Integrated Curriculum Pro Osed Syllabus Outline for 3 Topics

Sec 3 SIO for SMTP (Chemistry) 2011

Specific Instructional Objectives Chemistry – Sec 3 SMTP 2011

TOPIC 1 EXPERINMENTAL TECHNIQUES (Self-study)

1.1 Experimental design

(a) name appropriate apparatus for the measurement of time, temperature, mass and volume, including burettes, pipettes, measuring cylinders and gas syringes (b) suggest suitable apparatus, given relevant information, for a variety of simple experiments, including collection of gases and measurement of rates of reaction

1.2 Methods of purification and analysis

(a) describe methods of separation and purification for the components of the following types of mixtures: (i) solid-solid (ii) solid-liquid (iii) liquid-liquid (miscible and immiscible) Techniques to be covered for separation and purification include: (i) use of a suitable solvent, filtration and crystallization or evaporation (ii) sublimation (iii) distillation and fractional distillation (iv) use of a separating funnel (v) paper chromatography (b) describe paper chromatography and interpret chromatograms including comparison with ‘known’ samples and the use of Rf values (c) explain the need to use locating agents in the chromatography of colourless compounds (d) deduce from the given melting point and boiling point the identities of substances and their purity (e) explain that the measurement of purity in substances used in everyday life, e.g. foodstuffs and drugs, is important

1.3 Identification of ions and gases

(a) describe the use of aqueous sodium hydroxide and aqueous ammonia to identify the following aqueous cations: aluminium, ammonium, calcium, copper(II), iron(II), iron(III), lead(II) and zinc (formulae of complex ions are not required) (b) describe tests to identify the following anions: carbonate (by the addition of dilute acid and subsequent use of limewater); chloride (by reaction of an aqueous solution with nitric acid and aqueous silver nitrate); iodide (by reaction of an aqueous solution with nitric acid and aqueous lead(II) nitrate); nitrate (by reduction with aluminium and aqueous sodium hydroxide to ammonia and subsequent use of litmus paper) and sulfate (by reaction of an aqueous solution with nitric acid and aqueous barium nitrate) (c) describe tests to identify the following gases: ammonia (using damp red litmus paper); carbon dioxide (using limewater); chlorine (using damp litmus paper); hydrogen (using a burning splint); oxygen (using a glowing splint) and sulfur dioxide (using acidified potassium dichromate(VI))

* Refers to SMTP curriculum extensions 1 Sec 3 SIO for SMTP (Chemistry) 2011

TOPIC 2 ATOMIC STRUCTURE & BONDING

2.1 Atomic Structure

(a) identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses (b) deduce the behaviour of beams of protons, neutrons and electrons in both electric and magnetic fields (c) describe the distribution of mass and charges within an atom (d) define proton (atomic) number and nucleon (mass) number 12 (e) interpret and use such symbols as 6 C (f) deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge) (g) (i) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number (ii) define the term isotopes and distinguish between isotopes on the basis of different numbers of neutrons present (h) describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals (i) describe the shapes of s and p orbitals (j) state the electronic configuration of atoms and ions given the proton number (and charge) (k) i) *Explain the factors influencing the ionization energies of elements ii) * Explain the trends in ionization energies across a period and down a group of the Periodic Table (l) *Deduce the electronic configurations of elements from successive ionization energy data (m) * Interpret successive ionization energy data of an element in terms of the position of that element within the Periodic Table

2.2 Chemical bonding

(a) describe the formation of ions by electron loss/gain in order to obtain the electronic structure of a noble gas (b) describe the formation of ionic (electrovalent) bonds between metals and non-metals as the electrostatic force which holds two oppositely charged ions together., e.g. NaCl, MgCl2 and MgO, including the use of ‘dot and cross’ diagrams (c) state that ionic materials contain a giant lattice in which the ions are held by electrostatic attraction, e.g. NaCl (candidates will not be required to draw diagrams of ionic lattices or explain the effect of ionic charge and ionic radii on the numerical magnitude of a lattice energy) (d) deduce the formulae of other ionic compounds from diagrams of their lattice structures, limited to binary compounds (e) relate the physical properties (including electrical property) of ionic compounds to their lattice structure (f) describe the formation of a covalent bond by the sharing of a pair of electrons in order to gain the electronic configuration of a noble gas (g) describe, including the use of ‘dot and cross’ diagrams, the formation of covalent bonds between i) non-metallic elements e.g. H2; O2; N2; Cl2; HCl; CO2; CH4; C2H4 ii) *compounds with co-ordinate (dative covalent) bonding, as in BF3.NH3 (h) deduce the arrangement of electrons in other covalent molecules (i) *explain the shapes of, and bond angles in, molecules, such as BF3 (trigonal); CO2 (linear); CH4 (tetrahedral); NH3 (pyramidal); H2O( non-linear); SF6 (octahedral) by using the Valence Shell Electron Pair Repulsion theory (j) describe covalent bonding in terms of orbital overlap, giving  and  bonds (k) *explain the shape of, and bond angles in, the ethane, ethene and benzene molecules in relation to sigma and pi carbon-carbon bonds. (l) describe hydrogen bonding, using ammonia and water as examples of molecules containing – NH and –OH groups (m) outline the importance of hydrogen bonding to the physical properties including ice and water

