Upper secondaryCode 9701
Chemistry
90 topics and 366 objectives, each one something you should be able to do.
Written by Kindred in its own plain words. Your syllabus is still the official list, so check it for anything exact.
What this subject asks you to be able to do
1.1Particles in the atom and atomic radius7 objectives
- understand that most of an atom is empty space, with a tiny, dense nucleus made of protons and neutrons at its centre; electrons sit in shells within that empty region around the nucleus
- identify protons, neutrons and electrons and describe each one by its relative charge and relative mass
- understand what is meant by the terms atomic number and proton number, and by mass number and nucleon number
- describe how mass and charge are arranged inside an atom
- describe how beams of protons, neutrons and electrons, all travelling at the same velocity, behave when passing through an electric field
- work out how many protons, neutrons and electrons atoms and ions contain, when supplied with the atomic or proton number, mass or nucleon number and charge
- state and qualitatively account for how atomic and ionic radii change moving along a period and moving down a group
1.2Isotopes4 objectives
- define an isotope by referring to how many protons and neutrons it contains
- understand how isotopes are written as x yA, with x standing for the mass or nucleon number and y standing for the atomic or proton number
- state that isotopes of a single element have identical chemical properties, and account for why this is so
- state that isotopes of one element have different physical properties, limited to mass and density, and account for why this happens
1.3Electrons, energy levels and atomic orbitals9 objectives
- understand what is meant by shells, sub-shells and orbitals, the principal quantum number (n) and the ground state, with the scope limited to electronic configuration
- describe how many orbitals there are in s, p and d sub-shells, and how many electrons s, p and d sub-shells can each hold
- outline the sequence, from lowest to highest energy, of the sub-shells in the first three shells along with the 4s and 4p sub-shells
- outline electronic configurations so that they show how many electrons sit in every shell, sub-shell and orbital
- account for electronic configurations by reference to electron energy and inter-electron repulsion
- work out the electronic configuration of an atom or ion from its atomic or proton number and its charge, writing it in either the full form, e.g. for Fe: 1s2 2s2 2p6 3s2 3p6 3d6 4s2, or the shorthand form, [Ar] 3d6 4s2
- understand and apply the notation known as electrons in boxes, e.g. for Fe: [Ar]
- describe the shape of an s orbital and of a p orbital, and sketch both
- describe a free radical: any species that has at least one unpaired electron
1.4Ionisation energy8 objectives
- define the expression first ionisation energy, IE, and apply it
- construct equations showing the first, second and each later ionisation energy
- identify, and account for, how ionisation energies vary along a period and through a group of the Periodic Table
- identify, and account for, the differences between successive ionisation energies for a single element
- understand that the ionisation energies of an atom depend on how strongly its nucleus attracts the outer electron
- explain, with reference to spin-pair repulsion, shielding from inner shells and sub-shells, atomic/ionic radius and nuclear charge, why elements differ in ionisation energy
- deduce which electronic configurations elements have, drawing on successive ionisation energy data
- deduce where an element sits in the Periodic Table, drawing on successive ionisation energy data
2.1Relative masses of atoms and molecules2 objectives
- define the unified atomic mass unit, taking it as one twelfth of a carbon-12 atom's mass
- define, by reference to the unified atomic mass unit, each of these: relative isotopic mass, relative atomic mass, Ar, relative molecular mass, Mr, and relative formula mass
2.2The mole and the Avogadro constant1 objective
- define the mole by reference to the Avogadro constant, and use the word in that sense
2.3Formulas5 objectives
- write the formula of an ionic compound using ionic charges and oxidation numbers (indicated by a Roman numeral), including: (a) predicting an ion's charge from where an element sits in the Periodic Table (b) recalling the names and formulas of the following ions: NO3–, CO32–, SO42–, OH–, NH4+, Zn2+, Ag+, HCO3–, PO43–
- write and build equations that are balanced, including ionic equations that do not include spectator ions (b) apply suitable state symbols within equations
- define empirical formula and molecular formula, and use both terms
- appreciate the meaning of the terms anhydrous, hydrated and water of crystallisation, and apply them
- work out empirical and molecular formulas from supplied data
2.4Reacting masses and volumes (of solutions and gases)1 objective
- carry out calculations, including ones that rely on the mole concept, covering: (a) reacting masses found from formulas and equations, and percentage yield (b) gas volumes, e.g. when hydrocarbons burn (c) concentrations and volumes of solutions (d) which reagent is limiting and which is in excess (e) deduce stoichiometric relationships by working through calculations such as those in 2.4.1(a)–(d)
3.1Electronegativity and bonding4 objectives
- define electronegativity as an atom's ability to pull electrons towards itself
- account for the causes behind differences in electronegativity between elements, linking them to nuclear charge, atomic radius, and shielding from inner shells and sub-shells
- give and account for how electronegativity varies from left to right across a period and from top to bottom within one group of the Periodic Table
- use the difference between Pauling electronegativity values to predict whether ionic or covalent bonds will form (whether some ionic compounds show covalent character is not assessed) (Pauling electronegativity values are supplied where needed)
3.2Ionic bonding2 objectives
- define ionic bonding as an electrostatic attraction that acts between ions of opposite charge (cations carrying a positive charge and anions carrying a negative charge)
- describe the nature of ionic bonding as seen in compounds including sodium chloride, magnesium oxide and calcium fluoride
3.3Metallic bonding1 objective
- define metallic bonding as the electrostatic force holding positively charged metal ions together with delocalised electrons
3.4Covalent bonding and coordinate (dative covalent) bonding3 objectives
- define covalent bonding as the electrostatic attraction between a pair of electrons shared by two atoms and the nuclei of those atoms (a) describe the covalent bonding in molecules including: • hydrogen, H2 • oxygen, O2 • nitrogen, N2 • chlorine, Cl2 • hydrogen chloride, HCl • carbon dioxide, CO2 • ammonia, NH3 • methane, CH4 • ethane, C2H6 • ethene, C2H4 (b) understand that an element of period 3 can expand its octet, as happens in compounds including sulfur dioxide, SO2, phosphorus pentachloride, PCl5, and sulfur hexafluoride, SF6 (c) describe coordinate (dative covalent) bonding, including when gaseous ammonia and hydrogen chloride combine to give the ammonium ion, NH4+, and within the Al2Cl6 molecule
- describe how covalent bonds arise from orbital overlap, yielding σ and π bonds: • σ bonds result from orbitals overlapping directly between the bonding atoms • π bonds result from adjacent p orbitals overlapping sideways, above and below the σ bond (b) describe the way σ and π bonds are created in molecules including H₂, C₂H₆, C₂H₄, HCN and N₂ (c) apply the idea of hybridisation when describing sp, sp² and sp³ orbitals
