Chemistry_Assessment_Terms

Topic 1: Quantitative Chemistry 1.1.1 Apply the mole concept to substances The mole concept applies to all kinds of particles: atoms, molecules, ions, electrons, formula units, and son on. The amount of substance is measured in moles. Avogadro’s constant: 6.02 * 10^23 mol-1 1.1.2 Determine the number of particles and the amount of substance (in moles) Convert between the amount of stance (in moles) and the number of atoms, molecules, ions, electrons and formula units. 1.2 Formulas 1.2.1 Define the terms “relative atomic mass” (Ar) and “relative molecular mass” (Mr) Relative Atomic Mass: Mass of one mole of atoms Relative Molecular Mass: Mass of one mole of a compound 1.2.2 Calculate the mass of one mole of a species from its formula 1.2.3 Solve problems involving the relationship between the amount of substance in moles, mass and molar mass. 1.2.4 Distinguish between the terms “empirical formula” and “molecular formula” Chemical formula: the number of each type of atom in the smallest viable unit of the substance. The chemical formula will be an integral number of the empirical formula Empirical formula: the simplest possible ratio of elements in a substance. 1.2.5 Determine the empirical formula from the percentage composition or from other experimental data. 1.2.6 Determine the molecular formula when given both the empirical formula and experimental data. 1.3 Chemical Equations 1.3.1 Deduce chemical equations when all reactants and products are given. 1.3.2 Identify the mole ratio of any two species in a chemical equation 1.3.3 Apply the state symbols (s), (l), (g) and (aq) 1.4 Mass and Gaseous Volume Relationships in Chemical Reactions 1.4.1 Calculate theoretical yields from chemical equations. Given a chemical equation and the mass or amount (in moles) of one species, calculate the mass or amount of another species. 1.4.2 Determine the limiting reactant and the reactant in excess when quantities or reacting substances are given. 1.4.3 Solve problems involving theoretical, experimental, and percentage yield. 1.4.4 Apply Avogadro’s law to calculate reacting volumes of gases. 1.4.5 Apply the concept of molar volume at standard temperature and pressure in calculations The molar volume of an ideal gas under standard conditions is 2.24 x 10-2 m3 mol-1 (22.4 dm3mol-1) 1.4.6 Solve problems involving the relationship between temperature, pressure and volume for a fixed mass of an ideal gas. 1.4.7 Solve problems using the ideal gas equation, PV = nRT 1.4.8 analyse graphs relating to the ideal gas equation 1.5 Solutions 1.5.1 Distinguish between the terms “solute,” “solvent,” “solution,” and “concentration” Concentration in mol dm-1 is often represented by square brackets around the substance under consideration, for example, [HCl] 1.5.2 Solve problems involving concentration, amount of solute and volume of solution Topic 2: Atomic Structure 2.1 The Atom 2.1.1 State the position of protons, neutrons, and electrons in the atom 2.1.2 State the relative masses and relative charges of protons, neutrons, and electrons Proton relative mass: 1 Relative charge: +1 Neutron relative mass: 1 Relative charge: 0 Electron relative mass: 5*10-4 Relative charge: -1 2.1.3 Define the terms “mass number” (A), “atomic number” (Z), and “isotopes of an element” 2.1.4 Deduce the symbol for an isotope given its mass number and atomic number 126C using the formula AZX 2.1.5 Calculate the number of protons, neutrons, and electrons in atoms and ions from the mass number, atomic number and charge. 2.1.6 Compare the properties of the isotopes of an element 2.1.7 Discuss the uses of radioisotopes Examples should include 14C in radiocarbon dating, 60Co in radiotherapy, and 131I and 125I as medical tracers 2.2 The Mass Spectrometer 2.2.1 Describe and explain the operation of a mass spectrometer A simple diagram of a single beam mass spectrometer is required. The following stages of operation should be considered: vaporization, ionization, acceleration, deflection and detection 2.2.2 Describe how the mass spectrometer may be used to determine relative atomic mass using the 12C scale 2.2.3 Calculate non-integer relative atomic masses and abundance of isotopes from given data. 2.3 Electron Arrangement 2.3.1 Descrie the electromagnetic spectrum Identify ultraviolet, visible, and infrared regions, and describe the variation in wavelength, frequency, and energy across the spectrum. 2.3.2 Distinguish between a “continuous spectrum” and a “line spectrum” 2.3.3 Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels Be able to draw an energy level diagram, show transitions between different energy levels and recognize that lines in a line spectrum are directly related to these differences. An understanding of convergence is expected. Series should be considered in the ultraviolet, visible, and infrared regions of the spectrum. 2.3.4 Deduce the elctron arrangement for atoms and ions up to Z = 20 for example, 2.8.7 for Z = 17 Topic 3: Periodicity 3.1 The periodic table 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. 3.1.2 Distinguish between the terms “group” and “period” Be aware of the position of the transition elements in the periodic table 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table 3.2 Physical Properties 3.2.1 Define the terms “first ionization energy” and “electronegativity” 3.2.2 Descrie and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (LiCs) and the halogens (FI). 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, and electronegativities for elements across period three. 3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table. 3.3 Chemical Properties 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. The following reactions should be covered: * Alkali metals (Li, Na, and K) with water * Alkali metals (Li, Na, and K) with halogens (Cl2, Br2, and I2) * Halogens (Cl2, Br2, and I2) with halide ions (Cl-, Cr-, and I-) 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3 Equations are required for the reactions of Na2O, MgO, P4O10 and SO3 with water. Topic 4: Bonding 4.1 Ionic Bonding 4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions 4.1.2 Describe how ions can be formed as a result of electron transfer 4.1.3 Deduce which ions will be formed when elements in groups 1, 2, and 3 lose electrons 4.1.4 Deduce which ions will be formed when elements in groups 5, 6, and 7 gain electrons 4.1.5 State that transition elements can form more than one ion Include examples such as Fe2+ and Fe3+ 4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values 4.1.7 State the formula of common polyatomic ions formed by non-metals in periods 2 and 3 Examples include NO-3, OH-, SO2-4, CO2-3, PO43-, NH4+, HCO3-. 4.1.8 Describe the lattice structure of ionic compounds Students should be able to describe the structure of sodium chloride as an example of an ionic lattice. 