Chemistry 9701 — specification explorer

1. Particles in the atom and atomic radius

• 1.1.1understand that atoms are mostly empty space surrounding a very small, dense nucleus that contains protons and neutrons; electrons are found in shells in the empty space around the nucleus
• 1.1.2identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses
• 1.1.3understand the terms atomic and proton number; mass and nucleon number
• 1.1.4describe the distribution of mass and charge within an atom
• 1.1.5describe the behaviour of beams of protons, neutrons and electrons moving at the same velocity in an electric field
• 1.1.6determine the numbers of protons, neutrons and electrons present in both atoms and ions given atomic or proton number, mass or nucleon number and charge
• 1.1.6adeduce the presence of bromine and chlorine atoms in a compound using the [M + 2] peak
• 1.1.7state and explain qualitatively the variations in atomic radius and ionic radius across a period and down a group

2. Isotopes

• 1.2.1define the term isotope in terms of numbers of protons and neutrons
• 1.2.2understand the notation y A for isotopes, where is the mass or nucleon number and is the atomic or proton number
• 1.2.3state that and explain why isotopes of the same element have the same chemical properties
• 1.2.4state that and explain why isotopes of the same element have different physical properties, limited to mass and density

3. Electrons, energy levels and atomic orbitals

• 1.3.1

  understand the terms:

  • shells, sub-shells and orbitals
  • principal quantum number (n)
  • ground state, limited to electronic configuration
• 1.3.2describe the number of orbitals making up s, p and d sub-shells, and the number of electrons that can fill s, p and d sub-shells
• 1.3.3describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p sub-shells
• 1.3.4describe the electronic configurations to include the number of electrons in each shell, sub-shell and orbital
• 1.3.5explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion
• 1.3.6determine the electronic configuration of atoms and ions given the atomic or proton number and charge, using either of the following conventions: e.g. for Fe: 1s 2s 2p 3s 3p 3d 4s (full electronic configuration) or [Ar] 3d 4s (shorthand electronic configuration)
• 1.3.7understand and use the electrons in boxes notation e.g. for Fe: [Ar]
• 1.3.8describe and sketch the shapes of s and p orbitals
• 1.3.9describe a free radical as a species with one or more unpaired electrons

4. Ionisation energy

• 1.4.1define and use the term first ionisation energy, IE
• 1.4.2construct equations to represent first, second and subsequent ionisation energies
• 1.4.3identify and explain the trends in ionisation energies across a period and down a group of the Periodic Table
• 1.4.4identify and explain the variation in successive ionisation energies of an element
• 1.4.5understand that ionisation energies are due to the attraction between the nucleus and the outer electron
• 1.4.6explain the factors influencing the ionisation energies of elements in terms of nuclear charge, atomic/ionic radius, shielding by inner shells and sub-shells and spin-pair repulsion
• 1.4.7deduce the electronic configurations of elements using successive ionisation energy data
• 1.4.8deduce the position of an element in the Periodic Table using successive ionisation energy data

1. Relative masses of atoms and molecules

• 2.1.1define the unified atomic mass unit as one twelfth of the mass of a carbon-12 atom
• 2.1.2define relative atomic mass, A , relative isotopic mass, relative molecular mass, M , and relative formula mass in terms of the unified atomic mass unit

2. The mole and the Avogadro constant

• 2.2.1define and use the term mole in terms of the Avogadro constant

3. Formulas

• 2.3.1

  write formulas of ionic compounds from ionic charges and oxidation numbers (shown by a Roman numeral), including: HCO , PO:

  • (a) the prediction of ionic charge from the position of an element in the Periodic Table
  • 2– 2– – + 2+ +
  • (b) recall of the names and formulas for the following ions: NO , CO , SO , OH , NH , Zn , Ag ,
  • 3–
• 2.3.2

  (a) write and construct equations (which should be balanced), including ionic equations (which should not include spectator ions):

  • (b) use appropriate state symbols in equations
• 2.3.3define and use the terms empirical and molecular formula
• 2.3.4understand and use the terms anhydrous, hydrated and water of crystallisation
• 2.3.5calculate empirical and molecular formulas, using given data