(n) explain the terms bond energy, bond length and bond polarity and use them to compare the reactivities of covalent bonds (o) *describe intermolecular forces (van der Waals’ forces), based on permanent and induced dipoles, as in CHCl3 (l), Br2 (l) and the liquid noble gases (p) relate the physical properties (including electrical property) of covalent substances to their structure and bonding (q) describe metallic bonding in terms of a lattice of positive ions surrounded by mobile electrons (‘sea of electrons’) (r) relate the electrical conductivity of metals to the mobility of the electrons in the structure (s) describe, interpret and/or predict the effect of different types of bonding (ionic bonding; covalent bonding and metallic bonding) on the chemical and physical properties of substances (t) show understanding of chemical reactions in terms of energy transfer associated with the breaking and making of chemical bonds (u) compare the structure of simple molecular substances, e.g. methane, iodine, with those of giant molecular substances, e.g. poly(ethene); sand (silicon dioxide); diamond; graphite in order to deduce their properties (v) compare the bonding and structures of diamond and graphite in order to deduce their properties such as electrical conductivity, lubricating or cutting action (candidates will not be required to draw the structures) (w) suggest from quoted physical data the type of structure and bonding present in a substance (x) deduce the physical and chemical properties of substances from their structures and bonding

TOPIC 3 ACIDS BASES AND SALTS

3.1 The characteristic properties of acids and bases (a) state the properties of acids in reactions with metals (the ability of metals to react with acids should be linked to their position in the reactivity series), bases and carbonates (b) state the uses of sulfuric acid as in the manufacture of detergents and fertilisers; and as a battery acid (c) describe the characteristic properties of bases in reactions with acids and with ammonium salts (d) describe importance of water for acidity, i.e., water causes acid molecules to ionise and form hydrogen ions (e) define acid as a substance that produces hydrogen ions as the only positive ions in water (f) explain basicity of common acids and relate to concentration of hydrogen ions (g) describe qualitatively the difference between strong and weak acids in terms of extent of ionisation of acid in water (h) describe an alkali as a basic hydroxide which is soluble in water to produce hydroxide ions (i) describe the reaction between hydrogen ions and hydroxide ions to produce water, H + + OH-  H2O as neutralisation (j) describe aqueous ammonia as a weak alkali (k) identify aqueous sodium hydroxide and potassium hydroxide as strong alkalis (l) describe how to test hydrogen ion concentration and hence, relative acidity using Universal indicator and the pH scale (m) describe the importance of controlling pH in soil and how excess acidity can be treated using calcium hydroxide (n) classify oxides as either acidic, basic, amphoteric or neutral based on metallic / non-metallic character (o) classify sulfur dioxide as an acidic oxide and state its uses as a bleach, in the manufacture of wood pulp for paper and as a food preservative (by killing bacteria)

3.2 Preparation of salts

(a) define a salt as a substance formed when the hydrogen ions of an acid are partly or completely replaced by metallic or ammonium ions (b) describe the general rules of solubility for common salts to include nitrates, chlorides (including silver and lead), sulfates (including barium, calcium and lead), carbonates, hydroxides, Group I cations and ammonium salts (c) describe the techniques used in the preparation, separation and purification of salts as examples of the techniques specified in section 1.2(a) (methods for preparation should include precipitation and titration together with reactions of acids with metals, insoluble bases and insoluble carbonates) (d) suggest a method of preparing a given salt from suitable starting materials, given appropriate information

TOPIC 4 PERIODIC TABLE

4.1 Periodic Trends

(a) describe the Periodic Table as an arrangement of the elements in the order of increasing proton (atomic) number (b) describe how the position of an element in the Periodic Table is related to proton number and electronic structure (c) describe the relationship between Group number and the ionic charge of an element (d) explain the similarities between the elements in the same Group of the Periodic Table in terms of their electronic structure (e) describe the change from metallic to non-metallic character from left to right across a Period of the Periodic Table (f) predict the properties of elements in Group I, VII and the Transition elements using the Periodic Table (g) describe the relationship between Group number, number of valence electrons and metallic/non-metallic character. (h) *Describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements in the third period (sodium to argon)* i. Explain qualitatively the variation in atomic radius and ionic radius ii. Interpret the variation in melting point and in electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements iii. Explain the variation in first ionization energy

iv. Describe the reactions, if any, of the elements with oxygen (to give Na2O; MgO)

