Chemistry AS - Spec Check

1.1.1 Atoms

 atomic structure;

 relative masses.

(a) describe protons, neutrons and electrons in

terms of relative charge and relative mass;

(b) describe the distribution of mass and charge

within an atom;

(c) describe the contribution of protons and

neutrons to the nucleus of an atom, in terms

of atomic (proton) number and mass

(nucleon) number;

(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;

(e) explain the term isotopes as atoms of an

element with different numbers of neutrons

and different masses;

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

measurement of relative masses;

(g) define the terms relative isotopic mass and

relative atomic mass, based on the 12C scale;

(h) calculate the relative atomic mass of an

element given the relative abundances of its

isotopes;

(i) use the terms relative molecular mass and

relative formula mass and calculate values

from relative atomic masses.

1.1.2 Moles and Equations

 the mole;

 reacting masses and equations

(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);

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

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

(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;

(d) calculate empirical and molecular formulae,

using composition by mass and percentage

compositions;

(e) construct balanced chemical equations for

reactions studied and for unfamiliar reactions

given reactants and products;

(f) carry out calculations, using amount of

substance in mol, involving:

(i) mass,

(ii) gas volume,

(iii) solution volume and concentration;

(g) deduce stoichiometric relationships from

calculations;

(h) use the terms concentrated and dilute as

qualitative descriptions for the concentration

of a solution.

1.1.3 Acids

 acids and bases;

 salts

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

aqueous solution;

(b) state the formulae of the common acids:

hydrochloric, sulfuric and nitric acids;

(c) state that common bases are metal oxides,

metal hydroxides and ammonia

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

releases OH– ions in aqueous solution;

(e) state the formulae of the common alkalis:

sodium hydroxide, potassium hydroxide and

aqueous ammonia

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

ion of an acid is replaced by a metal ion or

NH4+;

(g) describe the reactions of an acid with

carbonates, bases and alkalis, to form a salt;

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

from an acid: eg OH– forming H2O; NH3

forming NH4+;

(i) explain the terms anhydrous, hydrated and

water of crystallisation;

(j) calculate the formula of a hydrated salt from

given percentage composition, mass

composition or experimental data;

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

structured titrations.

1.1.4 Redox

 oxidation number;

 redox reactions

(a) apply rules for assigning oxidation number to

atoms in elements, compounds and ions;

(b) describe the terms oxidation and reduction in

terms of:

(i) electron transfer,

(ii) changes in oxidation number;

(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);

(d) write formulae using oxidation numbers;

(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;

(f) describe the redox reactions of metals with

dilute hydrochloric and dilute sulfuric acids;

(g) interpret and make predictions from redox

equations in terms of oxidation numbers and

electron loss/gain.

 ionisation energies;

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

(a) Define the terms first ionisation energy and

successive ionisation energy;

(b) Explain that ionisation energies are

influenced by nuclear charge, electron

shielding and the distance of the outermost

electron from the nucleus;

(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;

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

first four shells;

 ionisation energies;

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

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

up to two electrons, with opposite spins;

(a) Define the terms first ionisation energy and

successive ionisation energy;

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

(b) Explain that ionisation energies are

influenced by nuclear charge, electron

shielding and the distance of the outermost

electron from the nucleus;

(g) state the number of:

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

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

(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;

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

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

4p orbitals;

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

first four shells;

(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;

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

up to two electrons, with opposite spins;

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

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

1.2.2 Bonding and Structure

 ionic bonding;

 covalent bonding;

 the shapes of simple molecules and ions;

 electronegativity and polarity;

 intermolecular forces.

(g) state the number of:

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

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

(a) describe the term ionic bonding as

electrostatic attraction between oppositelycharged

ions;

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

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

4p orbitals;

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

describe ionic bonding;

(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;

(c) predict ionic charge from the position of an

element in the Periodic Table;

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

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

CO32–, SO42– and NH4+;

(e) describe the term covalent bond as a shared

pair of electrons;

1.2.2 Bonding and Structure

 ionic bonding;

 covalent bonding;

 the shapes of simple molecules and ions;

 electronegativity and polarity;

 intermolecular forces.