4. Reacting masses and volumes (of solutions and gases)

• 2.4.1

  perform calculations including use of the mole concept, involving: (When performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question. When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified (see also Mathematical requirements section).):

  • (a) reacting masses (from formulas and equations) including percentage yield calculations
  • (b) volumes of gases (e.g. in the burning of hydrocarbons)
  • (c) volumes and concentrations of solutions
  • (d) limiting reagent and excess reagent
  • (e) deduce stoichiometric relationships from calculations such as those in 2.4.1(a)–(d)

1. Electronegativity and bonding

• 3.1.1define electronegativity as the power of an atom to attract electrons to itself
• 3.1.2explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells
• 3.1.3state and explain the trends in electronegativity across a period and down a group of the Periodic Table
• 3.1.4use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds (the presence of covalent character in some ionic compounds will not be assessed) (Pauling electronegativity values will be given where necessary)

2. Ionic bonding

• 3.2.1define ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions)
• 3.2.2describe ionic bonding including the examples of sodium chloride, magnesium oxide and calcium fluoride

3. Metallic bonding

• 3.3.1define metallic bonding as the electrostatic attraction between positive metal ions and delocalised electrons

4. Covalent bonding and coordinate (dative covalent) bonding

• 3.4.1

  define covalent bonding as electrostatic attraction between the nuclei of two atoms and a shared pair of electrons:

  • (a) describe covalent bonding in molecules including:
  • hydrogen, H
  • oxygen, O
  • nitrogen, N
  • chlorine, C l
  • hydrogen chloride, HC l
  • carbon dioxide, CO
  • ammonia, NH
  • methane, CH
  • ethane, C H
• 3.4.2

  6 sulfur dioxide, SO , phosphorus pentachloride, PC l , and sulfur hexafluoride, SF hydrogen chloride gases to form the ammonium ion, NH , and in the A l C l molecule 2 (a) describe covalent bonds in terms of orbital overlap giving σ and π bonds:

  • ethene, C H
  • (b) understand that elements in period 3 can expand their octet including in the compounds
  • (c) describe coordinate (dative covalent) bonding, including in the reaction between ammonia and
  • σ bonds are formed by direct overlap of orbitals between the bonding atoms
  • π bonds are formed by the sideways overlap of adjacent p orbitals above and below the σ bond
  • (b) describe how the σ and π bonds form in molecules including H₂, C₂H₆, C₂H₄, HCN and N₂
  • (c) use the concept of hybridisation to describe sp, sp² and sp³ orbitals
• 3.4.3

  (a) define the terms: gaseous state molecules:

  • bond energy as the energy required to break one mole of a particular covalent bond in the
  • bond length as the internuclear distance of two covalently bonded atoms
  • (b) use bond energy values and the concept of bond length to compare the reactivity of covalent

5. Shapes of molecules

• 3.5.1

  state and explain the shapes of, and bond angles in, molecules by using VSEPR theory, including as simple examples:

  • BF (trigonal planar, 120°)
  • CO (linear, 180°)
  • CH (tetrahedral, 109.5°)
  • NH (pyramidal, 107°)
  • H O (non-linear, 104.5°)
  • SF (octahedral, 90°)
  • PF (trigonal bipyramidal, 120° and 90°)
• 3.5.2predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.5.1

6. Intermolecular forces, electronegativity and bond properties

• 3.6.1

  (a) describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples:

  • (b) use the concept of hydrogen bonding to explain the anomalous properties of H₂O (ice and water):
  • its relatively high melting and boiling points
  • its relatively high surface tension
  • the density of the solid ice compared with the liquid water
• 3.6.2use the concept of electronegativity to explain bond polarity and dipole moments of molecules
• 3.6.3

  (a) describe van der Waals’ forces as the intermolecular forces between molecular entities other than those due to bond formation, and use the term van der Waals’ forces as a generic term to describe all intermolecular forces permanent dipole–permanent dipole forces between molecules where hydrogen is bonded to a highly electronegative atom:

  • (b) describe the types of van der Waals’ forces:
  • instantaneous dipole–induced dipole (id-id) forces, also called London dispersion forces
  • permanent dipole–permanent dipole (pd-pd) forces, including hydrogen bonding
  • (c) describe hydrogen bonding and understand that hydrogen bonding is a special case of
• 3.6.4state that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces

7. Dot-and-cross diagrams

• 3.7.1use dot-and-cross diagrams to illustrate ionic, covalent and coordinate bonding including the representation of any compounds stated in 3.4 and 3.5 (dot-and-cross diagrams may include species with atoms which have an expanded octet or species with an odd number of electrons)

1. The gaseous state: ideal and real gases and pV = nRT

• 4.1.1explain the origin of pressure in a gas in terms of collisions between gas molecules and the wall of the container
• 4.1.2understand that ideal gases have zero particle volume and no intermolecular forces of attraction
• 4.1.3state and use the ideal gas equation pV = nRT in calculations, including in the determination of M

2. Bonding and structure

• 4.2.1

  describe, in simple terms, the lattice structure of a crystalline solid which is:

  • (a) giant ionic, including sodium chloride and magnesium oxide
  • (b) simple molecular, including iodine, buckminsterfullerene C and ice
  • (c) giant molecular, including silicon( IV ) oxide, graphite and diamond
  • (d) giant metallic, including copper
• 4.2.2describe, interpret and predict the effect of different types of structure and bonding on the physical properties of substances, including melting point, boiling point, electrical conductivity and solubility
• 4.2.3deduce the type of structure and bonding present in a substance from given information

1. Enthalpy change, Δ H

• 5.1.1understand that chemical reactions are accompanied by enthalpy changes and these changes can be exothermic ( Δ H is negative) or endothermic ( Δ H is positive)
• 5.1.2construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy
• 5.1.3

  define and use the terms: neutralisation, Δ H neut:

  • (a) standard conditions (this syllabus assumes that these are 298 K and 101 kPa) shown by .
  • (b) enthalpy change with particular reference to: reaction, Δ H , formation, Δ H , combustion, Δ H ,
• 5.1.4understand that energy transfers occur during chemical reactions because of the breaking and making of chemical bonds
• 5.1.5use bond energies ( Δ H positive, i.e. bond breaking) to calculate enthalpy change of reaction, Δ H
• 5.1.6understand that some bond energies are exact and some bond energies are averages
• 5.1.7calculate enthalpy changes from appropriate experimental results, including the use of the relationships q = mc Δ T and Δ H = – mc Δ T /

2. Hess’s law

• 5.2.1apply Hess’s law to construct simple energy cycles
• 5.2.2

  carry out calculations using cycles and relevant energy terms, including:

  • (a) determining enthalpy changes that cannot be found by direct experiment
  • (b) use of bond energy data

1. Redox processes: electron transfer and changes in oxidation number (oxidation state)

• 6.1.1calculate oxidation numbers of elements in compounds and ions
• 6.1.2use changes in oxidation numbers to help balance chemical equations
• 6.1.3explain and use the terms redox, oxidation, reduction and disproportionation in terms of electron transfer and changes in oxidation number
• 6.1.4explain and use the terms oxidising agent and reducing agent
• 6.1.5use a Roman numeral to indicate the magnitude of the oxidation number of an element

1. Chemical equilibria: reversible reactions, dynamic equilibrium

• 7.1.1

  (a) understand what is meant by a reversible reaction reactions being equal and the concentration of reactants and products remaining constant:

  • (b) understand what is meant by dynamic equilibrium in terms of the rate of forward and reverse
  • (c) understand the need for a closed system in order to establish dynamic equilibrium
• 7.1.2define Le Chatelier’s principle as: if a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change
• 7.1.3use Le Chatelier’s principle to deduce qualitatively (from appropriate information) the effects of changes in temperature, concentration, pressure or presence of a catalyst on a system at equilibrium
• 7.1.4deduce expressions for equilibrium constants in terms of concentrations, K
• 7.1.5use the terms mole fraction and partial pressure
• 7.1.6deduce expressions for equilibrium constants in terms of partial pressures, K (use of the relationship between K and K is not required)
• 7.1.7use the K and K expressions to carry out calculations (such calculations will not require the solving of quadratic equations)
• 7.1.8calculate the quantities present at equilibrium, given appropriate data
• 7.1.9state whether changes in temperature, concentration or pressure or the presence of a catalyst affect the value of the equilibrium constant for a reaction
• 7.1.10describe and explain the conditions used in the Haber process and the Contact process, as examples of the importance of an understanding of dynamic equilibrium in the chemical industry and the application of Le Chatelier’s principle

