number ratio of the numbers of the different atoms present in the molecule.
(divide by Ar and get ratio)
(molecular formula is a simple multiple of the empirical formula)

69. Acids / Bases Arrhenius Acid + Base Bronsted Lowry Acid + Base
Neutralisation
Conjugate Acid/ conjugate base
conjugate pair

70. Arrhenius Acid and Base
is a substance that dissociates in water to produce H+ ions.
Arrhenius Base
is a substance that dissociates in water to produce OH- ions.

71. Bronsted Lowry Acid /Base
is a proton (H+) donor
Bronsted Lowry Base
is a proton (H+) acceptor
Exam Q (Hons)
’07/Q7

72. Neutralisation is the reaction between an acid and a base forming
a salt and water
(acid + base -> salt + water)
SALT = is formed when the H of an acid is replaced by a metal

73. Conjugate Acid / Conjugate Base
is formed when a base accepts a proton
Conjugate Base
is formed when an acid donates a proton.

74. Conjugate Pair an acid and a base that differ by a proton
Exam Q (Hons)
‘07/Q7

75. Primary Standard is a substance of high Mr which can be obtained
in a pure stable soluble solid form
so that it can be weighed out and
dissolved in water to give
a solution of accurately known concentration.
(Know why high Mr matters)

76. Titration is a laboratory procedure where a
a measured volume of one solution is added to
a known volume of another solution until the reaction is complete.
(concentration of one solution known accurately at start)
(indicator used to show by colour change when reaction is complete)

77. Oxidation Reduction
Revision

78. Definitions Oxidation is addition of loss of increase in Exam Q (Hons)
‘08/Q10(b)
’06/Q10(b)
oxygen
electrons
oxidation number

79. Reduction is loss of oxygen gain of electrons
decrease in oxidation number

80. More…
An oxidising agent causes oxidation
and is itself reduced.
A reducing agent causes reduction
and is itself oxidised.
What is a redox reaction?

81. What is oxidised and reduced in each of the following?
Br2 + 2Fe 2+ → 2Br– + 2Fe 3+
Cu 2+ + Zn  Cu + Zn 2+
2Na + Cl2  2NaCl

82. Oxidation Number Rules
The oxidation number of
an Element is
group One elements is +1
group Two elements is +2 in
compounds

83. The oxidation number of
an ion is equal to the
charge on the ion
halogens is
-1 (in binary compounds)
(except ……????)

84. The oxidation number of H in a compound
is +1
except
in metal hydrides when it is -1

85. The oxidation number of O in a compound is-2
except (x2)
in peroxides when it is -1 (H2O2)
in OF2 when it is +2 (why?)

86. Oxidation numbers add up to zero in a compound
add up to the charge of a complex ion

87. What is the oxidation number of each element in :-
MnO4¯ I2
KBrO3
Na2S2O3
H2O2
NaClO

88. KMnO4 (end point = first permanent pink) oxidising agent purple
read top of meniscus
is reduced from
Mn (VII)  Mn (II) in presence of H+
purple  colourless
own indicator
(end point = first permanent pink)

89. KMnO4 not primary standard (x2)
get brown Mn (IV) if H+ absent
(which acid MUST be used – why x2)
not primary standard (x2)
standardised by titrating against standard solution of acidified Fe 2+

90. H2SO4
added during KMnO4 titrations to provide H+ and ensure the complete reduction of Mn (VII)  Mn (II)
and prevent formation of Mn (IV)
(brown)
added during prep. of Fe (II) solutions to
prevent oxidation of Fe 2+ to Fe 3+ by oxygen in the air ( why does this matter?)

91. Na2S2O3 S2O3 2- ion reducing agent used in photography
not primary standard – why ?
standardised by titrating against I 2
starch indicator – when added and why
colour change at end point ?

