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User:OldakQuill/Revision/Chemistry

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Topic 6

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  • demonstrate understanding of:
  • recall the characteristic physical (increase in melting/boiling points|decrease in electronegativity) and chemical properties of the elements of group 7 and their compounds limited to:
    • the reactions of halides with the:
      • halogens - Cl oxidises Br and I; Br oxidises I - Cl2 + 2Br- = 2Cl- + Br2
      • phosphoric acid - KCl + H3PO4 -> KH2PO4 + HCl
      • sulphuric acid - Generally NaCl(s) + H2SO4(l) = NaHSO4(s) + HCl(g). Chlorine cannot further reduce. Other two however follow pattern: NaHSO4(s) + 2HBr(g) = SO2(g) + Br2(g) 2H2O(l). Iodine is strong enough reducing agent to further reduce - SO2(g) + 6I2(g) = H2S(g) + 3I2(s) + 2H2O(l)
      • silver nitrate - form silver halide precipitate. Silver chloride is white and soluble in NH3; AgBr is cream, soluble in conc; AgI is yellow, insoluble
    • the properties of silver halides (above)
    • the formation and properties of the hydrogen halides - highly acidic acids - use sulphuric acid to form
  • identify, and make predictions from, the trends in the physical and chemical properties of the halogens and their compounds - boiling/melting points increase down (due to greater number of electrons, hence more van der Waals'); enthalpy hence acidity (in HX) increase down, electronegativity decreases down group (due to greater radius and electron shells), oxidising power decreases due to less attraction at edge (due to more shells)
  • recall uses of the halogens and their compounds - Chlorine kills bacteria in water, chlorine is in bleach, metabolism

Topic 7

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  • demonstrate understanding of the covalent bond in terms of electron sharing - a shared pair of electrons
  • interpret and construct dot-and-cross diagrams of simple covalent compounds
  • demonstrate understanding of the terms bond length and bond angle: paris of electrons move as far away as possible. Non-bonding pairs stronger repulsion. Greater bond energy - smaller length
  • predict and interpret the molecular shapes and bond angles in simple molecules and ions using electron-pair repulsion theory (up to four outer pairs of electrons present as bonding pairs and lone pairs): 4 pairs are tetrahedral (109.5); bonded pairs: 2 pairs are linear (180), 3 pairs are trigonal planar (120); lone pairs: 2 pairs are non-linear (104.5), 3 pairs are pyramidal (107)
  • demonstrate understanding of the terms σ-bond (line joining centres of atoms) and π-bond (two electron clouds on either side - asymmetric density), including the electron density in each type
  • demonstrate understanding of the term electronegativity: the power of an atom in a molecule to attract electrons to itself
  • select data in order to predict the nature of the structure and bonding in a given substance (simple molecular or giant covalent) - covalant is all non-metal; ionic is metal-non-metal; metallic is metal metal, including dative covalency (pair of shared electrons come from one atom), bonding of intermediate type, bond polarity (unequal sharing of electrons) and delocalization (electrons free to move - not fixed between pair)
  • demonstrate understanding of the terms: enthalpy change of atomisation (is the enthalpy change taht takes place when one mole of gaseous atoms is made from the element in its standard state under standard conditions), enthalpy change of combustion (is the enthalpy change that takes place when one mole of a substance reacts compltely with oxygen under standard conditions, all reactants in their standard states) and bond energy (energy required to break bond)
  • calculate bond energies, using Hess’s Law (enthalpy change is independent of the route the reaction takes) and selecting appropriate data
  • understand the relationship of bond energies to the activation energy of a reaction and how molecular collisions contribute to the rate of a reaction(qualitative only)
  • understand that catalysts speed up chemical reactions by providing alternative routes of lower activation energy
  • understand that increases of temperature speed up chemical reactions by increasing the proportion of molecules with the necessary activation energy (qualitative only)
  • recall that some reactions are reversible and understand the dynamic nature of equilibrium reactions; be able to predict the effect of a change in concentration (more on both sides), temperature (endothermic favours products) or pressure (fewer molecules) on the position of equilibrium of a reaction (in simple qualitative cases only).

