3AB+Chemistry

bonding capacity position on Periodic Table physical and chemical properties ||
 * = =Week= ||= =Content - 3A Chemistry= ||
 * = 1 - 2 || == Atomic structure and Periodic Table ==
 * explain the structure of the atom in terms of protons, neutrons and electrons
 * write the electron configuration using the shell model for the first twenty elements e.g. Na. 2, 8, 1
 * explain trends in ionisation energy, atomic radius and electronegativity across periods and down groups (for main group elements) in the Periodic Table
 * describe and explain the relationship between the number of valence electrons and an element’s
 * = 3 -6 || == Bonding ==


 * describe and apply the relationships between the physical properties and the structure of ionic, metallic, covalent network and covalent molecular substances
 * use the Valence Shell Electron Pair Repulsion (VSEPR) theory and Lewis structure diagrams to explain and predict and draw the shape of molecules and polyatomic ions (octet only)
 * explain polar and non-polar covalent bonds in terms of the electronegativity of the atoms involved in the bond formation
 * use the relationship between molecule shape and bond polarity to predict and explain the polarity of a molecule
 * explain the differences between intermolecular and intramolecular forces
 * describe and explain the origin and relative strength of the following intermolecular
 * interactions for molecules of a similar size:

dispersion forces dipole-dipole attractions hydrogen bonds ion-dipole interactions such as solvation of ions in aqueous solution


 * explain the relationships between physical properties such as melting and boiling point, and the types of intermolecular forces present in substances of similar size
 * apply an understanding of intermolecular interactions to explain the trends in melting and boiling points of hydrides of groups 15, 16 and 17 accounting for the anomalous behaviour of NH3, H2O and HF
 * explain and describe the interaction between solute and solvent particles in a solution
 * use the nature of the interactions, including the formation of ion-dipole and hydrogen bonds to explain water’s ability to dissolve ionic, polar and non-polar solutes. ||
 * = 7 - 9 || == Reactions, and equations ==


 * use observable properties, such as the colour of ions, to help predict and explain the formation of products in chemical processes
 * apply the solubility rules to predict if a precipitate will form when two dilute ionic solutions are mixed (see data sheet)
 * describe, write ionic equations for and interpret observations for:

precipitation reactions
 * write the chemical formulae for molecular compounds based on the number of atoms of each element present as inferred from the systematic names
 * write the molecular formulae of commonly encountered molecules that have non- systematic names

Stoichiometry
> conversion between Celsius and Kelvin temperature scales > mass, molar mass, number of moles of solute, concentration and volume of solution and gas volume using PV=nRT > percentage purity of reactants or percentage yield in industrial processes > a limiting reagent, including: > o identification of limiting reagents > o calculation of excess reagents. > the presence of catalysts > changes in temperature > pressure of whole system concentration state of subdivision > and reactants in a system in dynamic chemical equilibrium > following changes to a system initially at chemical equilibrium: changes in temperature changes in solution concentration changes in partial pressure of a gas addition of a catalyst * describe, write equations for > physical and chemical equilibrium > interpret observations, such as the colour changes, of physical and chemical > systems at equilibrium > > apply the concept of equilibrium in biological, environmental or laboratory situations > where a system is in dynamic chemical equilibrium > > explain the reasons for compromises between the ideal and actual conditions used > in industrial processes that involve reversible reactions ||
 * Use the Kinetic Theory to explain the concept of absolute zero.
 * perform calculations involving
 * perform concentration calculations (mol L-1,g L-1 ,ppm, percentage composition) ||
 * = 10-13 || Chemical equilibrium
 * explain by applying the collision theory how changes in rates of reactions can be accomplished by:
 * describe and explain the characteristics of a system in dynamic chemical equilibrium
 * write equilibrium law expressions for homogeneous and heterogeneous systems
 * use K and equilibrium law expression to explain the relative proportions of products
 * apply and explain how Le Châtelier’s principle can be used to predict the impact of the
 * = 14-15 ||= Semester 1 examination ||

> strong acid solutions > > strong base solutions > > the resulting solution when strong acid-base solutions are mixed > > explain using Le Châtelier’s Principle how buffers respond to the addition of H+ and OH- > > • describe, write equations for and interpret observations for the following reaction types: > neutralisation > > hydrolysis of salts of weak acids and weak bases ||
 * **Week** ||= ** Content - 3B Chemistry ** ||
 * 1 - 4 || Acids and bases in aqueous solutions
 * apply an understanding of the concept of an electrolyte to explain the self-ionisation of water
 * explain and apply the Arrhenius and Brønsted-Lowry models to describe acids and bases
 * apply the relationship between Kw and temperature to explain the pH value of a neutral solution at different temperatures
 * apply the relationship pH = - log H+ to calculate the pH of: (aq)
 * describe, write equations for and interpret observations for the solvation of ions in aqueous solution
 * apply the Brønsted-Lowry model to the hydrolysis of salts to predict and explain the acidic, basic or neutral nature of salts derived from monoprotic and polyprotic acids, and bases
 * describe and explain the conjugate nature of buffer solutions
 * explain qualitatively the concept of buffering capacity.
 * Reactions, and equations
 * 5 - 6 || Chemical reactions Stoichiometry

perform the calculation of concentration and volume involved in the dilution of solutions and the addition of solutions

