Chemistry AS - Spec Check

Spec Check.

1 of 144

1.1.1

1.1.1 Atoms

 atomic structure;

 relative masses.

2 of 144

Atomic Structure - 1

(a) describe protons, neutrons and electrons in

terms of relative charge and relative mass;

3 of 144

Atomic Structure - 2

(b) describe the distribution of mass and charge

within an atom;

4 of 144

Atomic Structure - 3

(c) describe the contribution of protons and

neutrons to the nucleus of an atom, in terms

of atomic (proton) number and mass

(nucleon) number;

5 of 144

Atomic Structure - 4

(d) deduce the numbers of protons, neutrons and

electrons in:

(i) an atom given its atomic and mass

number,

(ii) an ion given its atomic number, mass

number and ionic charge;

6 of 144

Atomic Structure - 5

(e) explain the term isotopes as atoms of an

element with different numbers of neutrons

and different masses;

7 of 144

Relative Masses - 1

(f) state that 12C is used as the standard

measurement of relative masses;

8 of 144

Relative Masses - 2

(g) define the terms relative isotopic mass and

relative atomic mass, based on the 12C scale;

9 of 144

Relative Masses - 3

(h) calculate the relative atomic mass of an

element given the relative abundances of its

isotopes;

10 of 144

Relative Masses - 4

(i) use the terms relative molecular mass and

relative formula mass and calculate values

from relative atomic masses.

11 of 144

1.1.2

1.1.2 Moles and Equations

 the mole;

 reacting masses and equations

12 of 144

The mole - 1

(a) explain the terms:

(i) amount of substance,

(ii) mole (symbol ‘mol’), as the unit for

amount of substance,

(iii) the Avogadro constant, NA, as the

number of particles per mole (6.02 ×

1023 mol–1);

13 of 144

The mole - 2

(b) define and use the term molar mass (units g

mol–1) as the mass per mole of a substance;

14 of 144

Empirical and molecular formulae - 1

(c) explain the terms:

(i) empirical formula as the simplest whole

number ratio of atoms of each element

present in a compound,

(ii) molecular formula as the actual number

of atoms of each element in a molecule;

15 of 144

Empirical and molecular formulae - 2

(d) calculate empirical and molecular formulae,

using composition by mass and percentage

compositions;

16 of 144

Chemical equations - 1

(e) construct balanced chemical equations for

reactions studied and for unfamiliar reactions

given reactants and products;

17 of 144

Calculation of reacting masses, mole concentration

(f) carry out calculations, using amount of

substance in mol, involving:

(i) mass,

(ii) gas volume,

(iii) solution volume and concentration;

18 of 144

Calculation of reacting masses, mole concentration

(g) deduce stoichiometric relationships from

calculations;

19 of 144

Calculation of reacting masses, mole concentration

(h) use the terms concentrated and dilute as

qualitative descriptions for the concentration

of a solution.

20 of 144

1.1.3

1.1.3 Acids

 acids and bases;

 salts

21 of 144

Acids and bases - 1

(a) explain that an acid releases H+ ions in

aqueous solution;

22 of 144

Acids and bases - 2

(b) state the formulae of the common acids:

hydrochloric, sulfuric and nitric acids;

23 of 144

Acids and bases - 3

(c) state that common bases are metal oxides,

metal hydroxides and ammonia

24 of 144

Acids and bases - 4

(d) state that an alkali is a soluble base that

releases OH– ions in aqueous solution;

25 of 144

Acids and bases - 5

(e) state the formulae of the common alkalis:

sodium hydroxide, potassium hydroxide and

aqueous ammonia

26 of 144

Salts - 1

(f) explain that a salt is produced when the H+

ion of an acid is replaced by a metal ion or

NH4+;

27 of 144

Salts - 2

(g) describe the reactions of an acid with

carbonates, bases and alkalis, to form a salt;

28 of 144

Salts - 3

(h) explain that a base readily accepts H+ ions

from an acid: eg OH– forming H2O; NH3

forming NH4+;

29 of 144

Salts - 4

(i) explain the terms anhydrous, hydrated and

water of crystallisation;

30 of 144

Salts - 5

(j) calculate the formula of a hydrated salt from

given percentage composition, mass

composition or experimental data;

31 of 144

Salts - 6

(k) perform acid–base titrations, and carry out

structured titrations.