Topic 8

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Topic 9

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  • interpret changes of state and the associated energy changes in terms of the particles present, their forces of attraction (intermolecular forces broken from solid->gas) and their arrangements (structural in solid - lattice, etc.)
  • demonstrate understanding of van der Waals’ forces (weak forces of attraction between molecules) and dipole-dipole interactions (electronegativities induced by electronegativity in neighbouring molecules)
  • interpret using the concept of van der Waals’ forces and dipole-dipole interactions:
    • properties which imply weak cohesive forces between all molecules (transient dipole-dipole interaction produce coheisve forces between neighbouring - due to attraction)
    • increase in boiling point with increasing size among similar molecules (great size=more electrons=greater asymmetry=greater force), and with increasing surface area among isomers (more electrons in contact=greater foce)
    • van der Waals’ radius (half distance between nuclei of atoms in adjacent molecules) and its relationship to atomic and covalent radius (halfd distance between nuclei of atoms in same molecule)
  • demonstrate understanding of hydrogen bonding, and identify the atoms involved in suchbonding in specified cases (One molecule must have a hydrogen atoms which is very highly positively polarised; so highly, in fact, that it is almost ready to be donated as a proton to a base. The other molecule must have one of the small strongly electronegative atoms of the elements nitrogen, oxgyen or fluorine and this atom must have available lone pair of electrons)
  • interpret using the concept of hydrogen bonding:
    • anomalous physical properties among hydrides - of the elements HCl, HBr and HI increase from one to next; HFl is highest however! HCl, etc. incrase due to van der Waals', HFl large due to hydrogen bond
    • anomalous physical properties among organic compounds -
  • demonstrate understanding of the importance of hydrogen bonding in determining the structures of some materials - used in DNA and enzymes (determin shape) - reason for shape of ice
  • predict some of the properties of an unfamiliar substance which contains hydrogen bonds

Topic 10

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  • demonstrate understanding of the nomenclature and corresponding displayed and structural formulae for halogenoalkanes: Begins with prefix "fluoro-", "chloro-", "bromo-", "iodo-" - follows normal naming conventions for organic compounds. "di-" and "tri-" prefices may be added for compounds with multiple halogens
  • recall the typical behaviour of halogenoalkanes, limited to:
    • combustion - a chlorine ion departs from the alkane
    • treatment with:
      • aqueous alkali: chlorine atoms departs with bonding electrons
      • alcoholic alkali: molecule loses hydrogen and halogen
      • aqueous silver nitrate: No reaction if halogens are in halogenoalkane
      • alcoholic ammonia: two ammonia molecules will react with a halogenoalkane. One replaces the halogen - the other forms a halogen ammonium
  • interpret the reactions of halogenoalkanes in terms of the processes of bond-breaking and bond-making by nucleophilic attack and by reference, as appropriate, to electron pair availability, bond polarisation and bond energy: nucleophilic attack is by a species seeking a nucleus (ie. too many electrons) - an example is alcoholic ammonium; which has a pair of unpaired electrons
  • recall uses of halogenoalkanes: Used as anaesthetics
  • demonstrate understanding of the following terms as associated with organic reactions - homolytic (same electrons present in products) and hereolytic fission (different electrons in products); free radical (uncharged species with unpaired electron), photochemical reaction (reaction stimulated by light); chain reaction (sequence of reactions where a reactive product causes more additional reactions); initiation (net number of free radicals increase); propagation (net number of free radicals stays the same); termination (net number of free radicals decreases); nucleophile (a species seeking positive charge) and electrophile (a species seeking negative charge); addition (atoms in species increases), substitution (atoms in species stay the same), elimination (atoms in species decreases) and hydrolysis reactions (reactions in which water is necessary for species breakdown).

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Topic 11

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  • demonstrate understanding of the terms: rate of reaction (proceedure of reaction in terms of concentration change per time), rate equation (rate=k[reagent] determined by slowest step), order of reaction (rate at which reaction occurs), rate constant (proportionality between rate and reagents), half-life (time taken for the concentration to half), rate-determining step (the slowest step, which causes the order to be what it is), activation energy (minimum energy which needs to be surpassed for the reaction to proceed), catalyst (lowers activation energy), heterogeneous catalysis (when in different phase to reactants)
  • deduce from experimental data for simple zero, first and second order reactions only:
    • half-life - zero: half life decreases; first: half life is constant; second: half life increases
    • order of reaction - zero: double concentration rate remains; first: double concentration rate doubles; second: double concentration rate quadruples
    • rate equation - superscript number is order, add numbers together for reaction order
    • rate-determining step, related to possible reaction mechanisms: slowest step, if you can vary reactants and rate remains - not in rate determining step. Rate determining step in rate equation.
    • activation energy (by graphical methods only; the Arrhenius equation will be given if needed): zero: shallow/flat; first:inverse conc-time; second: steeper
  • predict the effects of concentration, temperature and catalysis (all increase) on the rate of a reaction, and interpret them in terms of collision theory, the distribution of molecular kinetic energies, and alternative reaction pathways of lower activation energy (qualitative only); including links to topic 7
  • evaluate information by extraction from text and the Book of Data about catalysis.