(a) describe the term ionic bonding as

electrostatic attraction between oppositelycharged

ions;

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

describe:

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

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

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

N2 and CO2,

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

describe ionic bonding;

(iii) dative covalent (coordinate) bonding,

eg as in NH4+,

(c) predict ionic charge from the position of an

element in the Periodic Table;

(iv) molecules and ions analogous to those

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

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

CO32–, SO42– and NH4+;

(e) describe the term covalent bond as a shared

pair of electrons;

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

determined by repulsion between electron

pairs surrounding a central atom;

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

describe:

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

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

(h) state that lone pairs of electrons repel more

than bonded pairs;

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

N2 and CO2,

(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:

(iii) dative covalent (coordinate) bonding,

eg as in NH4+,

(i) BF3 (trigonal planar),

(ii) CH4 and NH4+ (tetrahedral),

(iii) SF6 (octahedral),

(iv) NH3 (pyramidal),

(v) H2O (non-linear),

(vi) CO2 (linear);

(iv) molecules and ions analogous to those

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

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

molecules and ions analogous to those

specified in (i);

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

determined by repulsion between electron

pairs surrounding a central atom;

(k) describe the term electronegativity as the

ability of an atom to attract the bonding

electrons in a covalent bond;

(h) state that lone pairs of electrons repel more

than bonded pairs;

(l) explain that a permanent dipole may arise

when covalently-bonded atoms have different

electronegativities, resulting in a polar bond;

(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:

(m) describe intermolecular forces based on

permanent dipoles, as in hydrogen chloride,

and induced dipoles (van der Waals’ forces),

as in the noble gases;

(i) BF3 (trigonal planar),

(ii) CH4 and NH4+ (tetrahedral),

(iii) SF6 (octahedral),

(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;

(iv) NH3 (pyramidal),

(v) H2O (non-linear),

(vi) CO2 (linear);

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;

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

molecules and ions analogous to those

specified in (i);

(k) describe the term electronegativity as the

ability of an atom to attract the bonding

electrons in a covalent bond;

(p) describe metallic bonding as the attraction of

positive ions to delocalised electrons;

(l) explain that a permanent dipole may arise

when covalently-bonded atoms have different

electronegativities, resulting in a polar bond;

(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,

(m) describe intermolecular forces based on

permanent dipoles, as in hydrogen chloride,

and induced dipoles (van der Waals’ forces),

as in the noble gases;

(iii) giant metallic lattices,

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

ice;

(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;

(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,

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;

(ii) different types of bonding (ionic

bonding, covalent bonding, metallic

bonding, hydrogen bonding, other

intermolecular interactions);

(s) deduce the type of structure and bonding

present from given information.

(p) describe metallic bonding as the attraction of

positive ions to delocalised electrons;

1.3.1 Periodicity

 the Periodic Table;

 trends in physical properties

(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,

(iii) giant metallic lattices,

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

ice;

(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,

(ii) different types of bonding (ionic

bonding, covalent bonding, metallic

bonding, hydrogen bonding, other

intermolecular interactions);

(s) deduce the type of structure and bonding

present from given information.

1.3.1 Periodicity

 the Periodic Table;

 trends in physical properties

(a) describe the Periodic Table in terms of the

arrangement of elements:

(i) by increasing atomic (proton) number,

(ii) in periods showing repeating trends in

physical and chemical properties,

(iii) in groups having similar physical and

chemical properties;

(b) describe periodicity in terms of a repeating

pattern across different periods;

(c) explain that atoms of elements in a group

have similar outer shell electron

configurations, resulting in similar properties

(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

(ii) a decrease down a group in terms of

increasing atomic radius and increasing

electron shielding outweighing

increasing nuclear charge;

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

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

points and boiling points,

(ii) explain variations in melting and boiling

points in terms of structure and

bonding;

(f) interpret data on electron configurations,

atomic radii, first ionisation energies, melting

points and boiling points to demonstrate

periodicity.

1.3.2 Group 2

 redox reactions of Group 2 metals;

 Group 2 compounds.

(a) describe the redox reactions of the Group 2

elements Mg → Ba:

(i) with oxygen,

(ii) with water;

(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;

(c) describe the action of water on oxides of

elements in Group 2 and state the

approximate pH of any resulting solution

(d) describe the thermal decomposition of the

carbonates of elements in Group 2 and the

trend in their ease of decomposition

(e) interpret and make predictions from the

chemical and physical properties of Group 2

elements and compounds;

(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.

1.3.3 Group 7

 redox reactions of Group 7 elements;

 halide tests.

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

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

(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;

(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;

(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,

(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);

(e) interpret and make predictions from the

chemical and physical properties of the

Group 7 elements and their compounds;

(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);

(g) describe the precipitation reactions, including

ionic equations, of the aqueous anions Cl–,

Br– and I– with aqueous silver ions, followed

by aqueous ammonia;

(h) describe the use of the precipitation reactions

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