2. Brønsted–Lowry theory of acids and bases

• 7.2.1state the names and formulas of the common acids, limited to hydrochloric acid, HCl, sulfuric acid, H₂SO₄, nitric acid, HNO₃, ethanoic acid, CH₃COOH
• 7.2.24 3 3 2 state the names and formulas of the common alkalis, limited to sodium hydroxide, NaOH, potassium hydroxide, KOH, ammonia, NH
• 7.2.3describe the Brønsted–Lowry theory of acids and bases
• 7.2.4describe strong acids and strong bases as fully dissociated in aqueous solution and weak acids and weak bases as partially dissociated in aqueous solution
• 7.2.5appreciate that water has pH of 7, acid solutions pH of below 7 and alkaline solutions pH of above 7
• 7.2.6explain qualitatively the differences in behaviour between strong and weak acids including the reaction with a reactive metal and difference in pH values by use of a pH meter, universal indicator or conductivity
• 7.2.7understand that neutralisation reactions occur when H (aq) and OH (aq) form H O(l)
• 7.2.8understand that salts are formed in neutralisation reactions
• 7.2.9sketch the pH titration curves of titrations using combinations of strong and weak acids with strong and weak alkalis
• 7.2.10select suitable indicators for acid-alkali titrations, given appropriate data (p K values will not be used)

1. Rate of reaction

• 8.1.1explain and use the term rate of reaction, frequency of collisions, effective collisions and non-effective collisions
• 8.1.2explain qualitatively, in terms of frequency of effective collisions, the effect of concentration and pressure changes on the rate of a reaction
• 8.1.3use experimental data to calculate the rate of a reaction

2. Effect of temperature on reaction rates and the concept of activation energy

• 8.2.1define activation energy, E , as the minimum energy required for a collision to be effective
• 8.2.2sketch and use the Boltzmann distribution to explain the significance of activation energy
• 8.2.3explain qualitatively, in terms both of the Boltzmann distribution and of frequency of effective collisions, the effect of temperature change on the rate of a reaction

3. Homogeneous and heterogeneous catalysts

• 8.3.1

  explain and use the terms catalyst and catalysis: activation energy effective catalyst:

  • (a) explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower
  • (b) explain this catalytic effect in terms of the Boltzmann distribution
  • (c) construct and interpret a reaction pathway diagram, for a reaction in the presence and absence of an

1. Periodicity of physical properties of the elements in Period

• 9.1.1describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements
• 9.1.2explain the variation in melting point and electrical conductivity in terms of the structure and bonding of the elements

2. Periodicity of chemical properties of the elements in Period

• 9.2.1describe, and write equations for, the reactions of the elements with oxygen (to give Na O, MgO, A l O ,
• 9.2.22 3 P O , SO ), chlorine (to give NaC l , MgC l , A l C l , SiC l , PC l ) and water (Na and Mg only) 2 state and explain the variation in the oxidation number of the oxides (Na O, MgO, A l O , P O , SO and SO only) and chlorides (NaC l , MgC l , A l C l , SiC l , PC l only) in terms of their outer shell (valence shell)
• 9.2.32 3 4 5 electrons 3 describe, and write equations for, the reactions, if any, of the oxides Na O, MgO, A l O , SiO , P O , SO and SO with water including the likely pHs of the solutions obtained
• 9.2.4describe, explain, and write equations for, the acid / base behaviour of the oxides Na O, MgO, A l O , P O , SO and SO and the hydroxides NaOH, Mg(OH) and A l (OH) including, where relevant, amphoteric behaviour in reactions with acids and bases (sodium hydroxide only)
• 9.2.5describe, explain, and write equations for, the reactions of the chlorides NaC l , MgC l , A l C l , SiC l , PC l with water including the likely pHs of the solutions obtained
• 9.2.6explain the variations and trends in 9.2.2, 9.2.3, 9.2.4 and 9.2.5 in terms of bonding and electronegativity
• 9.2.7suggest the types of chemical bonding present in the chlorides and oxides from observations of their chemical and physical properties