92. Iodine I2 Oxidising agent NOT a primary standard (X2)
Produced when MnO4- oxidises I- to I2
(known concentration) (in excess)
Starch indicator – when added? why then?
Colour change at end point
Blue/black to colourless

93. remember dilution factor
Bleach
sodium hypochlorite Na+ClO-
bleach diluted x10 with distilled water
not de-ionised water (why? )
ClO- oxidises I- to I2
I2 v. thiosulfate
starch indicator as before NB
remember dilution factor
in calculations

94. Definitions Rate of reactions

95. Rates of Reactions The rate of reaction is the change in concentration
Exam Q (Hons)
2003 Q7
2004 Q8
2007/Q9
2011 /Q5
The rate of reaction is
the change in concentration
per unit time of any one reactant or product.

96. Factors affecting rate
nature of reactants
particle size
concentration
temperature
catalysts

97. Equations to know Write each equation then check
Decomposition of hydrogen peroxide using manganese dioxide as catalyst
MnO2
2H2O2 2H2O + O2
Sodium thiosulfate and hydrochloric acid
Na2S2O7 + 2HCl  S + 2NaCl + SO2 +H2O

98. Rate Graphs Concentration v. ( 1 /Time ) or Temp v. ( 1 /Time )
( 1 /Time )used as Rate and Time inversely related
(shorter time means faster rate)
be careful with units of 1/time

99. Catalyst is a substance that alters the rate of reaction
Exam Q (Hons)
2003 Q7
‘07/Q9
is a substance that alters the rate of reaction
but is not consumed in the reaction.

100. Homogeneous catalysis
occurs when the reactants and the catalyst are in the same phase.
example =?
liquids
KI catalyses 2H2O2 2H2O + O2 (iodine snake)
And any enzyme

101. Heterogeneous catalysis
Exam Q (Hons)
‘07/Q4
occurs when the reactants and the catalyst are in different phases.
(NB must be phases not states)
example = ?
Al2O3 (solid)catalyses
ethanol (gas) ethene
Methanol  methanal using platinum

102. Autocatalysis
occurs when one of the products of the reaction catalyses the reaction.
Example = ?
Mn2+ ions in KMnO4 titrations
(purple changes to colourless more quickly as titration proceeds)

103. Mechanism of Catalysis
Intermediate Formation theory
Surface Adsorption theory
Know details of each and evidence of intermediate formation theory

104. Enzymes Are biological catalysts made of protein
Examples of homogeneous catalysis
Need to know 2 examples
Amylase catalyses conversion of starch to maltose
Catalase catalyses conversion of hydrogen peroxide to hydrogen and water

105. Learning Check Do I know Definition for Rate of reaction Catalyst
Homogeneous catalysis
Heterogeneous catalysis
Auto catalysis
Two mechanisms of catalysis
Press enter to continue

106. Catalytic converter Catalysts = ?
Pt + Pd + Rh on honeycomb surface (ceramic)
Gases in CO NO
NO2
hydrocarbons
Gases out
CO2 and N2 and H20

107. Learning Check Do I know 3 metals in Catalytic converter
4 wastes in exhaust fumes
Problem of each
What each is converted to
What poisons catalytic converter
Type of catalysis occuring in catalytic converter
Press enter to continue

108. Collision Theory
for a reaction to occur the reacting particles must collide with each other
a collision only results in a product being formed if a certain minimum energy is exceeded (called activation energy)

109. Effective Collision Is one in which a reaction occurs
Exam Q (Hons)
2009 Q9
Is one in which a reaction occurs
The activation energy has been reached or exceeded.

110. Activation Energy
is the minimum energy which colliding particles must have for a reaction to occur
(minimum energy required for effective collisions between particles)
Exam Q (Hons)
2006/Q7
2009/Q9

111. Activation Energy 2
Catalysts lower the activation energy of a reaction
Compare E act
Without catalyst
with catalyst

112. Energy Profile Diagram
Sketch an energy profile diagram for an endothermic reaction.
Press enter when ready and
It should look like this

113. Endothermic Energy Profile Diagram
Products
Activation Energy
Energy In
Reactants
Note – axes should be labelled Time (x) and energy (y)
Curve should be smooth !

114. Energy Profile Diagram
Sketch an energy profile diagram for an exothermic reaction.
Press enter when ready and
It should look like this

115. Exothermic Energy Profile Diagram
Activation Energy
Reactants
Energy Out
Products
Note – axes should be labelled Time (x) and energy (y)
Curve should be smooth !

