OCR AS Chemistry

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  • Created on: 09-10-18 18:15
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  • OCR AS Chemistry
    • 1- Atoms and Moles
      • Made up of 3 sub-atomic particles: PROTON (+1) NEUTRON (0) ELECTRON (-1)
      • RAM - Average mass of an atom on a scale where 1/12th of Carbon-12 is exactly 12
      • Relative Isotopic Mass - Average mass of an atom of an isotope compared to 1/12th of the mass of Carbon-12
      • Relative Molecular Mass - Average mass of a molecule compared to 1/12th of the mass of Carbon-12
      • Relative Formula Mass - Average mass of a formula unit compared to 1/12th of the mass of Carbon-12
      • The Mole - unit used to measure number of particles in a sample - uses Avagadro's Constant - 6.022x10^23
        • Used to calculate mass, Mr, number of particles present, volume of solutions, volume of gases and concentrations
          • No. of particles present / Avagadro's = MOLES
          • Mass / Mr = MOLES
          • Vol. (aq/l) x Conc. = MOLES
          • Vol. (g) / 24 (dm^3) = MOLES
            • Gas Volume can also be Calculated using the IDEAL GAS EQUATION
              • pV=nRT
    • 2- Formulas and Equations
      • There are two  types of formulae in chemistry - EMIPIRICAL and MOLECULAR
        • EMPIRICAL - smallest whole number ratio of atoms of each element in a compound
        • MOLECULAR - actual number of atoms of elements in a compound
      • Equations must be balanced
        • IONIC EQUATIONS - reactions involving ions in solution - only reacting particles included i.e. H+ + OH- -> H2O
        • Calculations - use MOLES equations depending on info provided
    • 3- Reactions and Calculations
      • Acids, Bases and Salts
        • ACID - proton donor  -releases H+ ions when mixed with water i.e. HCl
          • Strong Vs Weak - Strong Acid/Base FULLY dissociates in water - Weak PARTIALLY  dissociates in water
        • SALT -  product of neutralisation reaction
          • Neutralisation can occur in many ways
            • Metal + Acid => Salt + Hydrogen
            • Metal Oxide + Acid => Salt + Water
            • Acid + Metal Hydroxide => Salt + Water
            • Metal Carbonate + Acid => Salt + Carbon Dioxide + Water
            • Acid + Base => Salt + Water
          • Hydrated and Anhydrous - Hydrated = contains water of crystallisation - Anhydrous = doesn't contain water of crystallisation
            • Calculating formula of Hydrated Salts -  1) mass of anhydrous - mass of hydrated        2) Find moles of water lost  3) Calculate number of moles of anhydrous produced      4) work out ratio of moles of anhydrous to moles of water
        • Titrations - allow us to see how much acid is needed in order to neutralise an alkali
          • Indicators used in order to identify end point and point of neutralisation
            • i.e. methyl orange, phenolphthalei-n
          • Precision - mass - reset balance + make sure all solid is transferred solution - volumetric flasks and pipette
            • Standard Solutions - solution that has a precisely known concentration - known amount of solid dissolved in known amount of water to create known concentration
        • BASES - proton acceptor -  remove H+ ions and release OH- ions i.e. NaOH - soluble bases = alkalis
          • Concentrated Vs Dilute - Concentrated Acid/Base has more H+/OH- ions - Diluted has less H+/OH- ions
            • ACID - proton donor  -releases H+ ions when mixed with water i.e. HCl
              • Strong Vs Weak - Strong Acid/Base FULLY dissociates in water - Weak PARTIALLY  dissociates in water
      • actual yield / theoretical yield x 100 = % yield
        • theoretical yield = mass/Mr to find moles and then multiply moles by molar mass
      • Atom economy - mass of desired (Mr)/mass of reactants (Mr)
      • Oxidation Numbers
        • Shows total number of electrons an element has accepted or donated
          • Uncombined = 0 - Simple ions = same as charge Molecular ions = overall charge = overall oxidation number Neutral compound = 0 H = +1 F = -1
            • Redox = gain in electrons - reducing agent donates electrons and gets oxidised
              • Oxidation no. will increase by 1 for every electron lost and it will decrease by 1 for every electron gained
    • 4- Electrons, Bonding and Structure
      • Bonding
        • Ionic
          • Electrostatic forces of attraction between positive and negative ions
          • Conduct  electricity when molten or in solution - ions carry charge and are free to flow when in liquid form
          • High MP and BP due to large amount of energy needed to overcome electrostatic forces of attraction
          • Soluble in water due to polar water molecules being able to separate ions
        • Covalent Bonding
          • Shared pair of electrons between 2 atoms
          • Dative covalent bonding - one atoms provides both of the shared electrons
          • Strength - depends on AVERAGE BOND ENTHALPY - energy required to break a covalent bond
        • Shapes of Molecules
          • Electron Repulsion Theory -  lone pairs repel more than bonding pairs
          • Linear - 180 2 bonds  - CO2 Trigonal planar - 120 3 bonds - BF3 Tetrahedral 109.5 4 bonds - NH4  Trigonal pyramidal - 107 3 bonds - NH3          Non-linear - 104.5 2 bonds - H2O    Trigonal Bipyramidal - 120, 90 5 bonds - PCl5 Octahedral - 90 6 bonds - SF6
          • Polarisation
            • Electronegativi-ty - ability of and tom to attract bonding electrons in covalent bond
            • Polar Bonds - bonding electrons pulled towards more electrone-gative atom
              • In a polar bond, difference in electronegativi-ty causes a permanent dipole which is the difference in charge between 2 atoms caused by shift in electron density in the bond