4.2 Covalent Bonding 4.2.1 Descrie the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei Single and multiple bonds should be considered. Examples should include O2, N2, CO2, HCN, C2H4 (ethene) and C2H2 (ethyne) 4.2.2 Describe how the covalent bond is formed as a result of electron sharing Dative covalent bonds are required. Examples include CO, NH4+, and H3O+ 4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom. 4.2.4 State and explain the relationship between the number of bonds, bond length, and bond strength The comparison should include the bond lengths and bond strengths of: * Two carbon atoms joined by single, double and triple bonds * The carbon atom and the two oxygen atoms in the carboxyl groups of a carboxylic acid 4.2.5 Predict whether a compound of two elements would be covalent from the position of the elements in the periodic table or from their electronegativity values 4.2.6 Predict the relative polarity of bonds from electronegativity values 4.2.7 Predict the shape and bond angles for species with four, three and two negative charge centers on the central atom using the valence shell electron pair repulsion theory (VSEPR) Examples should include CH4, NH3, H2O, NH4+, H3O+, BF3, C2H4, SO2, C2H2, and CO2 4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities. 4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite, and C60 fullerene). 4.2.10 Describe the structure of and bonding in silicon and silicon dioxide 4.3 Intermolecular Forces 4.3.1 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles, or hydrogen bonding) and explain how they arise from the structural features of molecules The term van der Waals’ forces can be used to describe the interaction between non-polar molecules 4.3.2 Describe and explain how intermolecular forces affect the boiling points of substances The presence of hydrogen bonding can be illustrated by comparing: * HF and HCl * H2O and H2S * NH3 and PH3 * CH3OCH3 and CH3CH2OH * CH3CH2CH3, CH3CHO, and CH3CH2OH 4.4 Metallic Bonding 4.4.1 Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons. 4.4.2 Explain the electrical conductivity and malleability of metals 4.5 Physical Properties 4.5.1 Compare and explain the properties of substances resulting from different types of bonding Examples should include melting and boiling points, volatility, electrical conductivity, and solubility in non=polar and polar solvents Topic 5: Energetics 5.1 Exothermic and Endothermic Reactions 5.1.1 Define the terms “exothermic reaction,” “endothermic reaction,” and “standard enthalpy change of reaction” (∆H˚) Standard enthalpy change is the heat energy transferred understandard conditions – pressure 101.3 kPa, temperature 298 K. Only ∆H can be measured, not H for the initial or final state of a system 5.1.2 State the combustion and neutralization are exothermic processes 5.1.3 Apply the relationship between temperature change, enthalpy change and the classification of a reaction as endothermic or exothermic 5.1.4 Deduce, from an enthalpy level diagram, the relative stabilities of reactants and products, and the sign of the enthalpy change for the reaction 5.2 Calculation of Enthalpy Changes 5.2.1 Calculate the heat energy change when the temperature of a pure substance is changed Be able to calculate the heat energy change for a substance given the mass, specific heat capacity and temperature change using q = mc∆t 5.2.2 Design suitable experimental procedures for measuring the heat energy changes of reactions Should consider reactions in aqueous solution and combustion reactions. 5.2.3 Calculate the enthalpy change for a reaction using experimental data on temperature change, quantities of reactants and mass of water 5.2.4 Evaluate the results of experiments to determine enthalpy changes Be aware of the assumptions made and errors due to heat loss 5.3 Hess’s Law 5.3.1 Determine the enthalpy change of a reaction that is the sum of two or three reactions with known enthalpy changes Students should be able to use simple enthalpy cycles and enthalpy level diagrams and to manipulate equations. Not required to state Hess’s Law. 5.4 Bond Enthalpies 5.4.1 Define the term “average bond enthalpy” 5.4.2 Explain, in terms of average bond enthalpies, why some reactions are exoxthermic and others are endothermic Topic 6: Kinetics 6.1 Rates of Reaction 6.1.1 Define the term “rate of reaction” 6.1.2 Describe the suitable experimental procedures for measuring rates of reactions 6.1.3 Analyze data from rate experiments Be familiar with graphs of changes in concentration, volume, and mass against time 6.2 Collision Theory 6.2.1 Describe the kinetic theory in terms of the movement of particles whose average energy is proportional to temperature in kelvins 6.2.2 Define the term “activation energy” Ea 6.2.3 Describe the collision theory Know that reaction rate depends on: * Collision frequency * Number of particles with E ≥ Ea * Appropriate collision geometry or orientation 6.2.4 Predict and explain, using the collision theory, the qualitative effects of particle size, temperature, concentration and pressure on the rate of a reaction 6.2.5 Sketch and explain qualitatively the Maxwell-Boltzmann energy distribution curve for a fixed amount of gas at different temperatures and its consequences for changes in reaction rate 6.2.6 Describe the effect of a catalyst on a chemical reaction 6.2.7 Sketch and explain Maxwell-Boltzmann curves for reactions with and without catalysts Topic 7: Equilibrium 7.1 Dynamic Equilibrium 7.1.1 Outline the characteristics of chemical and physical systems in a state of equilibrium 7.2 The Position of Equilibrium 7.2.1 Deduce the equilibrium constant expression (Kc) from the equation for a homogeneous reaction Consider gases, liquids and aqueous solutions 7.2.2 Deduce the extent of a reaction from the magnitude of the equilibrium When Kc>>1, the reaction goes almost to completion When Kc<<1, the reaction hardly proceeds 7.2.3 Apply Le Chatelier’s principle to redict the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium and on the value of the equilibrium constant Not required to state Le Chatelier’s principle 7.2.4 State and explain the effect of a catalyst on an equilibrium reaction 7.2.5 Apply the concepts of kinetics and equilibrium to industrial proceses Haber and Contact processes Topic 8: Acids and Bases 8.1 Theories of Acids and Bases 8.1.1 Define “acids” and “bases” according to the Brønsted-Lowry and Lewis theories 8.1.2 Deduce whether or not a species could act as a Brønsted-Lowry and/or a Lewis acid or base 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brønsted-Lowry base (or acid) Make clear location of proton transferred, for example CH3COOH/CH3COO- rather than C2H4O2/C2H3O2- 8.2 Properties of Acids and Bases 8.2.1 Outline the characteristic properties of acids and bases in aqueous solution Bases that are not hydroxides, such as ammonia, soluble carbonates and hydrogencarbonates, should be included. Alkalis are bases that dissolve in water. Consider effects on indicators and the reactions of acids with bases, metals, and carbonates. 