3. Chemical periodicity of other elements

• 9.3.1predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity
• 9.3.2deduce the nature, possible position in the Periodic Table and identity of unknown elements from given information about physical and chemical properties

1. Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and

• 10.1.1describe, and write equations for, the reactions of the elements with oxygen, water and dilute hydrochloric and sulfuric acids
• 10.1.2describe, and write equations for, the reactions of the oxides, hydroxides and carbonates with water and dilute hydrochloric and sulfuric acids
• 10.1.3describe, and write equations for, the thermal decomposition of the nitrates and carbonates, to include the trend in thermal stabilities
• 10.1.4describe, and make predictions from, the trends in physical and chemical properties of the elements involved in the reactions in 10.1.1 and the compounds involved in 10.1.2, 10.1.3 and 10.1.5
• 10.1.5state the variation in the solubilities of the hydroxides and sulfates

1. Physical properties of the Group 17 elements

• 11.1.1describe the colours and the trend in volatility of chlorine, bromine and iodine
• 11.1.2describe and explain the trend in the bond strength of the halogen molecules
• 11.1.3interpret the volatility of the elements in terms of instantaneous dipole–induced dipole forces

2. The chemical properties of the halogen elements and the hydrogen halides

• 11.2.1describe the relative reactivity of the elements as oxidising agents
• 11.2.2describe the reactions of the elements with hydrogen and explain their relative reactivity in these reactions
• 11.2.3describe the relative thermal stabilities of the hydrogen halides and explain these in terms of bond strengths

3. Some reactions of the halide ions

• 11.3.1describe the relative reactivity of halide ions as reducing agents
• 11.3.2

  describe and explain the reactions of halide ions with:

  • (a) aqueous silver ions followed by aqueous ammonia (the formation and formula of the [Ag(NH ) ]
• 11.3.3

  2 complex is not required):

  • (b) concentrated sulfuric acid, to include balanced chemical equations

4. The reactions of chlorine

• 11.4.1describe and interpret, in terms of changes in oxidation number, the reaction of chlorine with cold and with hot aqueous sodium hydroxide and recognise these as disproportionation reactions
• 11.4.2explain, including by use of an equation, the use of chlorine in water purification to include the production of the active species HOC l and C l O which kill bacteria

1. Nitrogen and sulfur

• 12.1.1explain the lack of reactivity of nitrogen, with reference to triple bond strength and lack of polarity
• 12.1.2

  describe and explain:

  • (a) the basicity of ammonia, using the Brønsted–Lowry theory
  • (b) the structure of the ammonium ion and its formation by an acid–base reaction
  • (c) the displacement of ammonia from ammonium salts by an acid–base reaction
• 12.1.3state and explain the natural and man-made occurrences of oxides of nitrogen and their catalytic removal from the exhaust gases of internal combustion engines
• 12.1.4understand that atmospheric oxides of nitrogen (NO and NO ) can react with unburned hydrocarbons to form peroxyacetyl nitrate, PAN, which is a component of photochemical smog
• 12.1.5describe the role of NO and NO in the formation of acid rain both directly and in their catalytic role in the oxidation of atmospheric sulfur dioxide

1. Formulas, functional groups and the naming of organic compounds

• 13.1.1define the term hydrocarbon as a compound made up of C and H atoms only
• 13.1.2understand that alkanes are simple hydrocarbons with no functional group
• 13.1.3understand that the compounds in the table on pages 29 and 30 contain a functional group which dictates their physical and chemical properties
• 13.1.4interpret and use the general, structural, displayed and skeletal formulas of the classes of compound stated in the table on pages 29 and 30
• 13.1.5understand and use systematic nomenclature of simple aliphatic organic molecules with functional groups detailed in the table on pages 29 and 30, up to six carbon atoms (six plus six for esters, straight chains only for esters and nitriles)
• 13.1.6deduce the molecular and/or empirical formula of a compound, given its structural, displayed or skeletal formula

2. Characteristic organic reactions

• 13.2.1

  interpret and use the following terminology associated with types of organic compounds and reactions: (in equations for organic redox reactions, the symbol [O] can be used to represent one atom of oxygen from an oxidising agent and the symbol [H] to represent one atom of hydrogen from a reducing agent):

