Unit 2: Energetics, Group Chemistry, Halogenoalkanes and Alcohols

6.1The enthalpy change, ∆H, is the heat energy change measured at constant pressure and that standard conditions are 100 kPa

Know that the enthalpy change, ∆H, is the heat energy change measured at constant pressure and that standard conditions are 100 kPa and a specified temperature, usually 298 K

6.2That, by convention, exothermic reactions have a negative enthalpy change and endothermic reactions have a positive enthalpy

Know that, by convention, exothermic reactions have a negative enthalpy change and endothermic reactions have a positive enthalpy change

6.3Construct and interpret enthalpy level diagrams, showing exothermic and endothermic enthalpy changes

Be able to construct and interpret enthalpy level diagrams, showing exothermic and endothermic enthalpy changes

6.4The definition of standard enthalpy change of: i reaction, ∆rH ii formation, ∆fH iii combustion, ∆cH iv neutralisation

Know the definition of standard enthalpy change of: i reaction, ∆rH ii formation, ∆fH iii combustion, ∆cH iv neutralisation, ∆neutH v atomisation, ∆atH

6.5Experimental data to calculate: i energy transferred in a reaction recalling and using the expression: energy transferred

Be able to use experimental data to calculate: i energy transferred in a reaction recalling and using the expression: energy transferred (J) = mass (g) × specific heat capacity (J g-1 °C-1) × temperature change (°C) ii enthalpy change of the reaction in kJ mol⁻¹ This will be limited to experiments where substances are mixed in an insulated container and combustion experiments using a suitable calorimeter.

6.6Hess’s Law

Know Hess’s Law and be able to apply it to: i constructing enthalpy cycles ii calculating enthalpy changes of reaction using data provided, or data selected from a table or obtained from experiments

6.7CORE PRACTICAL 2 Determination of the enthalpy change of a reaction using Hess’s Law

CORE PRACTICAL 2 Determination of the enthalpy change of a reaction using Hess’s Law.

6.8Evaluate the results obtained from experiments and comment on sources of error and uncertainty and any assumptions made

Be able to evaluate the results obtained from experiments and comment on sources of error and uncertainty and any assumptions made in the experiments Students will need to consider experiments where substances are mixed in an insulated container and combustion experiments using, for example, a spirit burner and be able to draw suitable graphs and use cooling curve corrections.

6.9The terms ‘bond enthalpy’ and ‘mean bond enthalpy’

Understand the terms ‘bond enthalpy’ and ‘mean bond enthalpy’, and be able to use bond enthalpies to calculate enthalpy changes, understanding the limitations of this method

6.10Mean bond enthalpies from enthalpy changes of reaction

Be able to calculate mean bond enthalpies from enthalpy changes of reaction

6.11Bond enthalpy data gives some indication about which bond will break first in a reaction, how easy or difficult it is

Understand that bond enthalpy data gives some indication about which bond will break first in a reaction, how easy or difficult it is and therefore how rapidly a reaction will take place at room temperature

Understand the nature of the following intermolecular forces: i London forces (instantaneous dipole-induced dipole) ii permanent dipole-permanent dipole interactions iii hydrogen bonds

Understand the interactions in molecules, such as H2O, liquid NH3 and liquid HF, which give rise to hydrogen bonding

Understand the following anomalous properties of water resulting from hydrogen bonding: i its high melting and boiling temperature when compared with similar molecules ii the density of ice compared to that of water

Be able to predict the presence of hydrogen bonding in molecules analogous to those mentioned in 7.2

Understand, in terms of intermolecular forces, physical properties shown by substances, including: i the trends in boiling temperatures of alkanes with increasing chain length ii the effect of branching in the carbon chain on the boiling temperatures of alkanes iii the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons iv the trends in boiling temperatures of the hydrogen halides HF to HI

Understand factors that influence the choice of solvents, including: i water, to dissolve some ionic compounds, in terms of the hydration of the ions ii water, to dissolve simple alcohols, in terms of hydrogen bonding iii water, as a poor solvent for compounds (to include polar molecules such as halogenoalkane), in terms of inability to form hydrogen bonds iv non-aqueous solvents, for compounds that have similar intermolecular forces to those in the solvent

Know what is meant by the term ‘oxidation number’ and understand the rules for assigning oxidation numbers

Be able to calculate the oxidation number of elements in compounds and ions, including in peroxides and metal hydrides

Be able to indicate the oxidation number of an element in a compound or an ion, using a Roman numeral

Be able to write formulae given oxidation numbers

Understand oxidation and reduction in terms of electron transfer and changes in oxidation number, and the application of these ideas to reactions of s-block and p-block elements