8.3 Strong and Weak Acids and Bases 8.3.1 Distinguish between “strong” and “weak” acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity 8.3.2 State whether a given acid or base is strong or weak Consider hydrochloric acid, nitric acid, and sulfuric acid as strong acids Carboxylic acids and carbonic acid (aqueous carbon dioxide) as weak acids Consider all group 1 hydroxides and barium hydroxide as strong bases, and ammonia and amines as weak bases 8.3.3 Distinguish between “strong” and “weak” acids and bases, and determine the relative strengths of acids and bases, using experimental data. 8.4 The pH Scale 8.4.1 Distinguish between aqueous solutions that are “acidic,” “neutral,” or “alkaline” using the pH scale 8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline using pH values Be familiar with a universal indicator 8.4.3 State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+ (aq)] expressed as powers of 10 Topic 9: Oxidation and Reduction 9.1 Introduction to Oxidation and Reduction 9.1.1 Define “oxidation” and “reduction” in terms of electron loss and gain. 9.1.2 Deduce the oxidation number of an element in a compount Oxidation numbers should be shown by a sign (+ or -) and a number, for example +7 for Mn in KMnO4 9.1.3 State the names of compounds using oxidation numbers Oxidation numbers in names of comounds are represented by Roman numerals, for example iron (II) oxide, iron (III) oxide 9.1.4 Deduce whether an element undergoes oxidation or reduction in reactions using oxidation numbers 9.2 Redox Equations 9.2.1 Deduce simple oxidation and reduction half-equations given the species involved in a redox reaction 9.2.2 Deduce redox equations using half-equations H+ and H2O should be used where necessary to balance half-equations in acid solution. The balancing of equations for reactions in alkaline solution will not be assessed. 9.2.3 Define the terms “oxidizing agent” and “reducing agent” 9.2.4 Identify the oxidizing and reducing agents in redox equations 9.3 Reactivity 9.3.1 Deduce a reactivity series based on the chemical behavior of a group of oxidizing and reducing agents Examples include displacement reactions of metals and halogens. Standard electrode potentials will not be assessed. 9.3.2 Deduce the feasibility of a redox reaction from a given reactivity series 9.4 Voltaic Cells 9.4.1 Explain how a redox reaction is used to produce electricity in a voltaic cell Should include diagram to show how two half-cells can be connected by a salt bridge. Examples of half-cells are Mg, Zn, Fe and Cu in solutions of their ions 9.4.2 State that oxidation occurs at the negative electrode (anode) and reduction occurs at the positive electrode (cathode) 9.5 Electrolytic Cells 9.5.1 Describe, using a diagram, the essential components of an electrolytic cell The diagram should include the source of electric current and conductors, positive and negative electrodes, and the electrolyte 9.5.2 State that oxidation occurs at the positive electrode (anode) and reduction occurs at the negative electrode (cathode). 9.5.3 Describe how current is conducted in an electrolytic cell 9.5.4 Deduce the products of the electrolysis of a molten salt Half-equations showing the formation of products at each electrode will be assessed Topic 10: Organic Chemistry 10.1 Introduction 10.1.1 Describe the features of a homologous series * Same general formula * Neighboring members differing by CH2 * Similar chemical properties * Gradation in physical properties 10.1.2 Predict and explain the trends in boiling points of members of a homologous series 10.1.3 Distinguish between “empirical,” “molecular,” and “structural” formulas A structural formula is one that shows unambiguously how the atoms are arranged together. A full structural formula (graphic formula/displayed formula) shows every atom and bond. A condensed structural formula can omit bonds between atoms and can show identical groups bracketed together. The use of R to represent alkyl group and benzene ring can be used in condensed structural formulas. 10.1.4 Describe structural isomers as compounds with the same molecular formula but with different arrangements of atoms 10.1.5 Deduce structural formulas for the isomers of the non-cyclic alkanes up to C6. 10.1.7 Deduce structural formulas for the isomers of the straight-chain alkenes up to C6. 10.1.8 Apply IUPAC rules for naming the isomers of the straight-chain alkenes up to C6. 10.1.9 Deduce structural formulas for compounds containing up to six carbon atoms with one of the following functional groups: alcohol, aldehyde, ketone, carboxylic acid and halide Condensed structural formulas can use OH, CHO, Co< COOH, and F/Cl/Br/I 10.1.10 Apply IUPAC rules for naming compounds containing up to six carbon atoms with one of the following functional groups: alcohol, aldehyde, ketone, carboxylic acid and halide. 10.1.11 Identify the following functional groups when the present in structural formulas: amino (NH2), benzene ring and esters (RCOOR). 10.1.12 Identify primary, secondary and tertiary carbon atoms in alcohols and halogenoalkanes 10.1.13 Discuss the volatility and solubility in water of compounds containing the functional groups listed in 10.1.9 10.2 Alkanes 10.2.1 Explain the low reactivity of alkanes in terms of bond enthalpies and bond polarity 10.2.2 Describe using equations, the complete and incomplete combustion of alkanes 10.2.3 Describe, using equations, the reactions of methane and ethane with chlorine and bromine 10.2.4 Explain the reactions of methane and ethane with chlorine and bromine in terms of a free-radical mechanism Reference should be made to hemolytic fission and the reaction steps of initiation, propagation and termination. The use of the half-arrow to represent the movement of a single electron is not required. 10.3 Alkenes 10.3.1 Describe, using equations, the reactions of alkenes with hydrogen and halogens 10.3.2 Describe, using equations, the reactions of symmetrical alkenes with hydrogen halides and water. 10.3.3 Distinguish between “alkanes” and “alkenes” using bromine water. 10.3.4 Outline the polymerization of alkenes Include the formation of poly(ethane) poly(chloroethene) and poly(propene) as examples of addition polymers. Include identification of the repeating unit. 10.3.5 Outline the economic importance of the reactions of alkenes 10.4 Alcohols 10.4.1 Describe, using equations, the complete combustion of alcohols. 10.4.2 Describe, using equations, the oxidation reactions of alcohols A suitable oxidizing agent is acidified potassium dichromate(VI). Equations may be balanced using the symbol [O] to represent oxygen supplied by the oxidizing agent. Include the different conditions needed to obtain good yields of different products, that is, an aldehyde by distilling off the product as it is formed, and a carboxylic acid by heating under reflux. 10.4.3 Determine the products formed by the oxidation of primary and secondary alcohols Assume that tertiary alcohols are not oxidized by potassium dichromate (VI) 10.5 Halogenoalkanes 10.5.1 Describe, using equations, the substitution reactions of halogenoalkanes with sodium hydroxide 10.5.2 Explain the substitution reactions of halogenoalkanes with sodium hydroxide in terms of Sn1 and Sn2 mechanisms Reference should be made to heterolytic fission. Curly arrows should e used ot represent the movement of electron pairs. For tertiary halogenoalkanes the predominant mechanism is Sn1 and for primary halogenoalkanes it is Sn2. Both mechanisms occur for secondary halogenoalkanes. 