32 of 144

1.1.4

1.1.4 Redox

 oxidation number;

 redox reactions

33 of 144

Oxidation Number - 1

(a) apply rules for assigning oxidation number to

atoms in elements, compounds and ions;

34 of 144

Oxidation Number - 2

(b) describe the terms oxidation and reduction in

terms of:

(i) electron transfer,

(ii) changes in oxidation number;

35 of 144

Oxidation Number - 3

(c) use a Roman numeral to indicate the

magnitude of the oxidation state of an

element, when a name may be ambiguous,

eg nitrate(III) and nitrate(V);

36 of 144

Oxidation Number - 4

(d) write formulae using oxidation numbers;

37 of 144

Redox reactions - 1

(e) explain that:

(i) metals generally form ions by losing

electrons with an increase in oxidation

number to form positive ions,

(ii) non-metals generally react by gaining

electrons with a decrease in oxidation

number to form negative ions;

38 of 144

Redox reactions - 2

(f) describe the redox reactions of metals with

dilute hydrochloric and dilute sulfuric acids;

39 of 144

Redox reactions - 3

(g) interpret and make predictions from redox

equations in terms of oxidation numbers and

electron loss/gain.

40 of 144

1.2.1

 ionisation energies;

 energy levels, shells, sub-shells, orbitals and electron configuration.

41 of 144

Ionisation energies - 1

(a) Define the terms first ionisation energy and

successive ionisation energy;

42 of 144

Ionisation energies - 2

(b) Explain that ionisation energies are

influenced by nuclear charge, electron

shielding and the distance of the outermost

electron from the nucleus;

43 of 144

Ionisation energies - 3

(c) predict from successive ionisation energies of

an element:

(i) the number of electrons in each shell of

an atom,

(ii) the group of the element;

44 of 144

Electrons: electronic energy levels, shells, sub-s

(d) state the number of electrons that can fill the

first four shells;

45 of 144

1.2.1

 ionisation energies;

 energy levels, shells, sub-shells, orbitals and electron configuration.

46 of 144

Electrons: electronic energy levels, shells, sub-s

(e) describe an orbital as a region that can hold

up to two electrons, with opposite spins;

47 of 144

Ionisation energies - 1

(a) Define the terms first ionisation energy and

successive ionisation energy;

48 of 144

Electrons: electronic energy levels, shells, sub-s

(f) describe the shapes of s and p orbitals;

49 of 144

Ionisation energies - 2

(b) Explain that ionisation energies are

influenced by nuclear charge, electron

shielding and the distance of the outermost

electron from the nucleus;

50 of 144

Electrons: electronic energy levels, shells, sub-s

(g) state the number of:

(i) orbitals making up s-, p- and d-subshells,

(ii) electrons that occupy s-, p- and d-subshells

51 of 144

Ionisation energies - 3

(c) predict from successive ionisation energies of

an element:

(i) the number of electrons in each shell of

an atom,

(ii) the group of the element;

52 of 144

Electrons: electronic energy levels, shells, sub-s

(h) describe the relative energies of s-, p- and dorbitals

for the shells 1, 2, 3 and the 4s and

4p orbitals;

53 of 144

Electrons: electronic energy levels, shells, sub-s

(d) state the number of electrons that can fill the

first four shells;

54 of 144

Electrons: electronic energy levels, shells, sub-s

(i) deduce the electron configurations of:

(i) atoms, given the atomic number, up to

Z = 36,

(ii) ions, given the atomic number and ionic

charge, limited to s and p blocks up to Z

= 36;

55 of 144

Electrons: electronic energy levels, shells, sub-s

(e) describe an orbital as a region that can hold

up to two electrons, with opposite spins;

56 of 144

Electrons: electronic energy levels, shells, sub-s

(j) classify the elements into s, p and d blocks.

57 of 144

Electrons: electronic energy levels, shells, sub-s

(f) describe the shapes of s and p orbitals;

58 of 144

1.2.2

1.2.2 Bonding and Structure

 ionic bonding;

 covalent bonding;

 the shapes of simple molecules and ions;

 electronegativity and polarity;

 intermolecular forces.

59 of 144

Electrons: electronic energy levels, shells, sub-s

(g) state the number of:

(i) orbitals making up s-, p- and d-subshells,

(ii) electrons that occupy s-, p- and d-subshells

60 of 144

Ionic bonding - 1

(a) describe the term ionic bonding as

electrostatic attraction between oppositelycharged

ions;

61 of 144

Electrons: electronic energy levels, shells, sub-s

(h) describe the relative energies of s-, p- and dorbitals

for the shells 1, 2, 3 and the 4s and

4p orbitals;

62 of 144

Ionic bonding - 2

(b) construct ‘dot-and-cross’ diagrams, to

describe ionic bonding;

63 of 144

Electrons: electronic energy levels, shells, sub-s

(i) deduce the electron configurations of:

(i) atoms, given the atomic number, up to

Z = 36,

(ii) ions, given the atomic number and ionic

charge, limited to s and p blocks up to Z

= 36;

64 of 144

Ionic bonding - 3

(c) predict ionic charge from the position of an

element in the Periodic Table;

65 of 144

Electrons: electronic energy levels, shells, sub-s

(j) classify the elements into s, p and d blocks.

