1.3 b) perform calculations, including. 1.4(i) reacting masses (from formulae. candidates answers should reflect the

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1 SUBJECT: CHEMISTRY YEAR: 12 Study PLAN st Term TEACHER: MS. MASOOMA MON TH DATE WEE K Sep 2 nd 1 LEARNING OBJECIIVES PRIOR LEARNING RESOURCES 1 st Term Commences U1: Moles and equation Determine the formula 1.1define and use the terms relative of an ionic compound atomic, isotopic, molecular and from the charges on the earn- formula masses, based on the 12 C ions present chemistry/resource/r scale Construct and use es /the- 1.2 a) write and construct balanced symbolic equations with determination-of- equations state symbols, including relative-atomic-mass 1.3 b) perform calculations, including ionic equations Various worksheets Sep 9 th 2 use of the mole concept, involving: 1.4(i) reacting masses (from formulae Define relative molecular mass, Mr, as on calculations: and equations) the sum of the relative uk/ 1.5(ii) volumes of gases (e.g. in the atomic masses (relative Titrating sodium burning of hydrocarbons) formula mass or Mr will hydroxide with (iii) volumes and concentrations of be used for ionic hydrochloric acid: solutions compounds When performing calculations, Define the mole in terms stry.org/experiments candidates answers should reflect the of a specific number of /titrating-sodium- number of significant figures given or particles called hydroxide-with-

2 asked for in the question. When Avogadro s constant hydrochloric- rounding up or down, candidates Use the molar gas acid,129,ex.html should ensure that significant figures volume, taken as 24 dm 3 are neither lost unnecessarily nor used at room temperature and beyond what is justified (see also pressure Practical Assessment, Paper 3, Display Calculate stoichiometric of calculation and reasoning on reacting masses and page 52) reacting volumes of c) deduce stoichiometric relationships solutions; solution from calculations such as those in concentrations will be 1.5(b) expressed in mol/dm 3 U2: Atomic structure Describe the structure of a) identify and describe protons, an atom in terms of neutrons and electrons in terms of their electrons and a nucleus uk (atoms) relative charges and relative masses containing protons and b) deduce the behavior of beams of neutrons watch?v=2afpfg0c protons, neutrons and electrons in State the relative omohttp:// Sep 16 th 3 electric fields c) describe the distribution of mass and charges and approximate relative chemistry.com/links/ AtomicStructure/Pau charge within an atom masses of protons, lihundsrule.htm d) deduce the numbers of protons, neutrons and electron neutrons and electrons present in both Describe the build-up Past papers atoms and ions given proton and of electrons in shells and Paper 11, Nov 2011, nucleon numbers and charge. understand the Q3 significance of the noble Paper 11, Nov 2012,

3 gas electronic structures Q2 2.1 Electrons in atoms and of valency electron Paper 11, June 2012, a) describe the number and relative Q1 energies of the s, p and d orbitals for Paper 12, June 2012, the principal quantum numbers 1, 2 and Q4 3 and also the 4s and 4p orbitals Paper 12, June 2013, b) describe and sketch the shapes of s Q31 and p orbitals Paper 13, June 2012, 2.2c) state the electronic configuration Q2 of atoms and ions given the proton Paper 13, June 2013, number and charge, using the Q32 convention 1s 2 2s 2 2p 6, etc. 2.3 Electrons in atoms d) (i) explain and use the term ionisation energy (ii) explain the factors influencing the ionisation energies of elements Sep 23 rd 4 (iii) explain the trends in ionisation energies across a Period and down a Group of the Periodic Table (see also Describe the formation of ions by electron loss or gain com/watch?v=ewf7 RlVNBSA Section 9.1) e) deduce the electronic configurations of elements from successive ionisation energy data f) interpret successive ionisation energy

4 data of an element in terms of the position of that element within the Periodic Table g) explain and use the term electron affinity Oct 30 5 U4: States of matter 4.1 a) state the basic assumptions of the kinetic theory as applied to an ideal gas 4.2 b) explain qualitatively in terms of intermolecular forces and molecular size: (i) the conditions necessary for a gas to approach ideal behaviour (ii) the limitations of ideality at very high pressures and very low temperature 4.3 c) state and use the general gas equation pv = nrt in calculations, including the determination of Mr Q2 (a) These are a collection of interactive Java applets. Q9 Paper 13, Nov 2013, Q7 Q2 (b) U5: Enthalpy changes Relate the terms, Practical Booklet 4 Oct a) explain that chemical reactions are accompanied by energy changes, exothermic and endothermic to the Textbooks Lainchbury principally in the form of heat energy; temperature changes Experiments 2.1, 2.2

5 the energy changes can be exothermic (ΔH is negative) or endothermic (ΔH is positive) 5.2 b) explain and use the terms: (i) enthalpy change of reaction and standard conditions, with particular reference to: formation, combustion, hydration, solution, neutralization, atomization (ii) bond energy (ΔH positive, i.e. bond breaking) (iii) lattice energy (ΔH negative, i.e. gaseous ions to solid lattice) 5.3 c) calculate enthalpy changes from appropriate experimental results, including the use of the relationship enthalpy change, ΔH = mcδt 5.4 d) explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy e) apply Hess Law to construct simple energy cycles, and carry out calculations involving such cycles and relevant energy terms, with particular reference to: during chemical reactions Understand that exothermic and endothermic changes relate to the transformation of chemical energy to heat (thermal energy), and vice versa Q2 (c)(i) Paper 21, Nov 2013, Q5 (c) Paper 23, Nov 2013, Q2 (c) Paper 41, Nov 2013, Q4 (a) Q3 (a) Paper 23, Nov 2013, Q2 (d) Lainchbury Experiment 2.3 Q12 Paper 21, June 2013, Q2 (c) Q3 (c)(i)