10.6 Reaction Pathways 10.6.1 Deduce reaction pathways given the starting materials and the product. Conversions with more than two stages will not be assessed. Reagents, conditions, and equations should be included. For example, the conversion of but-2-ene to butanone can be done in two stages: but-2-ene can be heated with steam and a catalyst to form butan-2-ol, which can then be oxidized by heating with acidified potassium dichromate(VI) to form butanone. <> Topic 11: Measurement and Data Processing 11.1 Uncertainty and Error in Measurement 11.1.1 Describe and give examples of random uncertainties and systematic errors. 11.1.2 Distinguish between “precision” and “accuracy” It is possible for a measurement to have great precision yet be inaccurate (for example, if top of meniscus is read in a pipette or a measuring cylinder). 11.1.3 Describe how the effects of random uncertainties may be reduced Repeat readings. 11.1.4 State random uncertainty as an uncertainty range (+/-). 11.1.5 State the results of calculations to the appropriate number of significant figures. 11.2 Uncertainties in calculated Results 11.2.1 State uncertainties as absolute and percentage uncertainties. 11.2.2 Determine the uncertainties in results. For functions such as addition and subtraction, absolute uncertainties can be added. For multiplication, division and powers, percentage uncertainties can be added. If one uncertainty is much larger than others, the approximate uncertainty in the calculated result can be taken as due to that quantity alone. 11.3 Graphical Techniques 11.3.1 Sketch graphs to represent dependences and interpret graph behavior Students should be able to gie a qualitative physical interpretation of a particular graph, for example, the variables are proportional or inversely proportional 11.3.2 Construct graphs from experimental data Choose axes and scale and plotting point. 11.3.3 Draw best-fit lines through data points on a graph 11.3.4 Determine the values of physical quantities from graphs Include measuring and interpreting the slope (gradient) and stating the units for these quantities. Topic 12: Atomic Structure 12.1 Electron Configuration 12.1.1 Explain how evidence from first ionization energies across periods accounts for the existence of main energy levels and sub-levels in atoms 12.1.2 Explain how successive ionization energy data is related to the electron configuration of an atom 12.1.3 State the relative energies of s, p, d and f orbitals in a single energy level 12.1.4 State the maximum number of orbitals in a given energy level. 12.1.5 Draw the shape of an s orbital and the shapes of the Px, Py, and Pz orbitals 12.1.6 Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 54 For Z = 23, the full electron configuration is 1s22s2sp63s23p64s23d3 and the abbreviated electron configuration is [Ar] 4s23d3 or [Ar] 3d34s2. Exceptions to the principle for copper and chromium should be known. Should be familiar with representation of spinning electron in an orbital as an arrow in a box. Topic 13: Periodicity 13.1 Trends Across Period 3 13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure. Include the following oxides and chlorides: * Oxides: Na2O, MgO, Al2O3, SiO2, P4O6, and P4O10, SO2, SO3, Cl2O and Cl2O7 * Chlorides: NaCl, MgCl2, Al2Cl6, SiCl4, PCl3, and PCl5, and Cl2. 13.1.2 Describe the reactions of chloride and the chlorides referred to in 13.1.1 with water. 13.2: First-Row d-Block Elements 13.2.1 List the characteristic properties of transition elements Examples should include variable oxidation number, complex ion formation, existence of colored compounds and catalytic properties 13.2.2 Explain why Sc and Zn are not considered to be transition elements 13.2.3 Explain the existence of variable oxidation number in ions of transition elements Students should known that all transition elements can show an oxidation number of +2. In addition, they should be familiar with the oxidation numbers of the following: Cr (+3, +6), Mn (+4, +7), Fe (+3) and Cu (+1) 13.2.4 Define the term “ligand” 13.2.5 Describe and explain the formation of complexes of d-block elements Include [Fe(H2O)6]3+, [Fe(CN)6]3-, [CuCl4]2-,and [Ag(NH3)2]+. Only monodenate ligands are required. 13.2.6 Explain why some complexes of d-block elements are colored. In complexes, the d sub-level splits into two sets of orbitals of different energy and the electronic transitions that take place between them are responsible for their colors. 13.2.7 State examples of the catalytic action of transition elements and their compounds. Examples should include: * MnO2 in the decomposition of hydrogen peroxide * V2O5 in the Contact process * Fe in the Haber process and in heme * Ni in the conversion of alkenes to alkanes * Co in vitamin B12 * Pd and Pt in catalytic converters. Mechanisms of action will not be assessed. 13.2.8 Outline the economic significance of catalysts in the Contact and Haber processes Topic 14: Bonding 14.1 Shapes of molecules and ions 14.1.1 Predict the shape and bond angles for species with five and six negative charge centers using th VSPR theory Examples should include PCl5, SF6, XeF4, and PF6- 14.2 Hybridization 14.2.1 Describe sigma and π bonds. Include: * Sigma bonds resulting from the axial overlap of orbitals * Π bonds resulting from the sideways overlap of parallel p orbitals * Double bonds formed by one sigma and one π bond * Triple bonds formed by one sigma and two π bonds 14.2.2 Explain hybridization in terms of the mixing of atomic orbitals to form new orbitals for bonding Consider sp, sp2, and sp3 hybridization and the shapes and orientation of these orbitals 14.2.3 Identify and explain the relationships between Lewis structures, molecular shapes and types of hybridization (sp, sp2, and sp3) 14.3 Delocalization of Electrons 14.3.1 Describe the delocalization of π electrons and explain how this can account for the structures of some species Examples should include NO3-, NO2-, CO32-, O3, RCOO- and benzene Topic 15: Energetics 15.1 Standard Enthalpy Changes of Reaction 15.1.1 Define and apply the terms “standard state<” “standard enthalpy change of formation (∆H˚)” and “standard enthalpy change of combustion” (∆Hc˚). 15.1.2 Determine the enthalpy change of a reaction using standard enthalpy changes of formation and combustion 15.2 Born-Haber Cycle 15.2.1 Define and apply the terms “lattice enthalpy” and “electron affinity” 15.2.2 Explain how the relative sizes and the charges of ions affec the lattice enthalpies of different ionic compounds The relative value of the theoretical lattice enthalpy increases with higher ionic charge and smaller ionic radius due to increased attractive forces 15.2.3 Construct a Born-Haber cycle for group 1 and 2 oxides and chlorides, and use it to calculate an enthalpy change 15.2.4 Discuss the difference between theoretical and experimental lattice enthalpy values of ionic compounds in terms of their covalent character. A significant difference between the two values indicates covalent character 15.3 Entropy 15.3.1 State and explain the factors that increase the entropy of a system 15.3.2 predict whether the entropy change (∆S) for a given reaction or process is positive or negative 15.3.3 Calculate the standard entropy change for a reaction (∆S˚) using standard entropy values (S˚) 15.4 Spontaneity 15.4.1 Predict whether a reaction or process will be spontaneous by using the sign of ∆G˚ 15.4.2 Calculate ∆G˚ for a reaction using the equation ∆G˚ = ∆H˚ - T∆S˚ and by using values of the standard free energy change of formation ∆G˚ 15.4.3 Predict the effect of a change in temperature on the spontaneity of a reaction using standard entropy and enthalpy changes and the equation ∆G˚ = ∆H˚ - T∆S˚ Topic 16: Kinetics 16.1 Rate Expression 16.1.1 Distinguish between the terms “rate constant,” overall order of reaction,” and “order of reaction” with respect to a particular reactant. 