- define these terms: • bond energy, the energy taken in when one mole of a specific covalent bond is broken in the gaseous state • bond length, the internuclear distance between two atoms joined by a covalent bond (b) use bond energy values, together with the idea of bond length, to compare how reactive covalent molecules are
3.5Shapes of molecules2 objectives
- state and account for, using VSEPR theory, the shape and bond angles of a molecule, including these simple cases: BF3, trigonal planar, 120°; CO2, linear, 180°; CH4, tetrahedral, 109.5°; NH3, pyramidal, 107°; H2O, non-linear, 104.5°; SF6, octahedral, 90°; PF5, trigonal bipyramidal, 120° and 90°
- predict the shapes of molecules and ions, together with their bond angles, that are similar to those set out in 3.5.1
3.6Intermolecular forces, electronegativity and bond properties4 objectives
- describe hydrogen bonding, limited to molecules with N–H and O–H groups, including the simple cases of ammonia and water (b) use the idea of hydrogen bonding to explain why H₂O (ice and water) behaves unusually: • its melting and boiling points are relatively high • its surface tension is also relatively high • how the density of solid ice compares with that of liquid water
- use ideas about electronegativity to explain how polar a bond is and what dipole moment a molecule has
- describe van der Waals’ forces as any intermolecular forces acting between molecular entities apart from those that arise from forming bonds, and use van der Waals’ forces as an umbrella name covering every kind of intermolecular force (b) describe the kinds of van der Waals’ forces: • id-id forces (instantaneous dipole–induced dipole), also known as London dispersion forces • pd-pd forces (permanent dipole–permanent dipole), including hydrogen bonding (c) describe hydrogen bonding, and understand it as a particular type of pd-pd force that acts between molecules in which hydrogen is joined to a highly electronegative atom
- state that, generally speaking, the bonding in ionic, covalent and metallic substances is stronger than intermolecular forces
3.7Dot-and-cross diagrams1 objective
- use dot-and-cross diagrams to show ionic, covalent and coordinate bonding, including for any compound named in 3.4 and 3.5 (these diagrams can cover species containing an atom with an expanded octet, or species whose electron count is odd)
4.1The gaseous state: ideal and real gases and pV = nRT3 objectives
- account for the source of pressure in a gas by considering impacts of gas molecules against the walls of the container
- understand that particles in an ideal gas occupy zero volume and exert no intermolecular forces of attraction
- state and apply the ideal gas equation, pV = nRT, to problems, including when finding Mr
4.2Bonding and structure3 objectives
- describe, in straightforward language, the way particles are packed in each of these kinds of crystal lattice: (a) the giant ionic type, including sodium chloride and magnesium oxide (b) the simple molecular type, including iodine, buckminsterfullerene C60 and ice (c) the giant molecular type, including silicon(IV) oxide, graphite and diamond (d) the giant metallic type, including copper
- describe, interpret and predict how the kind of structure and bonding a substance has affects its physical properties, including melting point, boiling point, electrical conductivity and solubility
- deduce which kind of structure and bonding a substance contains, using information provided
5.1Enthalpy change, ΔH7 objectives
- understand that enthalpy changes take place alongside chemical reactions, and that a change may be exothermic, with ΔH negative, or endothermic, with ΔH positive
- construct and interpret a reaction pathway diagram with regard to the reaction's enthalpy change and its activation energy
- define and apply the following terms: (a) standard conditions, taken as 298K and 101kPa and marked by ⦵; (b) enthalpy change, particularly in relation to: reaction, ΔHr, formation, ΔHf, combustion, ΔHc, neutralisation, ΔHneut
- appreciate that energy is exchanged in chemical reactions as a result of bonds being broken and formed
- apply bond energies (with ΔH positive, i.e. for bond breaking) in working out the enthalpy change of reaction, ΔHr
- understand that a bond energy is in some cases exact and in other cases an average
- work out enthalpy changes from suitable experimental data, including applying the relationships q = mcΔT and ΔH = –mcΔT/n
5.2Hess’s law2 objectives
- use Hess’s law to build straightforward energy cycles
- carry out calculations involving cycles and suitable energy terms, including: (a) working out enthalpy changes which cannot be obtained through direct experimental measurement (b) applying bond energy data
6.1Redox processes: electron transfer and changes in oxidation number (oxidation state)5 objectives
- determine the oxidation numbers that elements take on within compounds and ions
- apply changes in oxidation numbers to aid the balancing of chemical equations
- explain the terms redox, oxidation, reduction and disproportionation through the transfer of electrons and changes in oxidation number, and use these terms
- explain what an oxidising agent and a reducing agent are, and use both terms
- use a Roman numeral to show how large an element's oxidation number is
7.1Chemical equilibria: reversible reactions, dynamic equilibrium12 objectives
- understand the definition of a reversible reaction
- understand dynamic equilibrium as the condition in which the forward and reverse reactions proceed at identical rates while the concentrations of reactants and products stay steady
- understand why a closed system is necessary for dynamic equilibrium to be set up
- define Le Chatelier's principle as follows: after a change is imposed on a system at dynamic equilibrium, the equilibrium position shifts in the direction that minimises that change
- use Le Chatelier’s principle to infer, qualitatively and from suitable information, the consequences of altering temperature, concentration or pressure, or of the presence of a catalyst, for a system at equilibrium
- deduce expressions giving equilibrium constants using concentrations, Kc
- use mole fraction and partial pressure correctly as terms
- deduce the expressions for equilibrium constants that use partial pressures, Kp (applying the link between Kp and Kc is not necessary)
- use the expressions for Kc and Kp in calculations (these calculations will not involve solving quadratic equations)
- determine the quantities present once equilibrium is reached, from appropriate data
- state whether altering temperature, concentration or pressure, or introducing a catalyst, changes the numerical value of the equilibrium constant of a reaction
- describe and explain the conditions chosen for the Haber process and the Contact process, which show why an understanding of dynamic equilibrium and the use of Le Chatelier’s principle matter to the chemical industry
7.2Brønsted–Lowry theory of acids and bases10 objectives
- give the name and formula of each typical acid, limited to hydrochloric acid, HCl, sulfuric acid, H2SO4, nitric acid, HNO3, ethanoic acid, CH3COOH
- recall the name and formula of each typical alkali, limited to sodium hydroxide, NaOH, potassium hydroxide, KOH, ammonia, NH3
- describe how the Brønsted–Lowry theory defines an acid and a base
- outline how strong acids and strong bases dissociate fully in aqueous solution, while weak acids and weak bases dissociate partially
- appreciate that water has a pH of 7, that acid solutions have a pH below 7 and that alkaline solutions have a pH above 7
- account for, in qualitative terms, how strong and weak acids differ in behaviour, including their reactions with a reactive metal and their pH values measured using a pH meter, universal indicator or conductivity
- appreciate that neutralisation takes place when H+(aq) and OH–(aq) combine to give H2O(l)