Know that oxidising agents gain electrons and reducing agents lose electrons

Understand that a disproportionation reaction involves an element in a single species being simultaneously oxidised and reduced

Know that oxidation number is a useful concept in terms of the classification of reactions as redox and as disproportionation

Understand that metals, in general, form positive ions by loss of electrons with an increase in oxidation number whereas non-metals, in general, form negative ions by gain of electrons with a decrease in oxidation number

Be able to write ionic half-equations and use them to construct full ionic equations 8B: The elements of Groups 1 and 2 Students will be assessed on their ability to:

Understand reasons for the trend in ionisation energy down Groups 1 and 2

Understand reasons for the trend in reactivity of the elements down Group 1 (Li to K) and Group 2 (Mg to Ba)

Know the reactions of the elements of Group 1 (Li to K) and Group 2 (Mg to Ba) with oxygen, chlorine and water

Know the reactions of: i oxides of Group 1 and 2 elements with water and dilute acid ii hydroxides of Group 1 and 2 elements with dilute acid

Know the trends in solubility of the hydroxides and sulfates of Group 2 elements

Understand the reasons for the trends in thermal stability of the nitrates and the carbonates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involved

Understand the formation of characteristic flame colours by Group 1 and 2 compounds in terms of electron transitions Students will be expected to know the flame colours for Group 1 and 2 compounds.

Know experimental procedures to show: i patterns in the thermal decomposition of Group 1 and 2 nitrates and carbonates Students will be expected to know tests for carbon dioxide and oxygen; and to recognise nitrogen dioxide by its colour and acidic pH. ii flame colours in compounds of Group 1 and 2 elements

Know reactions, including ionic equations where appropriate, for identifying: (i) carbonate ions, CO3^2−, and hydrogencarbonate ions, HCO3−, using aqueous acid to form carbon dioxide and limewater to test the gas; (ii) sulfate ions, SO4^2−, using acidified barium chloride solution; and (iii) ammonium ions, NH4+, using sodium hydroxide solution and warming to form ammonia, tested with litmus and HCl fumes.

Be able to calculate solution concentrations, in mol dm-3 and g dm-3, including simple acid-base titrations using the indicators methyl orange and phenolphthalein

CORE PRACTICAL 3 Finding the concentration of a solution of hydrochloric acid.

Understand how to minimise the sources of measurement uncertainty in volumetric analysis and estimate the overall uncertainty in the calculated result

CORE PRACTICAL 4 Preparation of a standard solution from a solid acid and use it to find the concentration of a solution of sodium hydroxide.

Understand reasons for the trends for Group 7 elements in: i melting and boiling temperatures and physical state at room temperature ii electronegativity iii reactivity down the group

Understand the trend in reactivity of Group 7 elements in terms of the redox reactions of Cl2, Br2 and I2 with halide ions in aqueous solution Students are expected to know the colours of the elements in standard conditions, in aqueous solution and in a non-polar organic solvent.

Understand, in terms of changes in oxidation number, the following reactions of the halogens: i oxidation reactions with Group 1 and 2 metals ii the disproportionation reaction of chlorine with water and the use of chlorine in water treatment iii the disproportionation reaction of chlorine with cold, dilute aqueous sodium hydroxide to form bleach iv the disproportionation reaction of chlorine with hot alkali v reactions analogous to those specified above

Understand the following reactions: i solid Group 1 halides with concentrated sulfuric acid, to illustrate the trend in reducing ability of the hydrogen halides ii precipitation reactions of the aqueous anions Cl-, Br- and I- with aqueous silver nitrate solution and nitric acid, and the solubility of the precipitates in aqueous ammonia solution iii hydrogen halides with ammonia gas (to produce ammonium halides) and with water (to produce acids)

Be able to make predictions about fluorine and astatine and their compounds, in terms of knowledge of trends in halogen chemistry

Understand, in terms of the collision theory, the effect of changes in concentration, temperature, pressure and surface area on the rate of a chemical reaction

Understand that reactions take place only when collisions have sufficient energy, known as the activation energy

Be able to calculate the rate of a reaction from: i the time taken for a reaction, using rate = 1/time ii the gradient of suitable graph, by drawing a tangent, either for initial rate, or at a time, t

Understand qualitatively, in terms of the Maxwell-Boltzmann distribution of molecular energies, how changes in temperature affect the rate of a reaction

Understand the role of catalysts in providing alternative reaction routes of lower activation energy

Be able to draw the reaction profiles for uncatalysed and catalysed reactions, including the energy level of the intermediate formed with the catalyst

Understand the use of catalysts in industry to make processes more sustainable by using less energy and/or higher atom economy