16.1.2 Deduce the rate expression for a reaction from experimental data 16.1.3 Solve problems involving the rate expression 16.1.4 Sketch, identify and analyze graphical representations of zero, first, and second-order reactions Be familiar with both concentration-time and rate-concentration graphs 16.2 Reaction Mechanism 16.2.1 Explain that reactions can occur by more than one step and that the slowest step determines the rate of reaction (rate-determining step) 16.2.2 Describe the relationship between reaction mechanism, order of reaction, and rate-determining step Only examples with one or two-step reactions where the mechanism given will be assessed. 16.3 Activation Energy 16.3.1 Describe qualitatively the relationship between the rate constant (k) and temperature (T) 16.3.2 Determine activation energy (Ea) values from the Arrhenius equation by a graphical method Topic 17: Equilibrium 17.1 Liquid-Vapor Equilibrium 17.1.1 Describe the equilibrium established between a liquid and its own vapor and how it is affected by temperature changes 17.1.2 Sketch graphs showing the relationship between vapor pressure and temperature and explain them in terms of the kinetic theory 17.1.3 State and explain the relationship between enthalpy of vaporization, boiling point and intermolecular forces 17.2 The Equilibrium Law 17.2.1 Solve homogeneous equilibrium problems using the expression for Kc. Topic 18: Acids and Bases 18.1 Calculations involving acids and bases 18.1.1 State the expression for the ionic product constant of water (Kw) 18.1.2 Deduce [H+ (aq)] and [OH-(aq)] for water at different temperatures given Kw values 18.1.3 Solve problems involving [H+ (aq)] and [OH-(aq)], pH and pOH 18.1.4 State the equation for the reaction of any weak acid or weak base with water, and hence deduce the expressions for Ka and Kb 18.1.5 Solve problems involving solutions of weak acids and bases using the expressions Ka*Kb = Kw pHa + pKb = pKw pH +pOH = pKw Identify the relative strengths of acids and bases using values of Ka, Kb, pKa, and pKb 18.2 Buffer Solutions 18.2.1 Describe the composition of a buffer solution and explain its action 18.2.2 Solve problems involving the composition and pH of a specified buffer system Only examples involving the transfer of one proton will be assessed. Examples should invlude ammonia solution/ammonium chloride and ethanoic acid/sodium ethanoate 18.3 Salt Hydrolysis 18.3.1 deduce whether salts form acidic, alkaline or neutral aqueous solutions examples include salts formed from the four possible combinations of strong and weak acids and bases. The effect of the charge density of the cations in groups 1, 2, and 3 and d-block elements should also be considered. For example: [Fe(H2O)6]3+(aq) [Fe(OH)(H2O)5]2+ + H+(aq) 18.4 Acid-Base Titrations 18.4.1 Sketch the general shapes of graphs of pH against volume for titrations involving strong and weak acids and bases, and explain their important features Only examples involving the transfer of one proton will be assessed. Important features are: * Intercept with pH axis * Equivalence point * Buffer region * Points where pKa = pH or pKb = pOH 18.5 Indicators 18.5.1 Describe qualitatively the action of an acid-base indicator Use Hln(aq) H+(aq) + ln-(aq) (the first part is color a and the second part is color b) 18.5.2 State and explain how the pH range of an acid-base indicator relates to its pKa value 18.5.3 Identify an appropriate indicator for a titration, given the equivalence point of the titration and the pH range of the indicator Topic 19: Oxidation and Reduction 19.1 Standard Electrode Potentials 19.1.1 Describe the standard hydrogen electrode 19.1.2 Define the term “standard electrode potential(E˚) 19.1.3 Calculate cell potentials using standard electrode potentials 19.1.4 Predict whether a reaction will be spontaneous using standard electrode potential values 19.2 Electrolysis 19.2.1 Predict and explain the products of electrolysis of aqueous solutions Explanations should refer to E˚ values, nature of the electrode and concentration of the electrolyte. Examples include the electrolysis of water, aqueous sodium chloride and aqueous copper(II)sulfate 19.2.2 Determine the relative amounts of the products formed during electrolysis The factors to be considered are chanrge on the ion, current and duration of electrolysis 19.2.3 Describe the use of electrolysis in electroplating Topic 20: Organic Chemistry 20.1 Introduction 20.1.1 Deduce structural formulas for compounds containing up to six carbon atoms with one of the following functional groups: amine, amide, ester and nitrile 20.1.2 Apply IUPAC rules for naming compounds containing up to six carbon atoms with one of the following functional groups: amine, amide, ester and nitrile 20.2 Nucleophilic Substitution Reactions 20.2.1 Explain why the hydroxide ion is a better nucleophile than water. 20.2.2 Describe and explain how the rate of nucleophilic substitution in halogenoalkanes by the hydroxide ion depends on the identity of the halogen 20.2.3 Describe and explain how the rate of nucleophilic substitution in halogenoalkanes by the hydroxide ion depends on whether the halogenoalkane is primary, secondary or tertiary 20.2.4 Describe, using equations the substitution reactions of halogenoalkanes with ammonia and potassium cyanide 20.2.5 Explain the reactions of primary halogenoalkanes with ammonia and potassium cyanide in terms of the Sn2 mechanism 20.2.6 Describe, using equations, the reduction of nitriles using hydrogen and a nickel catalyst 20.3 Elimination Reactions 20.3.1 Describe, using equations, the elimination of HBr from bromoalkanes 20.3.2 Describe and explain the mechanism for the elimination of HBr from bromoalkanes 20.4 Condensation Reactions 20.4.1 Describe, using equations, the reactions of alcohols with carboxylic acids to form esters, and state the uses of esters 20.4.2 Describe, using equations, the reactions of amines with carboxylic acids 20.4.3 Deduce the structures of the polymers formed in the reactions of alcohols with carboxylic acids Emphasize the need for two functional groups on each monomer Include the poyester formed from ethane-1,2-diol and benzene-1,4-dicarboxylic acid. Include the identification of the repeating unit 20.4.4 Deduce the structures of the polymers formed in the reactions of amines with carboxylic acids Emphasize the need for two functional groups on each monomer Include the polyamide formed from 1,6-diaminohexane and hexanedioic acid. Include the identification of the repeating unit. 20.4.5 Outline the economic importance of condensation reactions 20.5 Reaction Pathways 20.5.1 Deduce reaction pathways given the starting materials and the product Conversions with more than two stages will not be assessed. Reagents, conditions, and equations should be included. For example, the conversion of 1-bromopropane to 1-butylamine can be done in two stages: 1-bromopropane can be reacted with potassium cyanide to form propanenitrile, which can then be reduced by heating with hydrogen and a nickel catalyst 20.6 Stereoisomerism 20.6.1 Describe stereoisomers as compounds with the same structural formula but with different arrangements of atoms in space. 