- understand that neutralisation reactions produce salts
- sketch the pH titration curves for each pairing of a strong or weak acid with a strong or weak alkali
- select a suitable indicator for an acid–alkali titration from appropriate data (pKa values do not feature)
8.1Rate of reaction3 objectives
- explain the meaning of each of these terms, and use them: the rate of reaction, the frequency of collisions, and effective and non-effective collisions
- account for qualitatively how raising or lowering concentration or pressure alters reaction rate, linking this to the frequency of effective collisions
- use results from experiments to work out a reaction's rate
8.2Effect of temperature on reaction rates and the concept of activation energy3 objectives
- define activation energy, EA, as the smallest amount of energy that colliding particles need for their collision to be effective
- sketch the Boltzmann distribution and use it to account for the importance of activation energy
- account for, in qualitative terms and drawing on both the Boltzmann distribution and the frequency of effective collisions, how changing temperature alters how quickly a reaction goes
8.3Homogeneous and heterogeneous catalysts1 objective
- explain the terms catalyst and catalysis and use them correctly; explain that when a catalyst is present a reaction follows a different mechanism, i.e. one with a lower activation energy; explain the catalytic effect by means of the Boltzmann distribution; construct and interpret the reaction pathway diagram of a reaction, both with and without an effective catalyst
9.1Periodicity of physical properties of the elements in Period 32 objectives
- outline, in qualitative terms, how atomic radius, ionic radius, melting point and electrical conductivity change across the elements, and indicate the periodicity shown by these changes
- account for how melting point and electrical conductivity vary across the elements, using their structure and bonding as the explanation
9.2Periodicity of chemical properties of the elements in Period 37 objectives
- describe what happens when each element reacts with oxygen (forming Na2O, MgO, Al2O3, P4O10 or SO2), with chlorine (forming NaCl, MgCl2, AlCl3, SiCl4 or PCl5) and with water (Na and Mg only), giving equations for these reactions
- state and explain how oxidation number changes across the oxides (only Na2O, MgO, Al2O3, P4O10, SO2 and SO3) and the chlorides (only NaCl, MgCl2, AlCl3, SiCl4, PCl5), explained by the electrons in their outer (valence) shell
- describe the oxides Na2O, MgO, Al2O3, SiO2, P4O10, SO2 and SO3 as they react with water, if they react at all, and give equations for these reactions, including the probable pHs of the solutions formed
- describe and explain how the oxides Na2O, MgO, Al2O3, P4O10, SO2 and SO3, and the hydroxides NaOH, Mg(OH)2 and Al(OH)3, behave as acids or bases, writing equations for this, including any amphoteric behaviour when they react with acids and bases (sodium hydroxide only)
- describe and account for what happens when the chlorides NaCl, MgCl2, AlCl3, SiCl4 and PCl5 are added to water, write equations for these reactions, including the probable pHs of the solutions produced
- explain the trends and variations described in 9.2.2, 9.2.3, 9.2.4 and 9.2.5 by reference to bonding and electronegativity
- suggest, from what is observed of their chemical and physical properties, which kinds of chemical bonding the chlorides and oxides contain
9.3Chemical periodicity of other elements2 objectives
- predict, drawing on what is known about chemical periodicity, the typical properties of an element belonging to a particular group
- deduce what an unknown element is, its nature and where it might sit in the Periodic Table, using information supplied on its physical and chemical properties
10.1Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds5 objectives
- outline what happens when the elements react with oxygen and with water, and when they react with dilute hydrochloric and sulfuric acids, giving equations for these reactions
- outline the reactions of the oxides, hydroxides and carbonates with water and with dilute hydrochloric and sulfuric acids, giving the equations for these reactions
- describe the thermal decomposition of nitrates and of carbonates, giving equations, including the trend in thermal stability
- describe the trends shown by the elements that react in 10.1.1, and by the compounds in 10.1.2, 10.1.3 and 10.1.5, in their physical and chemical properties, and make predictions from these trends
- state how the solubility of the hydroxides and sulfates changes
11.1Physical properties of the Group 17 elements3 objectives
- outline the colours of chlorine, bromine and iodine and how their volatility changes from one to the next
- outline and account for the change in the strength of the bonds within halogen molecules as you move down the group
- interpret how volatile the elements are, drawing on instantaneous dipole–induced dipole forces
11.2The chemical properties of the halogen elements and the hydrogen halides3 objectives
- describe how the elements compare in relative reactivity when they act as oxidising agents
- describe how each element reacts with hydrogen, and explain the relative reactivity the elements show in these reactions
- describe how the hydrogen halides compare in relative thermal stability, and account for these differences using bond strengths
11.3Some reactions of the halide ions2 objectives
- outline how strongly each halide ion acts as a reducing agent, relative to the others
- describe and account for what happens when halide ions are treated first with aqueous silver ions and then with aqueous ammonia (the making and formula of the [Ag(NH3)2]+ complex are not needed), and when they are treated with concentrated sulfuric acid, including balanced equations
11.4The reactions of chlorine2 objectives
- describe and interpret, by following the changes in oxidation number, how chlorine reacts with aqueous sodium hydroxide, both cold and hot, and recognise both as disproportionation reactions
- account for chlorine's role in purifying water, including by way of an equation, covering the formation of the active species HOCl and ClO– that destroy bacteria
12.1Nitrogen and sulfur5 objectives
- account for why nitrogen is so unreactive, referring to the strength of its triple bond and the absence of polarity
- describe and explain: (a) the basicity of ammonia, explained with the Brønsted–Lowry theory (b) the ammonium ion's structure, and how an acid–base reaction forms it (c) how an acid–base reaction displaces ammonia from ammonium salts
- state and explain where oxides of nitrogen come from, both naturally and through human activity, and how catalysts remove them from internal combustion engine exhaust gases
- appreciate that nitrogen oxides in the atmosphere, NO and NO2, can combine with unburned hydrocarbons to yield peroxyacetyl nitrate, PAN, which forms part of photochemical smog
- describe how NO and NO2 contribute to acid rain, both directly and by acting as catalysts in the oxidation of sulfur dioxide in the air
13.1Formulas, functional groups and the naming of organic compounds6 objectives
- define hydrocarbon to mean a compound whose molecules are built from C and H atoms only
- understand alkanes to be simple hydrocarbons that contain no functional group
- understand that each compound tabulated on pages 29 and 30 has a functional group, and that this group governs its physical and chemical properties
- interpret the general, structural, displayed and skeletal formulas of each class of compound tabulated on pages 29 and 30, and use them
- understand systematic nomenclature and use it to name simple aliphatic organic molecules that carry the functional groups tabulated on pages 29 and 30, with no more than six carbon atoms (six plus six in esters; for esters and nitriles, only straight chains)