116. Learning Check Do I know Definition for
Effective collision
Activation energy
Can I draw energy profile diagram for
Exothermic reaction
Endothermic reaction
Either of above with catalyst
The End

117. Chemical Equilibrium
is a state of dynamic balance where the rate of the forward reaction equals the rate of the reverse reaction.

118. Le Chatelier’s Principle
If a stress is applied to a system at equilibrium
the system readjusts to oppose the stress applied
reactions at equilibrium // oppose the applied stress(es)*
Exam Q (Hons)
’06/Q11b

119. Le Chatelier’s Principle and Gases
Le Chatelier’s Principle predicts that
in an all-gaseous reaction
an increase in pressure
will favour the reaction which takes place with a reduction in volume
( towards the side with the smaller number of molecules)

120. Equilibrium Constant Kc [ ] means concentration in moles per litre
[ C] c x [D]d
Kc = for aA + bB cC + dD
[A]a x [B]b
(product of products conc.
over
product of reactants conc.)

121. Le Chatelier and Industry
Ammonia and Haber Process
predict max yield at high press. /low temp
reality = 200 atm and 500o C
Sulfuric Acid and Contact Process
reality = one atm and 450oC

122. Kc large Kc => equilibrium far to right (lots of product produced)
small Kc => equilibrium far to left
(v. little product formed)
must quote temp.
units – depend on reaction
tells us how far not how fast a reaction occurs

123. pH pH = -log [H+] [ ] = moles per litre pH < 7 acid pH = 7 neutral
pH > 7 base
[ ] =
moles per litre

124. Kw Kw = [H+].[OH-] Also remember Kw = 1x10-14 ( at 25oC) ([H+] = √Kw)
so [H+]= 1x10-7
and pH =7
Exam Q (Hons)
08/Q8

125. Strong / Weak acid A strong acid is a good proton donor
or (is fully dissociated into ions in dilute aqueous soln.
[H+] = [acid] HCl
[H+] = 2x[acid] H2SO4 etc
A weak acid is a poor proton donor
or (slightly dissociated into ions in dil. aq. soln.)
[H+] = √Ka x Macid
Exam Q (Hons)
‘07/Q7

126. Strong / Weak base
A strong base is a good proton acceptor or one which is fully dissociated into ions in dilute aqueous solution
[OH-] = [base] NaOH
[OH-] = 2x[base] Ca(OH)2 etc
A weak base is a poor proton acceptor or one which is slightly dissociated in dli. aq. soln.
[OH-] = √Kb x Mbase

127. Indicator
An acid base indicator is a substance that changes colour according to the pH of the solution it is in.
(equilibrium HIn ↔ H+ + In-)

128. Methyl orange in acid (lower pH ) red in base ( higher pH) yellow

129. Phenolphthalein in acid (lower pH ) colourless in base ( higher pH)
pink
range
pH 8-11

130. Litmus in acid (lower pH ) red in base ( higher pH) blue range pH 5-8
(Not as reliable as others for accurate work)

131. Which indicator ? strong acid/strong base =
methyl orange / phenolphthalein /litmus( see above)
strong acid / weak base =
methyl orange
weak acid /strong base =
phenolphthalein
weak acid / weak base =
none (why?)

132. Hard Water is water that will not easily form a lather with soap
due to the presence of Ca 2+ or Mg 2+ ions in solution.
Exam Q (Hons)
’06/Q8

133. Temporary Hardness can be removed by boiling the water
due to Ca(HCO3)2
becomes CaCO3 on heating
leads to blocked pipes etc

134. Permanent Hardness is not removed by boiling the water
caused by CaSO4 or MgSO4

135. Methods of removing hardness
boiling (only works for temp. hardness)
distillation
washing soda
ion exchange

136. Calculations Total hardness = calcium hardness + magnesium hardness
but
Do calculations as if all hardness caused by CaCO3
expressed in p.p.m of CaCO3
p.p.m. = mg/litre

137. Water Treatment screening flocculation sedimentation filtration
chlorination
fluoridation
pH adjustment

138. B.O.D Biochemical Oxygen Demand is the amount of dissolved oxygen
consumed by biological action
when a sample of water is kept
at 20oC
in the dark
for five days.
(know reason for each of 3 conditions)

