Edexcel AS Chemistry Unit 1

Edexcel AS Chemistry Unit 1

1. Formulae, equation & moles

   • Terms
   • Equations
   • Concentrations
     • Safety precautions. If dilute acid spilt= add water to dilute it/ ass alkali to neutralise it/ mop it up with wet paper towel.  They are corrosive.  PPM
   • Moles
     • The amount of any substance containing the same number of entities as there are atoms in 12g of Carbon-12. Moles (solid)= mass/Mr Moles (liquid)=concentration times volume Moles gas= volume (dm^3)/24
   • Molar volume
     • Moles in a gas= volume (dm^3)/ 24
   • Percentage Yield
     • Actual yield/ theoretical yield multiplied by 100. Actual yield will be given in the question. Find the theoretical yield using moles/ mass equation.
   • Atom economy
   • Avogadro constant
     • Used to find the number of atoms/ particles in a molecule. Use the equation with moles and Av which is a specific number that they will tell you in the question. 
   • Experiments

2. Energetics & enthalpy changes

   • Terms
     • Activation energy: the minimum energy required for the particles to collide successfully and react. 
   • Enthalpy level diagrams
     • Energy Level Diagrams
   • Endothermic & exothermic
     • Endothermic: enthalpy change is positive. Requires heat energy and shown in reduction of temperature. Usually bond breaking.  Exothermic: enthalpy change is negative. Gives out heat energy and shown in an increase in temperature. Usually bond making
   • Definitions of standard enthalpy changes
     • Formation Combustion Atomisation Neutralisation Reaction (ones needed for Edexcel AS)
   • Hess's Law
     • Hess's Law states that the enthalpy change of a reaction is independent of the route taken. 
     • Limitations
       • Average bond enthalpies used. All reactants and products must be in their standard states. 
   • Coffee cup experiment
     • Q= mass times 4.18 times change in temperature. (Q=MCT)
     • Sources of error/assumptions
       • Heat loss, imcomplete combustion of fuel, evaporation of fuel from spirit lamp.  Assumes that the density of the liquid is the same as water. Assumes there are standard states. Assumes that the specific heat capacity of any solution is the same as water. 
   • Bond enthalpy
     • The energy needed to break one mole of a bond in a gaseous molecule, forming gaseous atoms. 
   • Standard states
     • 1 atmosphere of pressure (100KPa) and a stated temperature (298K)

3. Atomic structure & the periodic table

   • Definitions
     • Allotropes Mole Isotope Relative isotopic mass Relative atomic mass
   • Mass spectrometer
     • 1.  Vaporisation 2. Ionisation- bombarded with high energy electrons which knock off electrons. 3. Acceleration- in electric field.  4. Deflection- lighter/highly charged ions are deflected more. In magnetic field.  5. Detection- forms a mass spectrum
   • Uses of mass spectrometer
     • radioactive dating in space research. In sport for drug testing. 
   • Definition of ionisation energies
     • 1st ionisation energy: the energy required to remove 1 electron from 1 mol of gaseous atoms to form 1 mol of gaseous 1+ ions.  2nd ionisation energy: the energy required to remove 1 electron from 1 mol of gaseous 1+ ions to form 1 mol of gaseous 2+ ions.  Ionisation energies are always endothermic (positive)
   • Electronic structure
     • S, P orbitals and their shapes
   • Trends & properties
     • Ionisation energy increases across period and decreases down group due to shielding/ increased atomic radius/ decreased nuclear attraction. 
     • Polarisation
       • Small, highly charged cation and large anion. Will lead to some covalent character. This is proven when theoretical and experimental lattice enthalpies differ on Born Haber cycles. This is the reason for this. The small positive cation polarises the large anion. 
     • Periodicity
       • Repeating pattern or trends across a period.
   • Solubility rules
     • Remember from GCSE. Learn a table. This will help when constructing balanced equation to know if something is (aq) or (s). Usually for longer questions.
   • Indicators
     • Learn for titrations. Phenolphthlalein and methyl orange are the ones needed to know. Phenol= pink in alkali, colourless in acid. Methyl orange= red in acid, yellow in alkali

4. Bonding

   • Ionic bonding
     • Electrostatic forces of attraction between oppositely charged ions. Forms a lattice. Also repulsion between negative electron clouds.  Ionic compounds only conduct electricity in molten state as ions are free to move around the whole structure and carry a charge- they are fixed in solid state.
   • Covalent Bonding
     • attraction between the nucleus and the outer electrons.  Can be simple (S) or giant (Si). Simple only has weak intermolecular forces but giant/ macromolecular has strong covalent bonds. Electrostatic attraction.
   • Metallic Bonding
     • Electrostatic forces of attraction between positive ions and negative delocalised electrons. The more electrons, the more forces. The greater the charge of the ion, the stronger the forces. Giant lattices of metal ions surrounded by a sea of delocalised electrons. 

5. The alkanes

   • General
     • CnH2n+2 Burn with clean flame. Hydrocarbons( only have carbon and hydrogen atoms) Are saturated Only single bonds Methane, ethane, propane, butane, pentane, hexane, heptane, octane, nonane, decane etc...
   • Isomers
     • Can be shown in skeletal/ displayed/ structural formula. Find isomers by drawing out skeletal then adding methyl groups, shortening the longest chain but still has the same number of atoms.
   • Fuel
     • Definition of a fuel: something when burned produces heat/ energy.   
   • Sustainability
     • Not sustainable as from crude oil which is non-renewable.  Alternate greener methods could be biofuels (from plants= carbon neutral)
   • Mechanisms
     • Learn mechanism for free radical homolytic substitution reaction (with Cl2 and UV light).
     • Mechanisms
     • Substitution Reactions
   • Less reactive than alkenes
     • Weak pi bond. 
   • Boiling point increases with chain length
     • The number of electrons per molecule increases so the strength of the London forces increases.  More branching = lower boiling point. (Less branching allows closer contact between the molecules creating stronger London forces.)
   • Cycloalkanes
     • Same general formula as alkenes (CnH2n)

6. The alkenes

   • General
     • Same general formula as Cycloalkanes (CnH2n) Have a double bond. 
   • Addition Reactions
     • With hydrogen
       • Nickel catalyst, 150 degrees celcius. Forms an alkane. 
     • With halogens
       • Halogen should be dissolved in hexane, room temp.  (involved in the test for alkenes: Br2 water. Orange to colourless). Forms a di-substituted haloalkane
     • With hydrogen halides
       • The hydrogen halide should be in the gas phase, room temp. Forms a mono-substituted haloalkane. 
     • with Potassium manganate
       • Forms a diol (2 OH groups). Needs dilute sulphuric acid, room temp. Potassium manganate is an oxidising agent. Purple to colourless. 
   • Mechanisms
     • Electrophilic addition- 100% atom economy as only 1 product.  e.g alkene+ HBr or Br2.  REMEMBER: propene with HBr has a major and a minor product. Major more likely to form as a secondary haloalkane rather than a primary which is more stable due to alkyl groups being electron releasing.
     • Mechanisms
     • Electrophilic addition
   • Polymerisation
     • 100% yield as only 1 product.  Monomers join to form a polymer- exam may ask to draw repeat units of product. 
     • Polymerisation.
   • Geometric Isomerism
     • E/Z
       • Happens when cis trans breaks down. Z= cis E= trans  Whichever takes the highest priority (highest Mr)
     • cis/trans
     • restricted rotation around double bond if there are different groups on each carbon