- deduce a compound's molecular and/or empirical formula from its structural, displayed or skeletal formula
13.2Characteristic organic reactions2 objectives
- interpret the following terms, which relate to kinds of organic compound and reaction, and use them: (a) homologous series (b) saturated and unsaturated compounds (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 (when writing equations for organic redox reactions, [O] may stand for a single oxygen atom provided by an oxidising agent, and [H] for a single hydrogen atom provided by a reducing agent)
- understand the following terms for kinds of organic mechanism, and use them: (a) free-radical substitution (b) electrophilic addition (c) nucleophilic substitution (d) nucleophilic addition (when writing organic reaction mechanisms, show each movement of an electron pair with a curly arrow that starts from a bond or from a lone pair of electrons)
13.3Shapes of organic molecules; σ and π bonds4 objectives
- describe any organic molecule as one of three types: straight-chained, branched or cyclic
- describe and explain the shape taken by molecules made up of sp, sp2 and sp3 hybridised atoms, together with the bond angles present in them
- outline how σ and π bonds are positioned within molecules that include sp, sp2 and sp3 hybridised atoms
- understand and apply the word planar when describing how atoms are arranged in organic molecules, for example ethene
13.4Isomerism: structural isomerism and stereoisomerism6 objectives
- outline what structural isomerism is and how it splits into chain, positional and functional group isomerism
- outline stereoisomerism and how it is subdivided into geometrical (cis/trans) and optical types (the E/Z naming system may be used, though it is not required)
- outline geometrical (cis/trans) isomerism in alkenes, and account for its origin through restricted rotation caused by π bonds
- explain what a chiral centre is and that one such centre produces two optical isomers (enantiomers) (be aware that a compound may have more than one chiral centre; meso compounds, and names such as diastereoisomers, are not required)
- identify, in a molecule whose structural formula is given, any chiral centres and any geometrical (cis/trans) isomerism, including in cyclic compounds
- deduce the isomers an organic molecule could have, starting from its molecular formula
14.1Alkanes6 objectives
- recall the routes, with reagents and conditions, for making alkanes: (a) adding hydrogen across an alkene in a hydrogenation reaction, using H2(g), a Pt/Ni catalyst and heat (b) cracking a longer chain alkane by heating it with Al2O3
- describe: (a) how alkanes undergo complete and incomplete combustion (b) how alkanes undergo free-radical substitution with Cl2 or Br2 when ultraviolet light is present, taking the reactions of ethane as the example
- describe the free-radical substitution mechanism through its initiation, propagation and termination steps
- suggest how heavier fractions of crude oil can be cracked to give alkanes and alkenes of lower Mr that are more useful
- understand why alkanes are, in general, unreactive, including towards polar reagents, given the strength of their C–H bonds and the relatively low polarity of those bonds
- recognise the effects on the environment of the unburnt hydrocarbons, carbon monoxide and oxides of nitrogen released when alkanes burn in an internal combustion engine, and the effects of removing them catalytically
14.2Alkenes5 objectives
- recall how alkenes can be made, including the reagents and conditions: (a) elimination of HX from a halogenoalkane on heating with ethanolic NaOH (b) dehydration of an alcohol over a heated catalyst (e.g. Al2O3) or with a concentrated acid (e.g. concentrated H2SO4) (c) cracking a longer-chain alkane
- describe these reactions of alkenes: (a) electrophilic addition of (i) hydrogen, H2(g), with a Pt/Ni catalyst and heating, as a hydrogenation reaction (ii) steam, H2O(g), using an H3PO4 catalyst (iii) a hydrogen halide, HX(g), at room temperature (iv) a halogen, X2 (b) oxidation with cold dilute acidified KMnO4, forming the diol (c) oxidation with hot concentrated acidified KMnO4, which splits the carbon–carbon double bond, after which the products formed are identified to locate the alkene linkages in larger molecules (d) addition polymerisation, with ethene and propene as examples
- describe how aqueous bromine is used to test whether a C=C bond is present
- describe the electrophilic addition mechanism in alkenes, taking the reactions of bromine with ethene and of hydrogen bromide with propene as examples
- describe and explain how alkyl groups, through their inductive effects, alter the stability of the tertiary, secondary and primary cations formed in electrophilic addition, and use this to explain Markovnikov addition
15.1Halogenoalkanes7 objectives
- state the reagents and conditions used to make halogenoalkanes: (a) free-radical substitution of alkanes with Cl2 or Br2 under ultraviolet light, for example with ethane (b) electrophilic addition to an alkene using a halogen, X2, or hydrogen halide, HX(g), at room temperature (c) substitution of an alcohol, e.g. with HX(g); with KCl and concentrated H2SO4 or concentrated H3PO4; with PCl3 and heat; with PCl5; or with SOCl2
- classify halogenoalkanes as primary, secondary or tertiary
- describe these nucleophilic substitution reactions: (a) producing an alcohol when heated with NaOH(aq) (b) forming a nitrile when KCN in ethanol is heated (c) forming an amine when NH3 in ethanol is heated under pressure (d) the reaction with aqueous silver nitrate dissolved in ethanol, used to identify which halogen is present, using bromoethane as an example
- describe the elimination reaction in which NaOH dissolved in ethanol, with heating, forms an alkene, using bromoethane as an example
- describe and explain the mechanisms SN1 and SN2 by which halogenoalkanes undergo nucleophilic substitution, including the inductive effects arising from alkyl groups
- state that primary halogenoalkanes usually follow the SN2 mechanism; tertiary ones follow the SN1 mechanism; and secondary ones use a mix of both, depending on their structure
- describe and explain why halogenoalkanes differ in reactivity, focusing on how the strengths of their C–X bonds compare, as shown by their reactions with aqueous silver nitrates
16.1Alcohols5 objectives
- recall the reagents and conditions for producing alcohols via: (a) electrophilic addition of steam to an alkene, with H2O(g) and an H3PO4 catalyst (b) cold dilute acidified potassium manganate(VII) acting on alkenes, forming a diol (c) substituting a halogenoalkane by heating it with NaOH(aq) (d) reducing an aldehyde or ketone with NaBH4 or LiAlH4 (e) reducing a carboxylic acid with LiAlH4 (f) hydrolysing an ester with dilute acid or dilute alkali and heat
- describe: (a) combustion, the reaction with oxygen (b) substitution to give halogenoalkanes, e.g. using HX(g); or KCl with concentrated H2SO4 or concentrated H3PO4; or PCl3 with heat; or PCl5; or SOCl2 (c) how an alcohol reacts with Na(s) (d) oxidising it with acidified K2Cr2O7 or acidified KMnO4 into: (i) carbonyl compounds, collected by distillation (ii) carboxylic acids, obtained by heating under reflux (a primary alcohol forms an aldehyde that can be oxidised further to a carboxylic acid; a secondary alcohol forms a ketone; a tertiary alcohol cannot be oxidised) (e) dehydration to an alkene over a heated catalyst, e.g. Al2O3, or with a concentrated acid (f) making esters by reacting the alcohol with carboxylic acids, with concentrated H2SO4 as catalyst, as shown by ethanol