  • (a) homologous series
  • (b) saturated and unsaturated
  • (c) homolytic and heterolytic fission
  • (d) free radical, initiation, propagation, termination
  • (e) nucleophile, electrophile, nucleophilic, electrophilic
  • (f) addition, substitution, elimination, hydrolysis, condensation
  • (g) oxidation and reduction
• 13.2.2

  understand and use the following terminology associated with types of organic mechanisms: (in organic reaction mechanisms, the use of curly arrows to represent movement of electron pairs is expected; the arrow should begin at a bond or a lone pair of electrons):

  • (a) free-radical substitution
  • (b) electrophilic addition
  • (c) nucleophilic substitution
  • (d) nucleophilic addition

3. Shapes of organic molecules; σ and π bonds

• 13.3.1describe organic molecules as either straight-chained, branched or cyclic
• 13.3.23 2 describe and explain the shape of, and bond angles in, molecules containing sp, sp and sp hybridised atoms
• 13.3.3describe the arrangement of σ and π bonds in molecules containing sp, sp and sp hybridised atoms
• 13.3.4understand and use the term planar when describing the arrangement of atoms in organic molecules, for example ethene

4. Isomerism: structural isomerism and stereoisomerism

• 13.4.1describe structural isomerism and its division into chain, positional and functional group isomerism
• 13.4.2describe stereoisomerism and its division into geometrical (cis/trans) and optical isomerism (use of E/Z nomenclature is acceptable but is not required)
• 13.4.3describe geometrical (cis/trans) isomerism in alkenes, and explain its origin in terms of restricted rotation due to the presence of π bonds
• 13.4.4explain what is meant by a chiral centre and that such a centre gives rise to two optical isomers (enantiomers) (Candidates should appreciate that compounds can contain more than one chiral centre, but knowledge of meso compounds, or nomenclature such as diastereoisomers is not required.)
• 13.4.5identify chiral centres and geometrical (cis/trans) isomerism in a molecule of given structural formula including cyclic compounds
• 13.4.6deduce the possible isomers for an organic molecule of known molecular formula

1. Alkanes

• 14.1.1

  recall the reactions (reagents and conditions) by which alkanes can be produced:

  • (a) addition of hydrogen to an alkene in a hydrogenation reaction, H (g) and Pt/Ni catalyst and heat
  • (b) cracking of a longer chain alkane, heat with A l O
• 14.1.2

  3 2 describe: by the reactions of ethane:

  • (a) the complete and incomplete combustion of alkanes
  • (b) the free-radical substitution of alkanes by C l or Br in the presence of ultraviolet light, as exemplified
• 14.1.3describe the mechanism of free-radical substitution with reference to the initiation, propagation and termination steps
• 14.1.4suggest how cracking can be used to obtain more useful alkanes and alkenes of lower M from heavier crude oil fractions
• 14.1.5understand the general unreactivity of alkanes, including towards polar reagents in terms of the strength of the C–H bonds and their relative lack of polarity
• 14.1.6recognise the environmental consequences of carbon monoxide, oxides of nitrogen and unburnt hydrocarbons arising from the combustion of alkanes in the internal combustion engine and of their catalytic removal

2. Alkenes

• 14.2.1

  recall the reactions (including reagents and conditions) by which alkenes can be produced:

  • (a) elimination of HX from a halogenoalkane by ethanolic NaOH and heat
  • (b) dehydration of an alcohol, by using a heated catalyst (e.g. A l O ) or a concentrated acid
• 14.2.2

  3 (e.g. concentrated H SO ) 2 describe the following reactions of alkenes:

  • (ii) steam, H O(g) and H PO catalyst
  • (iii) a hydrogen halide, HX(g), at room temperature
  • (iv) a halogen, X double bond and the identities of the subsequent products to determine the position of alkene linkages in larger molecules:
  • (c) cracking of a longer chain alkane
  • (a) the electrophilic addition of
  • (i) hydrogen in a hydrogenation reaction, H (g) and Pt/Ni catalyst and heat
  • (b) the oxidation by cold dilute acidified KMnO to form the diol
  • (c) the oxidation by hot concentrated acidified KMnO leading to the rupture of the carbon–carbon
  • (d) addition polymerisation exemplified by the reactions of ethene and propene
• 14.2.3describe the use of aqueous bromine to show the presence of a C=C bond
• 14.2.4describe the mechanism of electrophilic addition in alkenes, using bromine/ethene and hydrogen bromide/propene as examples
• 14.2.5describe and explain the inductive effects of alkyl groups on the stability of primary, secondary and tertiary cations formed during electrophilic addition (this should be used to explain Markovnikov addition)