Topic 12

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Topic 13

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  • interpret the natural direction of change (spontaneous change) as the direction of increasing number of ways of sharing energy and therefore of increasing entropy (positive entropy change)
  • recall that the entropy change in any reaction is made up of the entropy change in the system added to the entropy change in the surroundings, summarised by the expression:

∆Stotal = ∆Ssystem + ∆Ssurroundings

  • recall the factors affecting the standard entropy of a substance, in particular its physical state (increase from solid-liquid-gas), and predict the relative entropies of different substances (qualitative only) (soft entropies are higher, complex substances are higher)
  • calculate the standard entropy change in the system for a stated chemical reaction using standard entropy data (∆Ssystem=∆Sproducts - ∆Sreactants
  • recall the expression:

∆Ssurroundings = –∆H/T and use it to calculate entropy changes in the surroundings and, hence, calculate ∆Stotal

  • recall that the feasibility of a reaction depends on the balance between ∆Ssystem and ∆Ssurroundings (has to be positive overall)

Topic 14

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  • demonstrate understanding of the term equilibrium as applied to physical and chemical systems, including links to topic 7 (a reversable reaction where a dynamic balance of products and reactants is met in a closed environment)
  • apply the Equilibrium Law to a chemical reaction in order to deduce the expression for the equilibrium constants, Kc (Kc=products/reactants) and Kp (Kp=products/reactants), and their units
  • perform simple calculations related to the Equilibrium Law
  • predict whether a system is capable of spontaneous change, using ∆S and Kc as indicators of thermodynamic feasibility, and the position of equilibrium (qualitatively only) (∆S needs to be positive, Kc will give position - 1 is in middle, 0.1 favours reactants, 10 favours products)
  • understand and use the term: acid (proton donor), base (proton acceptor), neutral (neither acid nor base), pH (logarithmic scale expressing hydrogen concentration), indicator (substance which visual indicates pH), buffer (a substance which resists change in pH)
  • demonstrate understanding of the term equilibrium as applied to acid-base systems (interchange of hydrogen ions)
  • apply the equilibrium law to an acid-base system in order to deduce the expression for the equilibrium constant, Ka (Ka=products/reactants;
  • demonstrate understanding of practical methods of determining acid and alkali concentrations by titration (datalogger, alkali in burette, including the interpretation of pH changes to select indicators (bromophenol blue for 3-4; bromothymol blue for 7; methyl orange for 3.5-4; phenolphthalein for 9)
  • concentrations by titration including the interpretation of pH changes to select indicators
  • plan an investigation using acid-alkali titrations, and justify the procedures involved
  • perform calculations from acid-alkali titration data
  • demonstrate understanding of the Brønsted-Lowery theory of acid-base behaviour, and use it to interpret the behaviour of strong acids/bases (great dissociation), weak acids/bases, conjugate acid-base pairs (H2O = H+ + OH-, and buffer solutions (resist change in pH
  • recall the terms pH (= -log [H+]), Ka and Kw (Kw = [H+] + [OH-]) and perform simple calculations using them, including calculating the pH of buffer solutions
  • deduce and interpret qualitatively the effect of changes in temperature and pressure on systems at equilibrium in terms of entropy changes (temperature favours products if endothermic - positive Enthalpy value; pressure favours side with fewer molecules); and the value of equilibrium constants.

Topic 15

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  • demonstrate understanding of the nomenclature and corresponding displayed and structural formulae for carbonyl compounds (in aldehydes -al; and ketones -one - ketones have a random group on the other carbon bond), carboxylic acids (-oic acid), esters (carbon double bonded to O; bonded to another O which is bonded to a random group - if methyl is the random group and butanoic acid was original - then methyl butanoate and acid chlorides (double bond O, single bond Cl - ethanoyl chloride)
  • recall the typical behaviour of aldehydes and ketones limited to:
  • recall the typical behaviour of carboxylic acids limited to:
    • solubility in water: only C1 to C4
    • acidity and formation of salts: weak acids in water (less than 1% ionised). Strong enough to displace carbon dioxide from sodium carbonate; will also neutralised sodium hydroxide to form salt.
    • reaction with alcohols to form esters: due to alcohol's lone pair, warmed - acid acts as catalyst (sulphuric)
    • relationship to alcohols, aldehydes and ketones:
  • recall the typical behaviour of esters and acyl chlorides limited to hydrolysis by water (acid or base catalyst necessary - acid removes alcohol; alkali forms an anion (like sodium) which displaces alcohol)
  • interpret the reactions of carboxylic acids, aldehydes and ketones, in terms of nucleophilic or electrophilic attack, including the relative electron attracting power of the different groups in acids and their derivatives, oxidation and reduction
  • interpret the infra-red spectra of acids and carbonyl compounds, including the effect of hydrogen bonding
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