- classify alcohols as primary, secondary or tertiary, including examples that carry more than one alcohol group (b) state the reactions, typical of each type, that tell them apart, e.g. mild oxidation with acidified K2Cr2O7, where the colour turns from orange to green
- deduce that an alcohol contains a CH3CH(OH)– group, written as CH3CH(OH)–R, from its reaction with alkaline I2(aq) producing a yellow precipitate of tri-iodomethane alongside an RCO2– ion
- account for the acidity of alcohols when set against water
17.1Aldehydes and ketones6 objectives
- recall the reagents and conditions that produce aldehydes and ketones: (a) oxidising a primary alcohol with acidified K2Cr2O7 or acidified KMnO4, distilling the product to give an aldehyde (b) oxidising a secondary alcohol with acidified K2Cr2O7 or acidified KMnO4, distilling the product to give a ketone
- describe: (a) reducing aldehydes and ketones with NaBH4 or LiAlH4 to form alcohols (b) reacting aldehydes and ketones with HCN, using KCN as a catalyst and heating, to form hydroxynitriles, with ethanal and propanone as examples
- describe the nucleophilic addition mechanism by which hydrogen cyanide reacts with aldehydes and ketones, as in 17.1.2(b)
- describe the way 2,4-dinitrophenylhydrazine (2,4-DNPH reagent) is employed in establishing whether carbonyl compounds are present
- deduce whether an unknown carbonyl compound is an aldehyde or a ketone from the outcomes of simple tests, namely Fehling’s and Tollens’ reagents and ease of oxidation
- deduce that an aldehyde or ketone, CH3CO–R, contains a CH3CO– group, based on its reaction with alkaline I2(aq) that forms a yellow precipitate of tri-iodomethane and an ion, RCO2–
18.1Carboxylic acids2 objectives
- recall the ways carboxylic acids can be made: (a) oxidising primary alcohols and aldehydes by refluxing with acidified K2Cr2O7 or acidified KMnO4 (b) hydrolysing nitriles using dilute acid or dilute alkali, then acidifying (c) hydrolysing esters by heating with dilute acid or dilute alkali, then acidifying
- describe: (a) the redox reaction between reactive metals and acids, giving a salt together with H2(g) (b) neutralisation of acids by alkalis, where the products are a salt and H2O(l) (c) the acid–base reaction of acids with carbonates, yielding a salt, H2O(l) and CO2(g) (d) esterification of acids with alcohols, using concentrated H2SO4 as a catalyst (e) reduction with LiAlH4, which forms a primary alcohol
18.2Esters2 objectives
- recall the reagents and conditions used to synthesise esters: (a) the condensation reaction of an alcohol with a carboxylic acid, using concentrated H2SO4 as a catalyst
- describe how esters undergo hydrolysis when treated with dilute acid, or with dilute alkali while being heated
19.1Primary amines1 objective
- recall the routes through which amines may be made: (a) a halogenoalkane treated with NH3 dissolved in ethanol and heated under pressure
19.2Nitriles and hydroxynitriles3 objectives
- recall the routes for making nitriles: (a) a halogenoalkane heated with KCN in ethanol
- recall the methods for making hydroxynitriles: (a) aldehydes and ketones combined with HCN, using KCN as a catalyst and warming the mixture
- describe how nitriles undergo hydrolysis by means of dilute acid or dilute alkali, with acidification afterwards, to yield a carboxylic acid
20.1Addition polymerisation4 objectives
- outline addition polymerisation using poly(ethene) and poly(chloroethene), PVC, as examples
- deduce the repeating unit that comes from a particular monomer in an addition polymer
- identify which monomer(s) make up a specified portion of an addition polymer molecule
- recognise why poly(alkene)s are hard to dispose of, namely their non-biodegradability and the harmful combustion products that burning them gives off
21.1Organic synthesis3 objectives
- where one organic molecule contains several functional groups: (a) identify those organic functional groups from the reactions covered in this course (b) predict the molecule's properties and reactions
- devise synthetic routes of more than one step to make organic molecules, drawing on the reactions covered in this course
- analyse a provided sequence of reactions, identifying for every stage the reaction category and the reagents involved, together with any side products that may form
22.1Infrared spectroscopy1 objective
- analyse the infrared spectrum of an uncomplicated molecule so as to pick out the functional groups it contains
22.2Mass spectrometry6 objectives
- analyse a mass spectrum using its m/e values and the isotopic abundances it shows
- work out an element's relative atomic mass when given the proportions in which its isotopes occur, or its mass spectrum
- deduce an organic molecule's molecular mass by reading the molecular ion peak of its mass spectrum
- suggest what molecules are produced by simple fragmentation as shown in a particular mass spectrum
- deduce how many carbon atoms, n, a compound holds by means of the [M + 1]+ peak, applying n = 100 × abundance of [M + 1]+ ion / (1.1 × abundance of M+ ion)
- deduce whether a compound contains bromine and chlorine atoms by making use of the [M + 2]+ peak
23.1Lattice energy and Born-Haber cycles7 objectives
- define the following terms and use them: (a) enthalpy change of atomisation, ΔHat (b) lattice energy, ΔHlatt (for gaseous ions coming together as a solid lattice)
- define first electron affinity, EA, and use this term correctly
- account for the influences on the electron affinities of elements
- describe and explain the trends shown by electron affinity in Group 16 and in Group 17
- construct a Born–Haber cycle for an ionic solid, and use it
- carry out Born–Haber cycle calculations
- explain qualitatively how the size of a lattice energy depends on ionic charge and on ionic radius
23.2Enthalpies of solution and hydration4 objectives
- define the term enthalpy change as it applies to hydration, ΔHhyd, and to solution, ΔHsol, and use it
- construct an energy cycle that connects lattice energy with the enthalpy changes of solution and of hydration, and use it
- carry out calculations based on each energy cycle described in 23.2.2
- explain, in qualitative terms, how an ion's charge and radius affect the numerical size of its enthalpy change of hydration
23.3Entropy change, ΔS3 objectives
- define entropy, S, as how many different arrangements are possible for the particles in a given system and for their energy
- predict and explain the sign of each entropy change that takes place: (a) while a substance changes state, e.g. melting, boiling and dissolving (and the reverse of each) (b) during a change in temperature (c) in a reaction where the number of gaseous molecules alters
- work out the entropy change, ΔS, of a reaction from the standard entropies, S⦵, of its reactants and of its products, using ΔS⦵ = ΣS⦵ (products) – ΣS⦵ (reactants)
23.4Gibbs free energy change, ΔG4 objectives
- give the Gibbs equation ΔG⦵ = ΔH⦵ – TΔS⦵ and apply it
- work out values by carrying out calculations with the equation ΔG⦵ = ΔH⦵ – TΔS⦵
- state, from the sign of ΔG, whether a process or reaction is feasible
- predict how a change in temperature bears on whether a reaction is feasible, with the standard enthalpy and entropy changes provided
24.1Electrolysis4 objectives
- predict which substances are released at the electrodes during electrolysis, using whether the electrolyte is molten or aqueous, where each species sits in the redox series (electrode potential) and its concentration
- state and apply the expression F = Le, which connects the Faraday constant (F), the Avogadro constant (L) and the electron's charge (e)
- calculate: (a) how much charge flows in electrolysis, from Q = It (b) the mass and/or volume of a substance set free by electrolysis
- describe how a value for the Avogadro constant is obtained through an electrolytic technique