calculate the concentration of ions in solution (mol L-1) for strong electrolytes perform volumetric analysis using either acid-base or redox context, and: > error > solutions > of an appropriate indicator in an acid-base titration > perform calculations based on acid-base and redox titrations || O, Cl , MnO –, Cr O 2–, ClO–, H+, concentrated sulfuric acid, concentrated nitric acid and common reducing agents (reductants) including Zn, C, H, Fe2+, C O 2– 224 > anode processes > > cathode processes > > electrolyte > > salt bridge and ion migration > > electron flow in external circuit > electric current >> Reactions and equations • describe, write equations for and interpret observations for the oxidation and reduction equations in an acidic environment || substitution reactions of alkanes
 * give a description of procedures used and methods for minimising experimental
 * describe and explain the characteristics of primary standards and standard
 * demonstrate an understanding of end point and equivalence point in the selection
 * explain the choice of indicators (in acid-base only) or use of self-indicators (redox)
 * 7 - 8 || Oxidation and reduction
 * apply the table of Standard Reduction Potentials to determine the relative strength of oxidising and reducing agents to predict reaction tendency
 * apply oxidation numbers to identify redox equations and/or oxidants and reductants identify by name and/or formula common oxidising and reducing agents including
 * describe and explain the role of the following in the operation of an electrochemical (galvanic) cell:
 * describe the electrical potential of a galvanic cell as the ability of a cell to produce an
 * describe and explain how an electrochemical cell can be considered as two half-cells
 * describe the role of the hydrogen half-cell in the table of Standard Reduction Potentials
 * describe the limitations of Standard Reduction Potentials table.
 * 9 - 12 || Organic chemistry
 * write balanced equations for the following reactions of hydrocarbons:

addition reactions of alkenes

combustion >> solubility in polar and non-polar solvents in terms of the intermolecular interactions >> reactions of alcohols:
 * draw and name structural isomers of alkanes and structural and geometric isomers of alkenes
 * recognise the functional groups—alcohols, aldehydes, ketones, carboxylic acids and esters and name simple straight chain examples to C8
 * explain the relationship between the presence of a functional group and chemical behaviour
 * alcohols:
 * name simple straight chain examples to C8
 * draw simple structural formula for primary, secondary and tertiary alcohols
 * explain physical properties of alcohols such as melting and boiling points and
 * describe, write equations for and predict and interpret observations for the following

with carboxylic acids with acidified Cr O 2- and MnO to produce: - aldehydes

- ketones

- carboxylic acids > > recognise primary amines > > name and draw simple structural formulae for primary amines only > >> amino acids and trans-fatty acids >> plastics >> organic compound from the analysis of combustion or other data ||
 * amines:
 * α amino acids:
 * recognise general structural formula for α amino acid
 * describe the chemistry of common organic substances such as soaps, detergents,
 * apply and explain the concept of polymerisation such as polypeptides, silicones or
 * determine by calculation the empirical and molecular formulae and the structure of an
 * 13 - 14 ||= ** Semester 2 examination ** ||
 * 15 ||= ** Revision program ** ||

**Exploring Chemistry Stage 3** **Problem solving and quantities in chemistry**


 * = Term ||= Week ||= Strand ||= Sets ||< Title ||
 * = 1 ||= 1 || Properties and Reactions ||= 2, 5 ||< * Limiting reagents
 * Percentage composition and yield ||
 * =  ||= 2 ||   ||= 1, 6 ||< * Gases
 * Empirical Formulae ||
 * =  ||= 3 || Atomic Structure and Bonding ||= 8 ||< * Electron configuration ||
 * =  ||= 4 ||   ||= 9 ||< * Ionisation energy ||
 * =  ||= 5 ||   ||= 10 ||< * Shape and polarity ||
 * =  ||= 6 ||   ||= 11 ||< * Periodic trends ||
 * =  ||= 7 || Properties and Reactions ||= 7 ||< * Reaction types ||
 * =  ||= 8 ||   ||= 3 ||< * Equilibrium constant expressions ||
 * =  ||= 9 ||   ||= 4 ||< * Application of Le Chatelier's principle ||
 * =  ||= 10 ||   ||=   ||<   ||
 * = 2 ||= 1 || Acids and Bases ||= 12 ||< * Acid and base strength ||
 * =  ||= 2 ||   ||= 13 ||< * Hydrolysis ||
 * =  ||= 3 ||   ||= 14 ||< * Water equilibrium ||
 * =  ||= 4 ||   ||= 15 ||< * Indicators and their use ||
 * =  ||= 5 ||   ||= 16 ||< * The pH scale ||
 * =  ||= 6 ||   ||= 17 ||< * Buffers ||
 * =  ||= 7 ||   ||= 18 ||< * Acid - base titrations 1 ||
 * =  ||= 8 ||   ||= 19 ||< * Acid - base titrations 2 ||
 * =  ||= 9 || Oxidation and Reduction ||= 20 ||< * Oxidation and reduction ||
 * =  ||= 10 ||   ||= 21, 22 ||< * Balancing redox equations
 * Redox titrations ||
 * = 3 ||= 1 ||  ||= 23 ||< * Electrochemistry ||
 * =  ||= 2 || Organic Chemistry ||= 24 ||< * Organic compunds ||
 * =  ||= 3 ||   ||= 25 ||< * Reactions of organic compounds ||
 * =  ||= 4 ||   ||= 26 ||< * Calculations involving carbon compounds ||
 * =  ||= 5 ||   ||= 27 ||< * Empirical and molecular formula from analysis ||
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