66 of 144

Ionic bonding - 4

(d) state the formulae for the following ions: NO3,

CO32–, SO42– and NH4+;

67 of 144

Covalent bonding and dative covalent (coordinate)

(e) describe the term covalent bond as a shared

pair of electrons;

68 of 144

1.2.2

1.2.2 Bonding and Structure

 ionic bonding;

 covalent bonding;

 the shapes of simple molecules and ions;

 electronegativity and polarity;

 intermolecular forces.

69 of 144

Ionic bonding - 1

(a) describe the term ionic bonding as

electrostatic attraction between oppositelycharged

ions;

70 of 144

Covalent bonding and dative covalent (coordinate)

(f) construct ‘dot-and-cross’ diagrams to

describe:

(i) single covalent bonding, eg as in H2,

Cl2, HCl, H2O, NH3, CH4, BF3 and SF6,

71 of 144

Covalent bonding and dative covalent (coordinate)

(ii) multiple covalent bonding, eg as in O2,

N2 and CO2,

72 of 144

Ionic bonding - 2

(b) construct ‘dot-and-cross’ diagrams, to

describe ionic bonding;

73 of 144

Covalent bonding and dative covalent (coordinate)

(iii) dative covalent (coordinate) bonding,

eg as in NH4+,

74 of 144

Ionic bonding - 3

(c) predict ionic charge from the position of an

element in the Periodic Table;

75 of 144

Covalent bonding and dative covalent (coordinate)

(iv) molecules and ions analogous to those

specified in (i), (ii) and (iii);

76 of 144

Ionic bonding - 4

(d) state the formulae for the following ions: NO3,

CO32–, SO42– and NH4+;

77 of 144

Covalent bonding and dative covalent (coordinate)

(e) describe the term covalent bond as a shared

pair of electrons;

78 of 144

The shapes of simple molecules and ions - 1

(g) explain that the shape of a simple molecule is

determined by repulsion between electron

pairs surrounding a central atom;

79 of 144

Covalent bonding and dative covalent (coordinate)

(f) construct ‘dot-and-cross’ diagrams to

describe:

(i) single covalent bonding, eg as in H2,

Cl2, HCl, H2O, NH3, CH4, BF3 and SF6,

80 of 144

The shapes of simple molecules and ions - 2

(h) state that lone pairs of electrons repel more

than bonded pairs;

81 of 144

Covalent bonding and dative covalent (coordinate)

(ii) multiple covalent bonding, eg as in O2,

N2 and CO2,

82 of 144

The shapes of simple molecules and ions - 3

(i) explain the shapes of, and bond angles in,

molecules and ions with up to six electron

pairs (including lone pairs) surrounding a

central atom, eg as in:

83 of 144

Covalent bonding and dative covalent (coordinate)

(iii) dative covalent (coordinate) bonding,

eg as in NH4+,

84 of 144

The shapes of simple molecules and ions - 4

(i) BF3 (trigonal planar),

(ii) CH4 and NH4+ (tetrahedral),

(iii) SF6 (octahedral),

85 of 144

The shapes of simple molecules and ions - 5

(iv) NH3 (pyramidal),

(v) H2O (non-linear),

(vi) CO2 (linear);

86 of 144

Covalent bonding and dative covalent (coordinate)

(iv) molecules and ions analogous to those

specified in (i), (ii) and (iii);

87 of 144

The shapes of simple molecules and ions - 6

(j) predict the shapes of, and bond angles in,

molecules and ions analogous to those

specified in (i);

88 of 144

The shapes of simple molecules and ions - 1

(g) explain that the shape of a simple molecule is

determined by repulsion between electron

pairs surrounding a central atom;

89 of 144

Electronegativity and bond polarity - 1

(k) describe the term electronegativity as the

ability of an atom to attract the bonding

electrons in a covalent bond;

90 of 144

The shapes of simple molecules and ions - 2

(h) state that lone pairs of electrons repel more

than bonded pairs;

91 of 144

Electronegativity and bond polarity - 2

(l) explain that a permanent dipole may arise

when covalently-bonded atoms have different

electronegativities, resulting in a polar bond;

92 of 144

The shapes of simple molecules and ions - 3

(i) explain the shapes of, and bond angles in,

molecules and ions with up to six electron

pairs (including lone pairs) surrounding a

central atom, eg as in:

93 of 144

Intermolecular forces - 1

(m) describe intermolecular forces based on

permanent dipoles, as in hydrogen chloride,

and induced dipoles (van der Waals’ forces),

as in the noble gases;

94 of 144

The shapes of simple molecules and ions - 4

(i) BF3 (trigonal planar),

(ii) CH4 and NH4+ (tetrahedral),

(iii) SF6 (octahedral),

95 of 144

Intermolecular forces - 2

(n) describe hydrogen bonding, including the role

of a lone pair, between molecules containing

–OH and –NH groups, ie as in H2O, NH3 and

analogous molecules;

96 of 144

The shapes of simple molecules and ions - 5

(iv) NH3 (pyramidal),

(v) H2O (non-linear),

(vi) CO2 (linear);

97 of 144

Intermolecular forces - 3

describe and explain the anomalous

properties of H2O resulting from hydrogen bonding, eg:

(i) the density of ice compared with water,

(ii) its relatively high freezing point and

boiling point;

98 of 144

The shapes of simple molecules and ions - 6

(j) predict the shapes of, and bond angles in,

molecules and ions analogous to those

specified in (i);

99 of 144

Electronegativity and bond polarity - 1

(k) describe the term electronegativity as the

ability of an atom to attract the bonding

electrons in a covalent bond;

100 of 144

Metallic bonding - 1

(p) describe metallic bonding as the attraction of

positive ions to delocalised electrons;

101 of 144

Electronegativity and bond polarity - 2

(l) explain that a permanent dipole may arise

when covalently-bonded atoms have different

electronegativities, resulting in a polar bond;

102 of 144

Bonding and physical properties - 1

(q) describe structures as:

(i) giant ionic lattices, with strong ionic bonding, ie as in NaCl,

(ii) giant covalent lattices, ie as in diamond

and graphite,

103 of 144

Intermolecular forces - 1

(m) describe intermolecular forces based on

permanent dipoles, as in hydrogen chloride,

and induced dipoles (van der Waals’ forces),

as in the noble gases;

104 of 144

Bonding and physical properties - 2

(iii) giant metallic lattices,

(iv) simple molecular lattices, ie as in I2 and

ice;

105 of 144

Intermolecular forces - 2

(n) describe hydrogen bonding, including the role

of a lone pair, between molecules containing

–OH and –NH groups, ie as in H2O, NH3 and

analogous molecules;

106 of 144

Bonding and physical properties - 3

(r) describe, interpret and/or predict physical

properties, including melting and boiling

points, electrical conductivity and solubility in terms of:

(i) different structures of particles (atoms,

molecules, ions and electrons) and the

forces between them,

107 of 144

Intermolecular forces - 3

describe and explain the anomalous

properties of H2O resulting from hydrogen bonding, eg:

(i) the density of ice compared with water,

(ii) its relatively high freezing point and

boiling point;

108 of 144

Bonding and physical properties - 4

(ii) different types of bonding (ionic

bonding, covalent bonding, metallic

bonding, hydrogen bonding, other

intermolecular interactions);

109 of 144

Bonding and physical properties - 5

(s) deduce the type of structure and bonding

present from given information.

110 of 144

Metallic bonding - 1

(p) describe metallic bonding as the attraction of

positive ions to delocalised electrons;

111 of 144

1.2.

1.3.1 Periodicity

 the Periodic Table;

 trends in physical properties

112 of 144

Bonding and physical properties - 1

(q) describe structures as:

(i) giant ionic lattices, with strong ionic bonding, ie as in NaCl,

(ii) giant covalent lattices, ie as in diamond

and graphite,

113 of 144

Bonding and physical properties - 2

(iii) giant metallic lattices,

(iv) simple molecular lattices, ie as in I2 and

ice;

114 of 144

Bonding and physical properties - 3

(r) describe, interpret and/or predict physical

properties, including melting and boiling

points, electrical conductivity and solubility in terms of:

(i) different structures of particles (atoms,

molecules, ions and electrons) and the

forces between them,

115 of 144

Bonding and physical properties - 4

(ii) different types of bonding (ionic

bonding, covalent bonding, metallic

bonding, hydrogen bonding, other

intermolecular interactions);

116 of 144

Bonding and physical properties - 5

(s) deduce the type of structure and bonding

present from given information.

117 of 144

1.2.