24.2Standard electrode potentials E⦵, standard cell potentials E⦵cell and the Nernst equation9 objectives
- define what is meant by each of these: (a) standard electrode (reduction) potential (b) standard cell potential
- outline the techniques employed to determine standard electrode potentials for: (a) a metal or non-metal placed in contact with its ions in aqueous solution (b) ions of one element that exist in different oxidation states
- work out a standard cell potential by putting together two standard electrode potentials
- use standard cell potentials in order to: (a) deduce each electrode's polarity, and from that explain/deduce the direction electrons take through a simple cell's external circuit (b) predict whether a reaction is feasible
- deduce from the E⦵ values how elements, compounds and ions compare with one another as oxidising agents or reducing agents, in terms of their reactivity
- construct full redox equations from the appropriate half-equations
- predict, in qualitative terms, how E, the value of an electrode potential, responds to changes in the concentrations of aqueous ions
- apply the Nernst equation, e.g. E = E⦵ + (0.059/z) log [oxidised species]/[reduced species], to predict in quantitative terms how an electrode potential depends on aqueous ion concentrations
- understand what the equation ΔG⦵ = –nE⦵cell F means and use it
25.1Acids and bases14 objectives
- understand what conjugate acids and conjugate bases are, and use the terms
- define conjugate acid–base pairs and pick out such pairs within reactions
- define mathematically pH, Ka, pKa and Kw, and apply them in calculations (Kb and the relation Kw = Ka × Kb is something you will not be tested on)
- work out the [H+(aq)] and pH values of the following: (a) strong acids (b) strong alkalis (c) weak acids
- define a buffer solution
- account for the way a buffer solution is prepared
- account for how buffer solutions regulate pH, with chemical equations included in those accounts
- describe, and account for, the applications of buffer solutions, including the part HCO3– plays in regulating pH in blood
- work out the pH of buffer solutions from suitable data
- understand what solubility product, Ksp, means, and use the term
- write out an expression that represents Ksp
- work out Ksp from concentrations, and the reverse
- understand and apply the common ion effect in accounting for why a compound's solubility differs in a solution containing a common ion
- carry out calculations that combine Ksp values with the concentration of a common ion
25.2Partition coefficients3 objectives
- state the meaning of the term partition coefficient, Kpc
- work out and apply a partition coefficient for a setup where the solute has the same physical state in both solvents
- understand that the numerical value of a partition coefficient depends on how polar the solute and the chosen solvents are
26.1Simple rate equations, orders of reaction and rate constants11 objectives
- explain each of these terms and use it: rate equation, rate constant, order of reaction, overall order of reaction, half-life, intermediate and rate-determining step
- understand, and apply, rate equations taking the form rate = k [A]m[B]n (with m and n each being 0, 1 or 2)
- deduce a reaction's order by reading concentration–time graphs, or by using data gathered through the initial rates and half-life methods
- interpret experimental results shown as graphs, including concentration–time and rate–concentration graphs
- work out an initial rate from concentration data
- construct a rate equation
- demonstrate an understanding that a first-order reaction's half-life does not depend on concentration
- apply a first-order reaction's half-life when carrying out calculations
- calculate a rate constant's numerical value, e.g. via: (a) combining the initial rates with the rate equation (b) applying the half-life, t½, and k = 0.693/t½
- where a reaction has several steps: (a) propose a mechanism agreeing with the rate equation and with the overall equation (b) forecast the order arising from a supplied mechanism together with its rate-determining step (c) derive a rate equation from a provided mechanism and its rate-determining step, for a specified reaction (d) spot an intermediate or catalyst in a given mechanism (e) pinpoint the rate-determining step using a rate equation alongside a mechanism provided
- describe in qualitative terms how raising or lowering the temperature alters the rate constant, and therefore the rate of a reaction
26.2Homogeneous and heterogeneous catalysts3 objectives
- explain that a catalyst may be either homogeneous or heterogeneous
- describe how a heterogeneous catalyst works, covering adsorption of the reactants, weakening of bonds and desorption of the products, for example: (a) iron, used in the Haber process (b) palladium, platinum and rhodium, which catalyse the removal of nitrogen oxides from car engine exhaust gases
- describe how a homogeneous catalyst works by taking part in one step and being re-formed in a later one, with examples: (a) oxides of nitrogen in the atmosphere, catalysing the oxidation of sulfur dioxide in the air (b) Fe2+ or Fe3+ in the I–/S2O82– reaction
27.1Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds2 objectives
- describe and explain qualitatively the trend in thermal stability shown by the nitrates and carbonates, including how ionic radius affects how strongly the large anion is polarised
- describe and explain qualitatively how, for the hydroxides and sulfates, solubility and the enthalpy change of solution, ΔH⦵sol, vary, in terms of how large the lattice energy is compared with the enthalpy change of hydration
28.1General physical and chemical properties of the first row of transition elements, titanium to copper6 objectives
- define what is meant by a transition element: a d-block element that produces one or more stable ions containing incomplete d orbitals
- sketch how the 3dxy and 3dz² orbitals are shaped in space
- understand that transition elements display the following characteristics: (a) oxidation states that can vary (b) acting as catalysts (c) producing complex ions (d) producing compounds that are coloured
- account for the existence of variable oxidation states in transition elements by reference to the near-equal energies of the 3d and 4s sub-shells
- explain why transition elements act as catalysts: they have two or more stable oxidation states, and they have vacant d orbitals which are energetically accessible and able to form dative bonds to ligands
- explain why transition elements form complex ions, by reference to their vacant, energetically accessible d orbitals
28.2General characteristic chemical properties of the first set of transition elements, titanium to copper10 objectives
- describe and explain how transition elements react with ligands to give complexes, including those of copper(II) and cobalt(II) ions formed from water and ammonia molecules, or from hydroxide and chloride ions
- define a ligand as any species with a lone pair of electrons which it uses to form a dative covalent bond with a metal atom/ion at the centre
- understand these kinds of ligand, including the examples given, and use the terms: (a) monodentate ligands, such as H2O, NH3, Cl– and CN– (b) bidentate ligands, such as 1,2-diaminoethane, en, H2NCH2CH2NH2, and the ethanedioate ion, C2O42– (c) polydentate ligands, such as EDTA4–
- define a complex as a molecule or ion made up of a metal atom or ion at the centre, bonded to one or more ligands
- describe the shape and bond angles (the geometry) of linear, square planar, tetrahedral and octahedral complexes of transition elements
- state the meaning of coordination number (b) predict a complex ion's formula and charge from the metal ion, the charge or oxidation state of that ion, the ligand, and either the coordination number or the geometry