                • Polar molecules have an overall dipole - arrangement of polar bonds determines whether it will have an overall dipole or not
            • Non Polar Bonds - atoms have equal electronegativi-ty
      • Electrons - Subatomic particles with -1 charge
        • Has sub-shells - s, p, d, f - can be used to identify certain elements i.e. s-block elements = alkali metals
      • Intermolecular forces
        • Forces between molecules - weaker than covalent, ionic or metallic
          • Induced dipole interactions - cause all atoms to be attracted to eachother - they are temporary - they are constantly being created and destroyed due to constantly moving electrons
          • Permanent dipole interactions - cause weak electrostatic forces of attraction between molecules - happen in addition to induced
          • Hydrogen bonding - only happens when H is covalently bonded to O,N or F - molecules with H bonding usually have OH or NH groups
    • 5- The Periodic Table
      • Modern Periodic table goes up in atomic mass (no. of protons)
      • 1st ionisation energy
        • The energy needed to remove 1 mole of electrons from one mole of gaseous atoms
        • Nuclear charge, atomic radius and shielding affect ionisation energy
      • Metals
        • Metallic bonding - cations in a sea of delocalised electrons
          • Allows metals to have high MP and BP due to more energy being required to over come bonds - they are also malleable and ductile due to the same size atoms sliding over each other - electrons carry current which allows them to conduct electricity - insoluble except for liquid metals
        • Group 2 Alkali Metal - down group, they react more due to shielding and radius (ionisation energy decreases) - used for neutralisation i.e. Ca(OH)2 or as antacids
      • Halogens - they form group 7
        • Boiling point down group increases due to increased strength of induced dipole interactions
        • Displacement reactions occur - more reactive at top of group
        • Testing for ions
          • CO3 + acid = bubbles
          • SO4+Ba(NO3)2 = white precipitate
          • Halide + AgNO3 = precipitate
          • NH4+/NH3 = red litmus turns blue
    • 6- Physical Chemistry
      • Enthalpy
        • Heat transferred in a reaction in standard conditions with all substances in standard states
        • Types include: reaction, formation, combustion and neutralisation
        • q=mc(delta)T
      • Hess's Law
        • Reaction's enthalpy change is the same regardless of the route taken
      • Rate
        • Maxwell-Boltzmann
        • Temp. = particles have higher KE = more frequent and successful collisions = higher rate
        • Conc. = more particles  = more frequent and successful collisions = higher rate
        • Catalysts - HOMO (same state) HETERO (different state) not used up themselves
          • Reduce costs - increase rate due to lower Ea
    • 7- Basic Organic Concepts and Hydrocarbons
      • Alkanes - branched or unbranched - un have lower BP/MP - no double bond
        • Substitution reactions - initiation, propagation and termination
          • Free radicals involved - OZONE
      • Combustion - incomplete and complete
        • CO poisoning
      • Alkenes - double bond (unsaturated)
        • E/Z and cis and trans are stereoisomers
          • Electrophilic addition
            • Curly arrow mechanism with ELECTROPHI-LE i.e. Br
            • Forms haloalkane
        • Can be used to make polymers
        • Reactions of alkenes
          • Hydrated, 300C and H3PO4 catalyst that is solid
          • Add bromine water to test for saturation - decolourises if is alkene
        • Markownikoff's rule - more stable carbocations form i.e. + charge is on 2nd C in propene
    • 8- Alcohols and Haloalkanes
      • Alcohols
        • General formula is CnH2n+1OH
          • Polar due to electronegativ-e OH - alcohol + water = HYDROGEN BONDS FORM
            • Undergo substitution reactions to form haloalkanes - requires H2SO4 as catalyst
              • Undergo nucleophilic substitution which results in the halogen being substituted in the molecule
                • Curly arrow mechanism
              • Undergo elimination reactions - water is eliminated = alkene
            • OXIDATION - P = aldehydes by distillation using H2SO4 and K2Cr2O7 and then carboxylic acids by reflux
              • S = ketones only using acidified dichromate
        • They can be Primary, Secondary or Tertiary
          • P - OH attached to C that is attached to 1 other C
          • S - OH attached to C that is attached to 2 other C
          • T - OH attached to C that is attached to 3 other C
      • Haloalkanes
        • Alkane with at least one halogen group
          • CFCs - C, F and Cl only - Cl radicals break down ozone layer in free radical substitution
        • C-halogen bond is polar  due to electronegativ-e nature of halogens
        • Can be identified by using AgNO3
        • Undergo nucleophilic substitution which results in the halogen being substituted in the molecule
          • Curly arrow mechanism
        • Hydrolysis - hydrolysed by warm aqueous alkali NaOH or water in nucleophilic substitution reaction
    • 9- Organic Synthesis and Analysis
      • IR Spectroscopy
        • IR absorbed by organic molecules' covalent bonds resulting in vibration  - different bonds vibrate more or less - wavenumber is used
        • Can be applied in numerous ways i.e. breathalysers or pollutant monitors
      • Mass Spectrometry
        • Mr and empirical formulae can be found with this info from M+ ion peak

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