139. Eutrophication
is the enrichment of water with nutrients which leads to the excessive growth of algae.
Nutrients – phosphates/nitrates
Algal bloom / oxygen depletion

140. Sewage Treatment Primary Treatment physical Secondary Treatment
biological
Tertiary Treatment
chemical

141. Water Analysis Atomic Absorption Spectrometry used to
detect heavy metals like Cd, Hg, Pb
pH meter
colorimetry
(Hach test Chlorine in pool water)

142. Electrolysis
is the use of electricity to bring about a chemical reaction.
KI/ Acidified water/ Na2SO4/CuSO4 / ions

143. Electrolyte
is a substance that conducts electricity as a result of the presence of ions.

144. Electroplating
is the process where electrolysis is used to put a layer of one metal on the surface of another.

145. Electrochemical Series
is a list of the elements in order of their standard electrode potentials.

146. Organic Chemistry is the study of compounds of carbon…
(except some simple compounds
like CO2, CO and carbonates)

147. Hydrocarbon is a compound that contains only carbon and hydrogen
includes
alkanes, alkenes, alkynes
excludes
alcohols, aldehydes, ketones, carboxylic acids, esters

148. Saturated compound is one with only carbon – carbon single bonds
alkanes

149. Unsaturated compound is one which contains at least
one carbon – carbon double or triple bond
and undergoes addition reactions
alkenes / alkynes
test for unsaturation
decolourise bromine solution
Exam Q (Hons)
08/Q9

150. Homologous Series
is a series of chemical compounds of uniform chemical type
showing gradations in physical properties
having a general formula for its members
each member has similar method of prep. and
each member differs by (CH2) from previous member

151. Structural isomers are compounds with the same molecular formula but
different structural formulas.
e.g. butane and methyl propane are both C4H10
need to know isomers up to C5H12

152. Aliphatic
an aliphatic compound is an organic compound that consists of
straight (open) chains of carbon atoms
and closed chain compounds with similar properties.

153. Aromatic
An aromatic compound is an organic compound that contains a benzene ring structure in their molecules.
(benzene – delocalised double bond)
(disc. by Michael Faraday )
(structure by Kekule)

154. Octane Number
of a fuel is a measure of the tendency of the fuel to resist knocking.
(Best fuels = high octane number
= 100 = 2,2,4 tri methyl pentane ) (Short chains, more branched chains, ring structures)
(Worst fuels = low octane number
=0 = heptane)
Exam Q (Hons)
‘08/Q6
’06/Q6

155. Ways to increase octane number
isomerisation
catalytic cracking
dehydro-cyclis-ation (re-forming)
add oxygenates

156. Isomerisation
changing straight chain alkanes into branched chain alkanes

157. Catalytic cracking is the breaking down of
long chain hydro- carbon molecules into short chain molecules by heat and catalysts
(for which there is a greater demand)
Exam Q (Hons)
‘07/Q6

158. Dehydrocyclisation ( Re-forming)
involves the use of catalysts to form ring structures
straight chain alkanes changed to cycloalkanes
cycloalkanes changed to aromatic compounds
petrol contains benzene = carcinogen
health concerns

159. Adding Oxygenates addition of to petrol to increase the octane number.
methanol
ethanol
MTBE
to petrol to increase the octane number.
(Methyl Tertiary Butyl Ether or
2 methoxy 2 methyl propane)

160. Exothermic reaction is one which produces heat.
∆H is minus ( giving away)

161. Endothermic reaction is one which takes in heat.
∆H is positive (add in)
ammonium nitrate dissolving in water

162. Heat of Reaction
is the heat change involved when the numbers of moles of reactants indicated in the balanced equation for the reaction react completely.

163. Heat of Combustion is the heat change involved when
one mole of a substance is
completely burned in
excess oxygen

164. Kilogram Calorific Value
of a fuel is the heat energy produced when 1 kg of a fuel is completely burned in oxygen.

165. Bond Energy
is the energy required to break one mole of covalent bonds and to separate the neutral atoms completely from each other.