- account for, in qualitative terms, the fact that ligand exchange may take place, including for copper(II) and cobalt(II) ions when combined with water molecules, ammonia molecules, hydroxide ions and chloride ions
- predict, with the help of E⦵ values, whether redox reactions involving transition elements and their ions are feasible
- describe the reactions of the following, and carry out calculations on them when suitable data are given: (a) MnO4–/C2O42– in acid solution (b) MnO4–/Fe2+ in acid solution (c) Cu2+ / I–
- use the data provided to carry out calculations on further redox systems
28.3Colour of complexes5 objectives
- define the terms degenerate and non-degenerate d orbitals, and apply both terms correctly
- outline how degenerate d orbitals split into two sets of non-degenerate d orbitals at higher energy, and apply ΔE to: (a) octahedral complexes, where there are two higher-energy d orbitals and three lower-energy ones (b) tetrahedral complexes, where three d orbitals lie higher and two lie lower
- explain why transition elements form coloured compounds, linking this to the frequency of the light absorbed when an electron is promoted from one non-degenerate d orbital to another
- describe qualitatively how changing the ligand alters ΔE, and so the frequency of light absorbed and the complementary colour seen
- use the complexes that copper(II) and cobalt(II) ions form with water and ammonia molecules and with hydroxide and chloride ions to illustrate how ligand exchange changes the observed colour
28.4Stereoisomerism in transition element complexes2 objectives
- describe the kinds of stereoisomerism that complexes can show, including those that arise with bidentate ligands: (a) cis/trans (geometrical) isomerism, e.g. square planar complexes such as [Pt(NH₃)₂Cl₂], and octahedral complexes like [Co(NH3)4(H2O)2]2+ and [Ni(H2NCH2CH2NH2)2(H2O)2]2+ (b) optical isomerism, shown for instance by [Ni(H2NCH2CH2NH2)3]2+ and [Ni(H2NCH2CH2NH2)2(H2O)2]2+
- deduce whether complexes such as those in 28.4.1(a) and 28.4.1(b) are polar overall
28.5Stability constants, Kstab4 objectives
- define Kstab, the stability constant of a complex, to be the equilibrium constant that applies when the complex ion forms in a solvent from the ions or molecules that make it up
- write a Kstab expression for a complex, with [H₂O] not included
- apply Kstab expressions to carry out calculations
- describe and explain ligand exchanges by reference to Kstab values, and understand that a large Kstab arises from the creation of a stable complex ion
29.1Formulas, functional groups and the naming of organic compounds4 objectives
- understand that each compound tabulated on page 47 has a functional group, and that this group governs its physical and chemical properties
- interpret the general, structural, displayed and skeletal formulas for every class of compound tabulated on page 47, and use them
- understand systematic nomenclature and use it to name simple aliphatic organic molecules, including cyclic compounds with one ring of no more than six carbon atoms, that carry the functional groups tabulated on page 47, with at most six carbon atoms (six plus six in esters and amides; for esters and nitriles, only straight chains)
- understand systematic nomenclature and use it for simple aromatic molecules that have one benzene ring and at least one simple substituent, for example 3-nitrobenzoic acid or 2,4,6-tribromophenol
29.2Characteristic organic reactions1 objective
- understand and apply these terms for classes of organic reaction mechanism: (a) electrophilic substitution (b) addition–elimination
29.3Shapes of aromatic organic molecules; σ and π bonds1 objective
- outline and account for the spatial arrangement that benzene and other aromatic molecules adopt, including sp² hybridisation, using σ bonds and a delocalised π system
29.4Isomerism: optical4 objectives
- understand that enantiomers share the same physical and chemical properties, except in how they rotate plane polarised light and in the biological activity they may have
- understand what optically active means and what a racemic mixture is, and use both terms
- describe how plane polarised light is affected by a substance's two optical isomers
- account for why chirality matters when drug molecules are made synthetically, including: how the two enantiomers may differ biologically, why a racemic mixture must be split into two pure enantiomers, and employing chiral catalysts to yield one pure optical isomer
30.1Arenes4 objectives
- describe the chemistry of arenes, illustrated through these processes involving benzene and methylbenzene: substitution by Cl₂ and by Br₂, with AlCl₃ or AlBr₃ acting as catalyst, to produce halogenoarenes (aryl halides); nitration by a mixture of the concentrated acids HNO₃ and H₂SO₄ at 25°C to 60°C; Friedel–Crafts alkylation with CH₃Cl and AlCl₃ under heating; Friedel–Crafts acylation with CH₃COCl and AlCl₃ under heating; complete oxidation of the side-chain with hot alkaline KMnO₄, followed by dilute acid, yielding a benzoic acid; hydrogenation, with H₂ and a Pt/Ni catalyst while heating, of the benzene ring to form a cyclohexane ring
- describe the mechanism of electrophilic substitution in arenes, taking the formation of nitrobenzene and of bromobenzene as examples, and use the effect of electron delocalisation in arenes (aromatic stabilisation) to explain why substitution predominates over addition
- predict, from the reaction conditions, whether an arene is halogenated in its side-chain or in its aromatic ring
- describe how, during electrophilic substitution of arenes, distinct substituent groups steer incoming groups to different positions on the ring (limited to the directing effects exerted by –NH₂, –OH, –R, –NO₂, –COOH and –COR)
31.1Halogen compounds2 objectives
- recall the routes that turn an arene into a halogenoarene through substitution, using Cl₂ or Br₂ together with a catalyst of AlCl₃ or AlBr₃, shown by benzene forming chlorobenzene and methylbenzene forming 2-chloromethylbenzene and 4-chloromethylbenzene
- account for why a halogenoalkane and a halogenoarene react differently, using chloroethane and chlorobenzene as the illustrative pair
32.1Alcohols1 objective
- describe how alcohols react with acyl chlorides to give esters, taking ethyl ethanoate as the example
32.2Phenol7 objectives
- state the routes (reagents and conditions) through which phenol can be made: treating phenylamine with HNO₂ or NaNO₂ plus dilute acid below 10°C to form the diazonium salt, then warming that diazonium salt with H₂O to give phenol
- recall phenol's behaviour through these examples: reaction with bases, for example NaOH(aq), giving sodium phenoxide; reaction with Na(s), giving sodium phenoxide and H₂(g); coupling with diazonium salts in NaOH(aq) to form azo compounds; nitrating the aromatic ring using dilute HNO₃(aq) at room temperature, producing a mixture of 2-nitrophenol and 4-nitrophenol; brominating the aromatic ring with Br₂(aq) to yield 2,4,6-tribromophenol
- account for the acidity shown by phenol
- describe and explain how water, phenol and ethanol compare in acidity
- account for why phenol needs different reagents and conditions for nitration and bromination than benzene does
- recall that, on a phenol's ring, substitution is directed by the hydroxyl group to the 2-, 4- and 6-positions
- apply your understanding of phenol's reactions to other phenolic compounds, e.g. naphthol
33.1Carboxylic acids5 objectives
- recall how benzoic acid can be made, namely by heating an alkylbenzene with alkaline KMnO₄ and subsequently treating the mixture with dilute acid, using methylbenzene as the example
- describe what happens when carboxylic acids are treated with PCl₃ and heated, or with PCl₅ or SOCl₂, to yield acyl chlorides