Be able to interpret the action of a catalyst in terms of a qualitative understanding of the Maxwell-Boltzmann distribution of molecular energies

Know that many reactions are readily reversible and that they can reach a state of dynamic equilibrium in which: i the rate of the forward reaction is equal to the rate of the backward reaction ii the concentrations of the reactants and the products remain constant

Be able to predict and justify the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium in a homogeneous system

Evaluate data to explain the necessity, for many industrial processes, to reach a compromise between the yield and the rate of reaction

Be able to classify reactions (including those in Unit 1) as addition, elimination, substitution, oxidation, reduction, hydrolysis or polymerisation

Understand the concept of a reaction mechanism

Understand that heterolytic bond breaking results in species that are electrophiles or nucleophiles

Know the definition of the term ‘nucleophile’

Understand the link between bond polarity and the type of reaction mechanism a compound will undergo 10B: Halogenoalkanes Students will be assessed on their ability to:

Understand the nomenclature of halogenoalkanes and be able to draw their structural, displayed and skeletal formulae

Understand the distinction between primary, secondary and tertiary halogenoalkanes

Understand the reactions of halogenoalkanes with: i aqueous alkali, including KOH(aq) to produce alcohols (where the hydroxide ion acts as a nucleophile) ii ethanolic potassium hydroxide to produce alkenes by an elimination reaction (where the hydroxide ion acts as a base) iii aqueous silver nitrate in ethanol (where water acts as a nucleophile) iv alcoholic ammonia under pressure to produce amines (where the ammonia acts as a nucleophile) v alcoholic potassium cyanide to produce nitriles (where the cyanide ion acts as a nucleophile) Students should know this is an example of increasing the length of the carbon chain.

Understand the mechanisms of the nucleophilic substitution reactions between primary halogenoalkanes and: i aqueous potassium hydroxide ii ammonia SN1 and SN2 substitution mechanisms will be tested in Unit 4.

Understand that experimental observations and data can be used to compare the relative rates of hydrolysis of: i primary, secondary and tertiary structural isomers of a halogenoalkane ii primary chloro-, bromo- and iodoalkanes using aqueous silver nitrate in ethanol

CORE PRACTICAL 5 Investigation of the rates of hydrolysis of some halogenoalkanes.

Know the trend in reactivity of primary, secondary and tertiary halogenoalkanes

Understand, in terms of bond enthalpy, the trend in reactivity of chloro-, bromo- and iodoalkanes

CORE PRACTICAL 6 Chlorination of 2-methylpropan-2-ol with concentrated hydrochloric acid.

Understand the nomenclature of alcohols and be able to draw their structural, displayed and skeletal formulae

Understand the distinction between primary, secondary and tertiary alcohols

Understand the reactions of alcohols with: i oxygen in air (combustion) ii halogenating agents • PCl5 to produce chloroalkanes (including its use as a qualitative test for the presence of the –OH group) • 50% concentrated sulfuric acid and potassium bromide to produce bromoalkanes • red phosphorus and iodine to produce iodoalkanes iii concentrated phosphoric acid to form alkenes by elimination Descriptions of the mechanisms of these reactions are not required.

Understand that potassium dichromate(VI) in dilute sulfuric acid can oxidise: i primary alcohols to produce aldehydes (which give a positive result with Benedict’s or Fehling’s solution) if the product is distilled as it forms ii primary alcohols to produce carboxylic acids (which give a positive result with sodium carbonate or sodium hydrogencarbonate) if the reagents are heated under reflux iii secondary alcohols to produce ketones In equations, the oxidising agent can be represented by [O].

Understand, the following techniques in the preparation and purification of a liquid organic compound: i heating under reflux ii extraction with a solvent using a separating funnel iii distillation iv drying with an anhydrous salt v boiling temperature determination

CORE PRACTICAL 7 The oxidation of propan-1-ol to produce propanal and propanoic acid.

Be able to interpret data from mass spectra to suggest possible structures of simple organic compounds using the m/z of the molecular ion and fragmentation patterns

Be able to use infrared spectra, or data from infrared spectra, to deduce functional groups present in organic compounds, and predict infrared absorptions, given wavenumber data, due to familiar functional groups including: i C–H stretching absorptions in alkanes, alkenes and aldehydes ii C=C stretching absorption in alkenes iii O–H stretching absorptions in alcohols and carboxylic acids iv C=O stretching absorptions in aldehydes, ketones and carboxylic acids v C–X stretching absorption in halogenoalkanes vi N-H stretching absorption in amines

CORE PRACTICAL 8 Analysis of some inorganic and organic unknowns.