4.2 Group properties

(a) describe lithium, sodium and potassium in Group I (the alkali metals) as a collection of relatively soft, low density metals showing a trend in melting point and in their reaction with water (b) describe chlorine, bromine and iodine in Group VII (the halogens) as a collection of diatomic non-metals showing a trend in colour, state and their displacement reactions with solutions of other halide ions (c) describe the elements in Group 0 (the noble gases) as a collection of monatomic elements that are chemically unreactive and hence important in providing an inert atmosphere, e.g. argon and neon in light bulbs; helium in balloons; argon in the manufacture of steel (d) describe the lack of reactivity of the noble gases in terms of their electronic structure

4.3 Transition elements

(a) describe the central block of elements (transition metals) as metals having high melting points, high densities, variable oxidation state and forming coloured compounds (b) state the uses of these elements and/or their compounds as catalysts, e.g. iron in the Haber process; vanadium(V) oxide in the Contact process; nickel in the hydrogenation of alkenes

TOPIC 5 MOLE CONCEPT

(a) state the symbols of the elements and formulae of the compounds mentioned in the syllabus (b) deduce the formulae of simple compounds from the relative numbers of atoms present and vice versa (c) determine the formulae of ionic compounds from the charges on the ions present and vice versa (d) interpret chemical equations with state symbols (e) write and/or construct chemical equations, with state symbols, including ionic equations (f) define the terms relative atomic, isotopic, molecular and formula masses, based on the 12C scale (g) define the term mole in terms of the Avogadro constant (h) analyse mass spectra in terms of isotopic abundances and molecular fragments [knowledge of the working of the mass spectrometer is not required] (i) calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum (j) calculate the percentage mass of an element in a compound when given appropriate information (k) define the terms empirical and molecular formulae (l) calculate empirical formulae and molecular formulae from relevant data (using combustion data or composition by mass) (m) calculate stoichiometric reacting mass (from formulae and equations) and volumes of gases (e.g. in the burning of hydrocarbons) (one mole occupies 24 dm3 at room temperature and pressure); calculations involving the idea of limiting reactants may be set (The gas laws and the calculations of gaseous volumes at different temperatures and pressures are not required) (n) deduce stoichiometric relationships from calculations such as those in (m) (o) apply the concept of solution concentrations (in expressed in mol/dm3 or g/dm3) to process the results of volumetric experiments and to solve simple problems (appropriate guidance will be provided where unfamiliar reactions are involved) (p) calculate % yield and % purity

TOPIC 6 ATMOSPHERIC CHEMISTRY

(a) describe the volume composition of clean air in terms of 79% nitrogen, 20% oxygen, and the remainder being noble gases (with argon as main constituent) and carbon dioxide. (b) name some common atmospheric pollutants (carbon monoxide; methane; nitrogen oxides (NO and NO2); ozone; sulfur dioxide; unburnt hydrocarbons) (c) state sources of these pollutants as (i) carbon monoxide from the incomplete combustion of carbon-containing substances (ii) nitrogen oxides from lightning activity and internal combustion engines. (iii) sulfur dioxide from volcanoes and combustion of fossil fuels. (d) describe the reactions used in possible solutions to the problems arising from some of the pollutants in (b) (i) redox reactions in catalytic converters to remove combustion pollutants, (ii) the use of calcium carbonate to reduce the effects of 'acid rain and in flue gas desulfurisation. (e) discuss some of the effects of these pollutants on health and on the environment (i) the poisonous nature of carbon monoxide (ii) the role of nitrogen dioxide and sulfur dioxide in the formation of acid rain and its effects on respiration, buildings and plants. (iii) the role of nitrogen dioxide, methane and unburnt hydrocarbons in the formation of photochemical smog. (f) describe the importance of the ozone layer and the problems involved with the depletion of ozone by reaction with chlorine containing compounds, chlorofluorocarbons (CFCs). (g) describe the carbon cycle in simple terms to include (i) the process of combustion, respiration and photosynthesis (ii) how carbon cycle regulates the amount of carbon dioxide in the atmosphere. (h) state that carbon dioxide and methane are greenhouse gases and may contribute to global warming, give the sources of these gases and discuss the possible consequences of an increase in global warming.

Advanced Modules

TOPIC 7 *Metals

7.1 General Properties (a) Describe the general properties of metals and non-metals (b) Explain why metals are used in alloying & identify representations of metals and alloys from diagrams of structures

7.2 Reactivity series (a) Place metals in order of reactivity calcium, copper, (hydrogen) iron, magnesium, potassium, silver, sodium and zinc by reference to i) Reactions with water/ steam and with dilute hydrochloric acid. ii) Displacement reactions from a solution of their ions iii) Apparent unreactive nature of aluminium

* Refers to SMTP curriculum extensions 6