20.6.2 Describe and explain geometrical isomerism in non-cyclic alkenes Include prefixes cis- and trans- and the term restricted rotation 20.6.3 Describe and explain geometrical isomerism in C3 and C4 cycloalkanes Include the dichloro derivatives of cyclopropane and cyclobutane 20.6.4 Explain the difference in the physical and chemical properties of geometrical isomers Include cis- and trans-1,2-dichloroethene as examples with different boiling points, and cis- and trans-but-2-ene-1,4-dioic acid as examples that react differently when heated 20.6.5 Describe and explain optical isomerism in simple organic molecules. Include examples such as butan-2-ol and 2-bromobutane. The term asymmetric can be used to describe a carbon atom joined to four different atoms or groups. The term chiral can be used to describe a carbon atom joined to four different atoms or groups, and also as a description of the molecule itself. Include the meanings of the terms enantiomer and racemic mixture. 20.6.6 Outline the use of a polarimeter in distinguishing between optical isomers Include the meaning of the term plane-polarized light 20.6.7 Compare the physical and chemical properties of enantiomers Option B: Human Biochemistry B.1 Energy B.1.1 Calculate the energy value of a food from enthalpy of combustion data. B.2 Proteins B.2.1 Draw the general formula of 2-amino acids B.2.2 Describe the characteristic properties of 2-amino acids Properties should include isoelecric point, formation of a zwitterion and buffer action B.2.3 Describe the condensation reaction of 2-amino acids to form polypeptides Reactions involving up to three amino acids will be assessed B.2.4 Describe and explain the primary, secondary (a-helix and ß-pleated sheets), tertiary and quaternary structure of proteins. Include all bonds and interactions (both intramolecular and intermolecular) responsible for the protein strucure B.2.5 Explain how proteins can be analyzed by chromatography and electrophoresis B.2.6 List the major functions of proteins in the body B.3 Carbohydrates B.3.1 Describe the structural featres of monosaccharides Monosaccharides contain a carbonyl group (C=O) and at least two –OH groups and have the empirical formula CH2O B.3.2 Draw the straight-chain and ring structural formulas of glucose and fructose B.3.3 Describe the condensation of monosaccharides to form disaccharides and polysaccharides Examples include: * Disaccharides—lactose, maltose and sucrose * Polysaccharides—starch (a-glucose), glycogen (a-glucose) and cellulose (ß-glucose) B.3.4 List the major functions of carbohydrates in the human body Include energy source (glucose), energy reserves (glycogen) and precursors for other biologically important molecules. B.3.5 Compare the structural properties of starch and cellulose, and explain why humans can digest starch but not cellulose Both are polymers of glucose units. Starch has two forms: amylose, which is a straight-chain polymer (a-1,4 linkage), and amylopectin, which is a branched structure with both a-1,4 linkage; this can be hydrolyzed by the enzyme cellulase, which is absent in most animals, including mammals. B.3.6 State what is meant by the term dietary fiber Dietary fiber is mainly plant material that is not hydrolyzed by enzymes secreted by the human digestive tract but may be digested by microflora in the gut. Examples include cellulose, hemicelulose, lignin and pectin B.3.7 Describe the importance of a diet high in dietary fiber B.4 Lipids B.4.1 Compare the composition of the three types of lipids found in the human body Lipids: poorly soluable in water, soluble in non-polar organic solvents Triglycerides (fats and oils): glycerol bonded to three fatty acids through a condensation reaction with ester bonds Phospholipids (lecithin): glycerol bonded to two fatty acids at one end, and phosphate group on the other end Steroids (cholesterol): polycyclic ring – three cyclohexane rings and one cyclopentane ring B.4.2 Outline the difference between HDL and LDL cholesterol and outline its importance Cholesterol can’t dissolve in blood – has to be transported by lipoproteins which coat cholesterol to make it easier to transport * LDL (low-density lipoprotein) contains more cholesterol than protein, and can build up in inner walls of arteries that feed brain and heart, also can cause atherosclerosis – plaque build up which increases heart disease * HDL (high-density lipoprotein) – contains more protein than cholesterol and protects against heart attack and may remove excess cholesterol B.4.3 Describe the difference in structure between saturated and unsaturated fatty acids * Saturated fats have no double bonds * Usually solids at room temperature, may pack closely together with others * Unsaturated fats have one or more double bonds formed by the removal of hydrogen atoms * Greater number of double bonds, more difficult to pack, lower melting point B.4.4 Compare the structures of the two essential fatty acids linoleic (omega-6 fatty acid) and linolenic (omega-3 fatty acid) and state their importance * Essential fatty acids cannot be made in the body – must obtain from food * Needed for brain function, growth and development, bone health, skin/hair growth * Lower cardiovascular disease * Ratio between Omega-6 and Omega-3 should be between 1:1 – 4:1 * Omega-3 (linolenic acid): found in vegetable oils/processed foods * Cardioprotective benefits and may reduce inflammation and prevent diseases * 3 carbons between first carbon and first double bond * Omega-6 (linoleic acid): imbalance may lead to heart disease, cancer, asthma, arthritis, infections * 6 carbons between first carbon and first double bond B.4.5 Define the term “iodine number” and calculate the number of C=C double bonds in an unsaturated fat/oil using addition reactions. * The number of moles of iodine reacting with one mole of fat/oil indicates the number of double bonds present in the fat/oil molecule * The iodine index/number is the number of grams of iodine (in solution) that adds to 100g of a triacylglyerol. Addition of iodine solution to an unsaturated molecule will cause the double bonds to break to form single-bonded carbon atoms. * Ratio of moles of iodine: one mole of fat = how many double bonds B.4.6 Describe the condensation of glycerol and three fatty acid molecules to make a triglyceride Three fatty acids>>bind to glycerol<< * Condensation reaction * Creates three ester bonds