6 (i) determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion (ii) average bond energies (iii) the formation of a simple ionic solid and of its aqueous solution (iv) Born-Haber cycles (including ionization energy and electron affinity) U6: Redox reaction 6.1 a) calculate oxidation numbers of Practical Booklet 4, elements in compounds and ions 5 Oct ,2 b) describe and explain redox processes in terms of electron transfer and changes in oxidation number Patterns in chemical behaviour and reactions Chemical bonds Q1 Paper 13, Nov 2013, 6.3 c) use changes in oxidation Q1 numbers to help balance chemical equations U8: Kinetics Describe the effect of Practical Booklet a) explain and use the term rate of concentration, particle Oct 21 8 reaction 8.2 b) explain qualitatively, in terms of collisions, the effect of concentration size, catalysis and temperature on the speeds of reactions Q8 Paper 13, Nov 2013, changes on the rate of a reaction Describe and explain the Q5 8.3 c) explain and use the terms rate effects of temperature

7 equation, order of reaction, rate and concentration in Q7 constant, half-life of a reaction, rate- terms of collisions Paper 42, June 2013, determining step between reacting Q1 (c) d) construct and interpret a reaction particles Textbooks pathway diagram, in terms of the Define catalyst as an Lainchbury enthalpy change of the reaction and of agent which increases Experiment 9.7 the activation energy rate but which remains e) explain and use the term catalysis, unchanged catalysts can be homogenous or uk heterogeneous: (i) explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower activation energy (ii) interpret this catalytic effect in terms of the Boltzmann distribution U9: Periodicity Distinguish between Group ii and vii metals and non-metals a) describe the reactions of the by their general physical earn- Oct/N ov 28 9 elements with oxygen, water and dilute acids 9.2 b) describe the behaviour of the and chemical properties Describe the reactivity series as related to the chemistry/resource/r es /thereaction-of- oxides, hydroxides and carbonates with tendency of a metal to magnesium-with- water and dilute acids form its positive ion, steam 9.3c) describe the thermal illustrated by its

8 decomposition of the nitrates and carbonates d) interpret, and make predictions from, the trends in physical and chemical properties of the elements and their compounds e) state the variation in the solubilities of the hydroxides and sulfates U13: Nitrogen and sulphur 13.1 a) explain the lack of reactivity of nitrogen b) describe and explain: (i) the basicity of ammonia (see also Section 7.2) (ii) the structure of the ammonium ion reaction, if any, with the aqueous ions of other listed metals the oxides of the other listed metals earnchemistry/resource/r es /theperiodic-tableproperties-of-group- 2-elements Paper 41, Nov 2013, Q1 (b), Paper 43, Nov 2013, Q1 (b)(i) Paper 41, June 2013, Q4 (a) Q19 Paper 13, Nov 2013, Q15 Nov 4 10 and its formation by an acid-base reaction Paper 21, June 2013, Q1 (c) (iii) the displacement of ammonia from Paper 21, June 2013, its salts Q1 (b) a) describe the formation of atmospheric sulphur dioxide from the combustion of sulphur-contaminated fossil fuels 13.2 b) state the role of sulphur dioxide

9 in the formation of acid rain and describe the main environmental consequences of acid rain U3: Chemical bonding 3.1 a) describe, including the use of dot-and-cross diagrams: (i) covalent bonding, in molecules such as hydrogen, oxygen, chlorine, hydrogen chloride, carbon dioxide, Paper 21, Nov 2013, Nov methane, ethene (ii) co-ordinate (dative covalent) bonding, such as in the formation of the ammonium ion and in the Al 2 Cl 6 molecule 3.2 b) describe covalent bonding in terms of orbital overlap, giving σ and π bonds, including the concept of hybridisation to form sp, sp 2 and sp 3 Describe the formation of ionic bonds between elements from Groups I and VII Explain the formation of ionic bonds between metallic and nonmetallic elements Q1 (b) Paper 23, Nov 2013, Q1 (a) & Q3 (b)(ii) Paper 41, Nov 2013, Q1 (a) Q1 (a) Q33 orbitals (see also Section 14.3) Paper 21, Nov 2013, 3.3 c) explain the shapes of, and bond Q1 (a) angles in, molecules by using the qualitative model of electron-pair repulsion (including lone pairs), using as simple examples: BF 3 (trigonal), CO 2 (linear), CH 4 (tetrahedral), NH 3

10 (pyramidal), H 2 O (non-linear), SF 6 (octahedral), PF 5 (trigonal bipyramidal) 3.4 & 3.5d) predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.2(b) (see also Section 14.3) Nov 18 th 12 Revision Nov Revision Dec 2nd 15 First Term Exam. Dec 9 th 16 First Term Exam. Dec Winter Break