1.3.1 Periodicity

 the Periodic Table;

 trends in physical properties

118 of 144

The structure of the Periodic Table in terms of gr

(a) describe the Periodic Table in terms of the

arrangement of elements:

(i) by increasing atomic (proton) number,

119 of 144

The structure of the Periodic Table in terms of gr

(ii) in periods showing repeating trends in

physical and chemical properties,

(iii) in groups having similar physical and

chemical properties;

120 of 144

The structure of the Periodic Table in terms of gr

(b) describe periodicity in terms of a repeating

pattern across different periods;

121 of 144

The structure of the Periodic Table in terms of gr

(c) explain that atoms of elements in a group

have similar outer shell electron

configurations, resulting in similar properties

122 of 144

Periodicity of physical properties of elements - 1

(d) describe and explain the variation of the first

ionisation energies of elements shown by:

(i) a general increase across a period, in

terms of increasing nuclear charge

123 of 144

Periodicity of physical properties of elements - 2

(ii) a decrease down a group in terms of

electron shielding outweighing

increasing nuclear charge;

124 of 144

Periodicity of physical properties of elements - 3

(e) for the elements of Periods 2 and 3:

(i) describe the variation in electron configurations, atomic radii, melting

points and boiling points,

125 of 144

Periodicity of physical properties of elements - 4

(ii) explain variations in melting and boiling

points in terms of structure and

bonding;

126 of 144

Periodicity of physical properties of elements - 5

(f) interpret data on electron configurations,

atomic radii, first ionisation energies, melting

points and boiling points to demonstrate

periodicity.

127 of 144

1.3.2

1.3.2 Group 2

 redox reactions of Group 2 metals;

 Group 2 compounds.

128 of 144

Redox reactions of Group 2 metals - 1

(a) describe the redox reactions of the Group 2

elements Mg → Ba:

(i) with oxygen,

(ii) with water;

129 of 144

Redox reactions of Group 2 metals - 2

(b) explain the trend in reactivity of Group 2

elements down the group due to the

increasing ease of forming cations, in terms

of atomic size, shielding and nuclear

attraction;

130 of 144

Reactions of Group 2 compounds - 1

(c) describe the action of water on oxides of

elements in Group 2 and state the

approximate pH of any resulting solution

131 of 144

Reactions of Group 2 compounds - 2

(d) describe the thermal decomposition of the

carbonates of elements in Group 2 and the

trend in their ease of decomposition

132 of 144

Reactions of Group 2 compounds - 3

(e) interpret and make predictions from the

chemical and physical properties of Group 2

elements and compounds;

133 of 144

Reactions of Group 2 compounds - 4

(f) explain the use of Ca(OH)2 in agriculture to

neutralise acid soils; the use of Mg(OH)2 in

some indigestion tablets as an antacid.

134 of 144

1.3.3

1.3.3 Group 7

 redox reactions of Group 7 elements;

 halide tests.

135 of 144

Characteristic physical properties - 1

(a) explain, in terms of van der Waals’ forces, the

trend in the boiling points of Cl2, Br2 and I2;

136 of 144

Redox reactions and trends in reactivity of Group

(b) describe the redox reactions, including ionic

equations, of the Group 7 elements Cl2, Br2

and I2 with other halide ions, in the presence

of an organic solvent, to illustrate the relative

reactivity of Group 7 elements;

137 of 144

Redox reactions and trends in reactivity of Group

(c) explain the trend in reactivity of Group 7

elements down the group from the

decreasing ease of forming negative ions, in

terms of atomic size, shielding and nuclear

attraction;

138 of 144

Redox reactions and trends in reactivity of Group

(d) describe the term disproportionation as a

reaction in which an element is

simultaneously oxidised and reduced,

illustrated by:

(i) the reaction of chlorine with water as

used in water purification,

139 of 144

Redox reactions and trends in reactivity of Group

(ii) the reaction of chlorine with cold, dilute

aqueous sodium hydroxide, as used to

form bleach,

(iii) reactions analogous to those specified

in (i) and (ii);

140 of 144

Redox reactions and trends in reactivity of Group

(e) interpret and make predictions from the

chemical and physical properties of the

Group 7 elements and their compounds;

141 of 144

Redox reactions and trends in reactivity of Group

(f) contrast the benefits of chlorine use in water

treatment (killing bacteria) with associated

risks (hazards of toxic chlorine gas and

possible risks from formation of chlorinated

hydrocarbons);

142 of 144

Characteristic reactions of halide ions - 1

(g) describe the precipitation reactions, including

ionic equations, of the aqueous anions Cl–,

Br– and I– with aqueous silver ions, followed

by aqueous ammonia;

143 of 144

Characteristic reactions of halide ions - 2

(h) describe the use of the precipitation reactions

in (g) as a test for different halide ions.

144 of 144