B.4.7 Describe the enzyme-catalyzed hydrolysis of triglycerides during digestion * Pancreatic lipase hydrolyzes triglycerides into two fatty acid chains and a 2-monoglyceride, which allows the small intestine to digest them. * Bile and pancreatic lipase dumped into small intentines when food present. Bile emulsifies fat and lipase hydrolyzes it. B.4.8 Explain the higher energy value of fats as compared to carbohydrates * Fats have very little oxygen = hard to oxidize fat * Carbohydrates have 2 times as many hydrogens as oxygens * Lipids have much higher hydrogen to oxygen concentrations * Therefore, the carbohydrate molecule is more oxidized than the fat molecule * Fat molecule must undergo more extensive oxidation, and the the more oxidation involved in the respiration releases more energy than the carbohydrate respiration B.4.9 Describe the important roles of lipids in the body and the negative effects that they can have on health Important roles: * Energy storage (excess fats) * Insulation and protection of organs * Steroid hormones * Structural component of cell membrane * Omega-3 poly-unsaturated fatty acids reduce risk of heart disease * Poly-unsaturated fats may lower levels of LDL cholesterol Negative effects include: * Increased risk of heart disease from elevated levels of LDL cholesterol and trans fatty acids; the major source of LDL cholesterol = saturated fats * Obesity B.5 Micronutrients and Macronutrients B.5.1 Outline the difference between micronutrients and macronutrients * Micronutrients are substances required in very small amounts that mainly function as a co-factor of enzymes * Vitamins and trace minerals (Fe, Cu, F, Zn, I, Se, Mn, Mo, Cr, Co and B) * Macronutrients are chemical substances that are required in relatively large amounts * Proteins, fats, carbohydrates, and minerals (Na, Mg, K, Ca, P, S and Cl) B.5.2 Compare the structures of retinol (vitamin A), calciferol (vitamin D) and ascorbic acid (vitamin C) * All have polar parts, but have to take size and O:C ratio into account * Vitamin A – long carbon chain with many conjugated C=C double bonds and one –OH group * Fat soluble * Affects maintenance of skin, mucous membranes, bones, teeth, vision * Deficiency may cause night blindness or xerophthalmia (malfunction of tear glands) * Vitamin C – formation and maintenance of collagen: protein that supports body structures like holding together skin, blood vessels and scar tissues * Formation of teeth/bones, enhances absorption of iron * Deficient = scurvy: loss of cementing action of collagen * More polar and water soluble (higher O:C ratio) * Vitamin D – normal bone formation and retention of calcium/phosphorus * Absorbs calcium ions into blood stream/to teeth * Deficient = rickets: bones weaken/bend * Very small O:C ratio (like in vitamin A), similar hydrophobic reactions * Both fat soluble, so stored with fat B.5.3 Deduce whether a vitamin is water or fat-soluable from its structure * Water soluble: vitamins B and C (O:C ratio = high and many –OH groups – capable of hydrogen bonding with water) * Fat-soluble: vitamins A, D, E, K (O:C ratio = small) B.5.4 Discuss the causes and effects of nutrient deficiencies in different countries and suggest solutions * Micronutrient deficiencies * Iron – anemia * Iodine – goiter * Retinol (vitamin A) – xerophtalmia, night blindness * Niacin (vitamin B3) – pellagra: comes from not being able to digest food as well * Thiamin (vitamin B1) – beriberi: inadequate bodily stores of thiamine – loss of sensation, swelling, increased heart rate * Ascorbic acid (vitamin C) – scurvy: several weeks of diet without vitamin C: body can’t be rid of cholesterol/immune system fight off disease * Calciferol (vitamin D) – rickets: where not sunny – need vitamin D for calcium and phosphorus absorption * Macronutrient deficiencies: * Protein = marasmus and kwashiorkor: due to lack of food/malnutrition * Solutions: * Provide food rations composed of fresh, vitamin and mineral rich foods * Adding nutrients missing in commonly consumed foods * Genetic modification of food * Providing nutritional supplements * Providing selenium supplements to people eating foods grown in selenium-poor soil B.6 Hormones B.6.1 Outline the production and function of hormones in the body Hormone | Production location | Derived from | Role in body | Adrenalin/epinephrine | Adrenal cortex (adjacent to kidneys | Amino acid tyrosine | Responsible for flight/fight, increased hart rate. Affects glucagon release into liver and glucose by liver to blood. | Thyroxine: iodine containing amino acid | Thyroid glands | Small molecule from amino acids | Regulating metabolism | Insulin | Pancreas | Protein | Decreases blood glucose level; increases glucose/amino acid uptake by cells | Antidiuretic hormone | Hypothalamus, released by pituitary gland | Protein | Prevents production of dilute urine | Aldosterone | Adrenal cortex | Cholesterol | Balance of salt/water | Sex hormones1.androgens(testosterone)2. estrogen (estradiol)3. progestins (progesterone) | 1. testes2. ovaries3. ovaries | cholesterol | 1. Development & maintenance of male reproductive system/secondary sexual characteristics2. similar role as above3. prepare/maintain uterus | B.6.2 Compare the structures of cholesterol and the sex hormones * Common steroid backbone, difference in functional groups * Backbone = 17 carbon atoms in 3 cyclohexane rings with cyclopentane ring * Functional groups = methyl/hydroxyl groups * * Testosterone * Keytone * Hydroxyl group * Androsterone * Keytone * Hydroxyl group * Cholesterol * Hydroxyl group * C=C double bond * Progesterone * Keytone * Carboxylic acid B.6.3 Describe the mode of action of oral contraceptives * Mimics hormone levels as though pregnant or builds up mucus to not allow sperm through * Synthetic progesterone and estrogen: stop the release of LHRH by the hypothalamus and FSH and LH by the pituitary * Norethynodrel and norethindrone are used * Molecular framework same as progesterone and can bind to receptor sites B.6.4 Outline the use and abuse of steroids * Promotes muscle growth and recover body weight * Male secondary sexual characteristics * Baldness, urinating problems in men * Secondary sex characteristics affected in women * Violent tempers, increased aggressive behavior, liver tumors, high blood pressure, heart attacks B7 Enzymes B.7.1 Describe the characteristics of biological catalysts (enzymes) include: enzymes are proteins; activity depends on tertiary and quaternary structure; and the specificity of enzyme action B.7.2 Compare inorganic catalysts and biological catalysts (enzymes) B.7.3 Describe the relationship between substrate concentration and enzyme activity B.7.4 Determine Vmax and the value of the Michaelis constant (Km) by graphical means and explain its significance. B.7.5 Describe the mechnism of enzyme action, including enzyme substrate complex, active site, and induced fit model B.7.6 Compare competitive inhibition and non-competitive inhibition. B.7.7 State and explain the effects of heavy-metal ions, temperature changes, and pH changes on enzyme activity. B8 Nucleic acids B.8.1 Describe the structure of nucleotides and their