166. Heat of Neutralisation
is the heat change involved when
one mole of H+ ions from an acid
reacts with
one mole of OH- from a base
forming one mole of H2O

167. Heat liberated Heat liberated = M x C x Rise in temp.
Kg kelvin
M=Mass of solution in Kg
c=specific heat capacity
rise in temp in Kelvin

168. Heat of formation of a compound is the heat change involved
when one mole of a compound
in its standard state
is formed from its elements
in their standard states.

169. Hess’s Law
states that if a chemical reaction takes place in a number of stages,
the sum of the heat changes in the separate stages is equal to the heat change if the reaction is carried out in one stage.
(overall heat change is independent of the pathway)

170. Law of Conservation of Energy
states that
energy cannot be created or destroyed
but
can be changed from one form of energy to another.

171. Functional Group
is an atom or group of atoms which is responsible for the characteristic properties of a series of organic compounds.

172. Substitution Reaction
is a chemical reaction in which an atom or group of atoms in a molecule is replaced by another atom or group of atoms
mechanism = free radical substitution
initiation (homolytic fission)
propagation
termination

173. Addition Reaction
is a chemical reaction is which two substances react together forming a single substance.
Mechanism = Ionic addition
Approach/ polarisation / heterolytic fission /carbonium ion / product formation
only happens to unsaturated compounds

174. Polymers
are long chain molecules made by joining together many small molecules called monomers.

175. Elimination reaction
is one in which a small molecule is removed from a larger molecule to leave a double bond in the larger molecule.

176. Organic Synthesis
is the process of making organic compounds from simpler starting materials.

177. Chromatography is a separation technique in which a mobile phase
carrying a mixture
moves in contact with
a selectively adsorbent stationary phase.

178. Instrumentation Mass Spec. AAS GC HPLC IR spec. UV spec.
X ray crystallography (option 2)

179. Mass Spec. Positively charged ions are separated
Principles
Processes
Positively charged ions are separated
according to different relative masses
when moving through magnetic field
Vapourisation
Ionisation
Acceleration
Separation
Detection
Used to
Analyse blood of race horses for drugs
Identify substances
Victor
A+

180. Atomic Absorption Spectrometry
Principles
Processes
Ground state atoms of an element absorb light characteristic of that element.
Absorption is directly proportional to concentration.
(higher absorbance means higher concentration of THAT ELEMENT present)
Dissolve
Atomise
Absorb
Measure
Detection
Used to
analyse water samples for heavy metals Cd Hg Pb

181. Gas Chromatography GC
Principles
Processes
Different components have different tendencies to dissolve in a non-volatile liquid, which is coated on fine particles of a solid in a the GC column
Injection …
Transport …
Separation …
Detection ...
Used with MS in drug testing
also blood alcohol levels
mobile phase ?
stationary phase? ?

182. HPLC
Principles
Processes
High Performance Liquid Chromatography
Different components of a mixture have different tendencies to adsorb onto fine particles of solid in HPLC column
Injection …
Transport …
Separation …
Detection …
Used to separate less volatile mixtures e.g. growth promoters in meat.
mobile phase ?
stationary phase? ?

183. IR Infra red spectrometry
Principles
Processes
Infra red spectrometry
Molecules of a substance absorb infra –red of different frequencies.
(different number/ type bonds)
The combination of frequencies absorbed is unique to the molecules of each substance
Prepare …
Transmit IR
Absorption…
Detection…
Spectrum obtained
Used to identify functional groups
and identify drugs

184. UV Ultra violet spectrometry Molecules absorb UV radiation
Principles
Processes
Ultra violet spectrometry
Molecules absorb UV radiation
Electrons promoted from ground state to higher energy states.
Absorption is directly proportional to concentration.
Prepare
Transmit UV through
Blank (o%abs)
Sample
(known + unknown)
Spectrum obtained
Quantatative
used to find amount of org. subs. e.g. drugs

185. X ray crystallography
Processes
Principles
Wavelengths of Xrays are comparable to distance between atoms in a crystal
Xrays are scattered when they hit a crystal surface
Pattern detected is analysed and structure worked out
Prepare ….
Transmit : x-ray detected on film
Pattern analysed and structure worked out
Used to determine structure of macro-molecules e.g. DNA