- recognise that certain carboxylic acids may undergo further oxidation: methanoic acid, HCOOH, is oxidised to carbon dioxide and water by Fehling's reagent, Tollens' reagent, acidified KMnO₄ or acidified K₂Cr₂O₇, while ethanedioic acid, HOOCCOOH, is oxidised to carbon dioxide by warm acidified KMnO₄
- outline and account for how the acidities of carboxylic acids, phenols and alcohols rank against one another
- describe and explain how chlorine-substituted carboxylic acids compare in acidity
33.2Esters1 objective
- recall the method for making esters from alcohols and acyl chlorides, with ethyl ethanoate and phenyl benzoate given as illustrations
33.3Acyl chlorides4 objectives
- state the methods (reagents and conditions) for making acyl chlorides: carboxylic acids heated with PCl₃, carboxylic acids treated with PCl₅, or carboxylic acids treated with SOCl₂
- describe the reactions that acyl chlorides undergo in the following cases: water added at room temperature, giving the carboxylic acid and HCl; an alcohol reacting at room temperature to make an ester and HCl; phenol reacting at room temperature to make an ester and HCl; ammonia reacting at room temperature to make an amide and HCl; a primary or secondary amine reacting at room temperature to make an amide and HCl
- describe how acyl chlorides undergo the addition–elimination mechanism in the reactions of 33.3.2(a)–(e)
- account for how comparatively easily hydrolysis occurs in acyl chlorides, alkyl chlorides and halogenoarenes (aryl chlorides)
34.1Primary and secondary amines3 objectives
- state the reagents and conditions needed to make primary and secondary amines in each of these cases: (a) halogenoalkanes heated with NH₃ in ethanol under pressure (b) halogenoalkanes heated with primary amines in ethanol in a sealed tube or under pressure (c) amides reduced by LiAlH₄ (d) nitriles reduced by LiAlH₄ or H₂/Ni
- describe how an amide forms when an acyl chloride and ammonia or an amine undergo condensation at room temperature
- outline and account for the basicity shown by aqueous solutions of amines
34.2Phenylamine and azo compounds4 objectives
- outline how phenylamine is made: benzene is nitrated to nitrobenzene, which is then reduced using hot Sn with concentrated HCl, and finally treated with NaOH(aq)
- outline: (a) what happens when phenylamine is combined with Br₂(aq) at room temperature (b) what happens when phenylamine is treated with HNO₂ or NaNO₂ and dilute acid below 10°C to yield the diazonium salt, and then how warming that diazonium salt with H₂O gives phenol
- outline and account for how aqueous ammonia, ethylamine and phenylamine rank against one another in base strength
- recall these points about azo compounds: (a) describe how benzenediazonium chloride couples with phenol in NaOH(aq) to give an azo compound (b) identify an azo group (c) state the frequent use of azo compounds as dyes (d) that further azo dyes may be made by a comparable route
34.3Amides3 objectives
- recall the reactions, along with their reagents and conditions, through which amides are formed: (a) ammonia reacting with an acyl chloride, kept at room temperature (b) a primary amine reacting with an acyl chloride, kept at room temperature
- describe the reactions of amides: (a) hydrolysis using either aqueous acid or aqueous alkali (b) LiAlH₄ reducing the CO group to give an amine
- state and explain why an amide is a much weaker base than an amine
34.4Amino acids3 objectives
- describe the acid/base properties of amino acids, along with how zwitterions come about, and make sure to cover the isoelectric point
- describe the way amide (peptide) bonds are generated between amino acids, giving di- and tripeptides
- interpret and predict the outcome of electrophoresis carried out at different pHs on a mixture containing amino acids and dipeptides (you will not be tested on setting up the apparatus)
35.1Condensation polymerisation4 objectives
- describe how polyesters are formed: (a) by reacting a diol with either a dicarboxylic acid or a dioyl chloride (b) by reacting a hydroxycarboxylic acid
- describe how polyamides are formed: (a) by reacting a diamine with either a dicarboxylic acid or a dioyl chloride (b) by reacting an aminocarboxylic acid (c) by joining amino acids together
- deduce the repeating unit for a condensation polymer produced from one specified monomer or from a specified pair of monomers
- identify the monomer(s) from which a given part of a condensation polymer chain was formed
35.2Predicting the type of polymerisation2 objectives
- predict whether a given monomer, or a given pair of monomers, will undergo addition or condensation polymerisation
- deduce which type of polymerisation reaction formed a given part of a polymer chain
35.3Degradable polymers3 objectives
- recognise that poly(alkenes) are chemically inert, which is why they may be hard to biodegrade
- recognise that a number of polymers may undergo degradation under the influence of light
- recognise that polyesters and polyamides are capable of biodegradation through hydrolysis under acidic and alkaline conditions
36.1Organic synthesis3 objectives
- where one organic molecule contains several functional groups: (a) identify those organic functional groups from the reactions covered in this course (b) predict the molecule's properties and reactions
- devise synthetic routes of more than one step to make organic molecules, drawing on the reactions covered in this course
- analyse a provided sequence of reactions, identifying for every stage the reaction category and the reagents involved, together with any side products that may form
37.1Thin-layer chromatography3 objectives
- describe each of the following terms and show that you understand it: (a) stationary phase, e.g. aluminium oxide coating a solid support (b) mobile phase, a solvent that can be either polar or non-polar (c) Rf value (d) baseline and solvent front
- interpret Rf values
- account for why Rf values differ, using the substance's interaction with the stationary phase and its relative solubility in the mobile phase
37.2Gas/liquid chromatography3 objectives
- describe and understand these terms: (a) stationary phase, a non-polar liquid with a high boiling point held on a solid support (b) mobile phase, which is a gas that does not react (c) retention time
- interpret gas /liquid chromatograms so as to express the makeup of a mixture as percentages
- explain retention times by considering how each substance interacts with the stationary phase
37.3Carbon-13 NMR spectroscopy2 objectives
- analyse and make sense of a carbon-13 NMR spectrum for a simple molecule so as to deduce: (a) the separate environments that the carbon atoms in it occupy (b) the structures that are possible for the molecule
- predict or explain how many peaks a given molecule gives in its carbon-13 NMR spectrum
37.4Proton (1H) NMR spectroscopy5 objectives
- analyse and interpret the proton (1H) NMR spectrum of a simple molecule in order to deduce: (a) which proton environments are present, from chemical shift values (b) how many of each type of proton there are, from relative peak areas (c) from the splitting pattern and the n + 1 rule (limited to singlet, doublet, triplet, quartet and multiplet), how many equivalent protons sit on the carbon atom next to the one carrying a given proton (d) which structures the molecule could have
- predict the chemical shift values and splitting patterns shown by the protons within a particular molecule
- describe how tetramethylsilane, TMS, is employed as the reference standard when measuring chemical shifts
- give the reason deuterated solvents, e.g. CDCl₃, are necessary when recording a proton NMR spectrum
- describe how O–H and N–H protons can be recognised by proton exchange with D₂O
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