condensation polymers (nucleic acids or polynucleotides). Nucleic acids are polymers made up of nucleotides. A nucleotide contains a phosphate group, a pentose sugar and an organic nitrogenous base. Students should recognize the structures of the five bases: Adenine, Cytosine, Guanine, Thymine and Uracil. Nucleic acids are joined by covalent bonds between the phosphate of one nucleotide and the sugar of the next, resulting in a backbone with a repeating pattern of sugar-phosphate-sugar-phosphate. Nitrogenous bases are attached to the sugar of the backbone. B.8.2 Distinguish between the structures of DNA and RNA. RNA has ribose as its pentose sugar; DNA has deoxyribose. Deoxyribose lacks an oxygen atom on C2. RNA has uracil instead of thymine as its base. RNA is a single-strnad nucleic acid; DNA is a double-strand nucleic acid. B.8.3 Explain the double helical structure of DNA The structure has two nucleic acid strands that spiral around an axis. Students should describe the hydrogen bonding between specific pairs of nucleotide bases. B.8.4 Describe the role of DNA as the repository of genetic information, and explain its role in protein synthesis. DNA is the genetic material that an individual inherits from its parents. It directs mRNA synthesis (transcription) and, though mRNA directs protein synthesis (translation) using a triplet code. B.8.5 Outline the steps involved in DNA profiling and state its use. B9 Respiration Compare aerobic and anaerobic respiration of glucose in terms of oxidation/reduction and energy released. In aerobic respiration, glucose is converted into pyruvate, which, in the presence of oxygen, changes to carbon dioxide and water. Overall, glucose undergoes oxidation and oxygen undergoes reduction. In anaerobic respiration, pyruvate is converted to lactate in human beings, whereas yeast converts pyruvate to ethanol and carbon dioxide. Redox equations should be used as appropriate. B.9.2 Outline the role of copper ions in electron transport and iron ions in oxygen transport Cytochromes and hemoglobin are suitable examples. Option G: Further Organic Chemistry G1 Electrophilic Addition Reactions G.1.1 Describe and explain the electrophilic addition mechanisms of the reactions of alkenes with halogens and hydrogen halides Include the application of Markonikov’s rule to predict the major product in the reactions of unsymmetrical alkenes with unsymmetrical reagents G.1.2 Predict and explain the formation of the major product in terms of the relative stabilities of carbocations G2 Nucleophilic Addition Reactions G.2.1 Describe, using equations, the addition of hydrogen cyanide to aldehydes and ketones G.2.2 Describe and explain the mechanism for the addition of hydrogen cyanide to aldehydes and ketones G.2.3 Describe using equations the hydrolysis of cyanohydrins to form carboxylic acids G3 Elimination Reactions G.3.1 Describe, using equations, the dehydration reactions of alcohols with phosphoric acid to form alkenes G.3.2 Describe and explain the mechanism for the elimination f water from alcohols Use H+ to represent the acid catalyst G4 Addition-Elimination Reactions G.4.1 Describe, using equations, the reactions of 2,4-dinitrophenylhydrazine with aldehydes and ketones G5 Arenes G.5.1 Describe and explain the structure of benzene using physical and chemical evidence Physical evidence: comparison of carbon-carbon bond lengths in alkanes, alkenes, and benzene and the number of structural isomers with the formula C6H4X2. Chemical evidence: comparison of enthalpies of hydrogenation of benzene, cyclohexene, 1,3-cyclohexadiene and 1,3,5-cyclohexatriene and the tendency of benzene to undergo substitution rather than addition reactions G.5.2 Describe and explain the relative rates of hydrolysis of benzene compounds halogenated in the ring and in the side-chain Only reactions with t he OH- ion will be assessed. G6 Organometallic Chemistry G.6.1 Outline the formation of Grignard reagents Include reaction of halogenoalkanes with magnesium metal G.6.2 Describe, using equations, the reactions of Grignard reagents with water, carbon dioxide, aldehydes, and ketones G7 Reaction Pathways G.7.1 Deduce reaction pathways given the starting materials and the product Reagents, conditions, and equations should be included. G8 Acid-Base Reactions G.8.1 Describe and explain the acidic properties of phenol and substituted phenols in terms of bonding Include comparison of the acidities of alcohols, phenol and 2,4,6-trinitrophenol G.8.2 Describe and explain the acidic properties of substituted carboxylic acids in terms of bonding G.8.3 Compare and explain the relative basicities of ammonia and amines Include primary, secondary and tertiary amines. Include the formation of salts from amines and the liberation of amines from salts using sodium hydroxide <> G9 Addition-Elimination Reactions G.9.1 Describe, using equations, the reactions of acid anhydrides with nucleophiles to form carboxylic acids, esters, amides and substituted amides Include the nucleophiles: water, alcohols, ammonia and amines. Aspirin and paracetamol can be made using reactions of this type G.9.2 Describe, using equations, the reactions of acyl chlorides with nucleophiles to form carboxylic acids, esters, amides and substituted amides. Include nucleophiles: water, alcohols, ammonia and amines. G.9.3 Explain the reactions of acyl chlorides with nucleophiles in terms of an addition-elimination mechanism G10 Electrophilic Substitution Reactions G.10.1 Describe, using equations, the nitration, chlorination, alkylation and acylation of benzene Include the use of benezene ring symbol as well as formulas such as C6H5NO2. The introduction of more than one group into the benzene ring will not be assessed here. G.10.2 Describe and explain the mechanisms for the nitration, chlorination, alkylation and acylation of benzene. Include the formation of NO2+ from the reaction between concentrated nitric and sulfuric acids and the formation of Cl+, R+, and RCO+ from reactions involving aluminum chloride as a halogen carrier catalyst. G.10.3 Describe, using equations, the nitration, chlorination, alkylation and acylation of methylbenzene G.10.4 Describe and explain the directing effects and relative rates of reaction of different substituents on a benzene ring Include the substituents –CH3, -OH and –NO2. Include the terms activating and deactivating. Only the introduction of one further group will be assessed, except for the formation of 2,4,6-trichlorophenol. The directing effects can be explained in terms of the charge distribution of the intermediates. The slightly increased reactivity due to the presence of –CH3 can be explained in terms of its electro-releasing nature. The creatly increased reactivity in terms of its partial donation of a non-bonded electron pair. The decreased reactivity due to the presence of –NO2 can be explained in terms of its electron-withdrawing nature and lack of a non-bonded electron pair. G11 Reaction Pathways G.11.1 Deduce reaction pathways given the starting materials and the product. Conversions with more than two stages will not be assessed. Reagents, conditions and equations should be included. <>