1. Halogenoalkanes

• 15.1.1

  recall the reactions (reagents and conditions) by which halogenoalkanes can be produced:

  • (a) the free-radical substitution of alkanes by C l or Br in the presence of ultraviolet light, as exemplified
• 15.1.2

  2 by the reactions of ethane concentrated H PO ; or with PC l and heat; or with PC l ; or with SOC l:

  • (b) electrophilic addition of an alkene with a halogen, X , or hydrogen halide, HX(g), at room temperature
  • (c) substitution of an alcohol, e.g. by reaction with HX(g); or with KC l and concentrated H SO or
• 15.1.3

  4 3 5 2 2 classify halogenoalkanes into primary, secondary and tertiary 3 describe the following nucleophilic substitution reactions: exemplified by bromoethane:

  • (a) the reaction with NaOH(aq) and heat to produce an alcohol
  • (b) the reaction with KCN in ethanol and heat to produce a nitrile
  • (c) the reaction with NH in ethanol heated under pressure to produce an amine
  • (d) the reaction with aqueous silver nitrate in ethanol as a method of identifying the halogen present as
• 15.1.4describe the elimination reaction with NaOH in ethanol and heat to produce an alkene as exemplified by bromoethane
• 15.1.5describe the S 1 and S 2 mechanisms of nucleophilic substitution in halogenoalkanes including the inductive effects of alkyl groups
• 15.1.6recall that primary halogenoalkanes tend to react via the S 2 mechanism; tertiary halogenoalkanes via the S 1 mechanism; and secondary halogenoalkanes by a mixture of the two, depending on structure
• 15.1.7describe and explain the different reactivities of halogenoalkanes (with particular reference to the relative strengths of the C–X bonds as exemplified by the reactions of halogenoalkanes with aqueous silver nitrates)

1. Alcohols

• 16.1.1

  recall the reactions (reagents and conditions) by which alcohols can be produced:

  • (a) electrophilic addition of steam to an alkene, H O(g) and H PO catalyst
• 16.1.2

  3 4 2 describe: H SO or concentrated H PO ; or with PC l and heat; or with PC l ; or with SOC l (ii) carboxylic acids by refluxing (primary alcohols give aldehydes which can be further oxidised to carboxylic acids, secondary alcohols give ketones, tertiary alcohols cannot be oxidised) exemplified by ethanol:

  • (b) reaction of alkenes with cold dilute acidified potassium manganate( VII ) to form a diol
  • (c) substitution of a halogenoalkane using NaOH(aq) and heat
  • (d) reduction of an aldehyde or ketone using NaBH or LiA l H
  • (e) reduction of a carboxylic acid using LiA l H
  • (f) hydrolysis of an ester using dilute acid or dilute alkali and heat
  • (a) the reaction with oxygen (combustion)
  • (b) substitution to form halogenoalkanes, e.g. by reaction with HX(g); or with KC l and concentrated
  • (c) the reaction with Na(s)
  • (d) oxidation with acidified K Cr O or acidified KMnO to:
  • (i) carbonyl compounds by distillation
  • (e) dehydration to an alkene, by using a heated catalyst, e.g. A l O or a concentrated acid
  • (f) formation of esters by reaction with carboxylic acids and concentrated H SO as catalyst as
• 16.1.3

  (a) classify alcohols as primary, secondary and tertiary alcohols, to include examples with more than one alcohol group from orange to green:

  • (b) state characteristic distinguishing reactions, e.g. mild oxidation with acidified K Cr O , colour change
• 16.1.4deduce the presence of a CH CH(OH)– group in an alcohol, CH CH(OH)–R, from its reaction with alkaline I (aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO
• 16.1.5explain the acidity of alcohols compared with water

1. Aldehydes and ketones

• 17.1.1

  recall the reactions (reagents and conditions) by which aldehydes and ketones can be produced:

  • (a) the oxidation of primary alcohols using acidified K Cr O or acidified KMnO and distillation to
• 17.1.2

  2 7 4 produce aldehydes produce ketones 2 describe: hydroxynitriles as exemplified by ethanal and propanone:

  • (b) the oxidation of secondary alcohols using acidified K Cr O or acidified KMnO and distillation to
  • (a) the reduction of aldehydes and ketones using NaBH or LiA l H to produce alcohols
  • (b) the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat to produce
• 17.1.3describe the mechanism of the nucleophilic addition reactions of hydrogen cyanide with aldehydes and ketones in 17.1.2(b)
• 17.1.4describe the use of 2,4-dinitrophenylhydrazine (2,4-DNPH reagent) to detect the presence of carbonyl compounds
• 17.1.5deduce the nature (aldehyde or ketone) of an unknown carbonyl compound from the results of simple tests (Fehling’s and Tollens’ reagents; ease of oxidation)
• 17.1.6deduce the presence of a CH CO – group in an aldehyde or ketone, CH CO–R, from its reaction with alkaline I (aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO

1. Carboxylic acids

• 18.1.1

  recall the reactions by which carboxylic acids can be produced:

  • (a) oxidation of primary alcohols and aldehydes with acidified K Cr O or acidified KMnO and refluxing
• 18.1.2

  2 7 4 2 describe:

  • (b) hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification
  • (c) hydrolysis of esters with dilute acid or dilute alkali and heat followed by acidification
  • (a) the redox reaction with reactive metals to produce a salt and H (g)
  • (b) the neutralisation reaction with alkalis to produce a salt and H O(l )
  • (c) the acid–base reaction with carbonates to produce a salt and H O(l) and CO (g)
  • (d) esterification with alcohols with concentrated H SO as catalyst
  • (e) reduction by LiA l H to form a primary alcohol

2. Esters

• 18.2.1

  recall the reaction (reagents and conditions) by which esters can be produced: catalyst:

  • (a) the condensation reaction between an alcohol and a carboxylic acid with concentrated H₂SO₄ as
• 18.2.2describe the hydrolysis of esters by dilute acid and by dilute alkali and heat

1. Primary amines

• 19.1.1

  recall the reactions by which amines can be produced:

  • (a) reaction of a halogenoalkane with NH₃ in ethanol heated under pressure
• 19.1.2Classification of amines will not be tested at AS Level.

2. Nitriles and hydroxynitriles

• 19.2.1

  recall the reactions by which nitriles can be produced:

  • (a) reaction of a halogenoalkane with KCN in ethanol and heat
• 19.2.2

  recall the reactions by which hydroxynitriles can be produced:

  • (a) the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat
• 19.2.3describe the hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification to produce a carboxylic acid

1. Addition polymerisation

• 20.1.1describe addition polymerisation as exemplified by poly(ethene) and poly(chloroethene), PVC
• 20.1.2deduce the repeat unit of an addition polymer obtained from a given monomer
• 20.1.3identify the monomer(s) present in a given section of an addition polymer molecule
• 20.1.4recognise the difficulty of the disposal of poly(alkene)s, i.e. non-biodegradability and harmful combustion products

1. Organic synthesis

• 21.1.1

  for an organic molecule containing several functional groups:

  • (a) identify organic functional groups using the reactions in the syllabus
  • (b) predict properties and reactions
• 21.1.2devise multi-step synthetic routes for preparing organic molecules using the reactions in the syllabus
• 21.1.3analyse a given synthetic route in terms of type of reaction and reagents used for each step of it, and possible by-products
• 21.1.4Analysis

1. Infrared spectroscopy

• 22.1.1analyse an infrared spectrum of a simple molecule to identify functional groups (see the Data section for the functional groups required)

2. Mass spectrometry

• 22.2.1analyse mass spectra in terms of m / e values and isotopic abundances (knowledge of the working of the mass spectrometer is not required)
• 22.2.2calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum
• 22.2.3deduce the molecular mass of an organic molecule from the molecular ion peak in a mass spectrum
• 22.2.4suggest the identity of molecules formed by simple fragmentation in a given mass spectrum
• 22.2.5deduce the number of carbon atoms, n , in a compound using the [M + 1] peak and the formula 100 × abundance of [M + 1] ion