Unit 1 - Energy Matters

Reaction Rate
1. The Course of a Reaction
  1. as the reaction proceeds the concentration of products increases & the concentration of reactants decreases
2. Factors Affecting Rate
  1. Collision Theory
    1. molecules must collide successfully for a reaction to take place
      1. particles must collide with a certain minimum energy (activation energy) in order to collide successfully
        if activation energy isn't reached particles will collide unsucessfully
  2. Surface Area
    1. increased surface area (smaller particles); increased no. of collisions
  3. Concentration
    1. increased no. of particles in a fixed volume
      1. more particles collide; increased reaction rate
        rate at the beginning of the reaction is greatest due to higher concentration of reactants
        as reactants are converted into products, the concentration of reactants decreases; no. of successful collisions decreases
  4. Temperature
    1. measure of the average kinetic energy of a substance
      1. kinetic energy of molecules in a reaction varies
        most particles will have kinetic energy near the average nut a small percentage will have values significantly more of less than the average
3. Energy Distribution Graphs
  1. lower temp: curve moves to the left & is higher
    1. higher temp: curve moves to the right & is lower
  2. photochemical reactions are ones in which light is used to increase the no. of particles with energy greater than the activation energy e.g. photosynthesis
4. Catalysts
  1. a substance which increases the rate of reaction & remains chemically unchanged at the end of the reaction
    1. lowers the activation energy
  2. heterogeneous catalysts - in a different state to the reactants
    1. homogeneous catalysts in the same state
    2. a catalyst provides an alternate reaction pathway which results in a lower activation energy
      1. reactant particles are adsorbed onto the active sites of the catalyst
        an activated complex is formed
        the products leave the catalyst , leaving it free for more reactants to adsorb onto it
  3. Enzymes are biological catalysts
    1. catalyse many reactions inside the cells of animals & plants
      1. they are specific
        rennin in cheese production
  4. Catalytic Poisoning
    1. if a molecule other than a reactant is preferentially adsorbed onto the active sites, the sites are no longer available
      1. this poisoning reduces the activity of the catalyst
        industrial catalysts can be poisoned by impurities in reactants
        can be removed & replaced
        or sometimes impurities can be removed from active sites
        reacting with gas
        e.g. catalyst in catalytic cracking can be poisoned by carbon, the carbon is removed by burning it off in a supply of air
        the catalyst is removed, regenerated & returned to the catalyst chamber
        lead poisons catalysts in catalytic converters
Patterns in the Periodic Table
1. Mendeleev
  1. similar chemical properties in groups, in order of increasing atomic number, left gaps for undiscovered
    1. Across a Period
      1. nuclear charge increases, same electron shell is filled up, the attraction between the nucleus & the electrons increases due to increasing size of the positive nuclear charge
        Down a Group
        an additional electron shell is added, the attraction of the positive nuclear charge to the outer electrons decreases due to electron shielding of the nucleus by the additional electron shells
        Atomic Size
        decreases across a period, the larger size of the nuclear charge the greater the attraction between it & the electrons resulting in a greater pull towards the nucleus & a decrease in atomic size
        increases down a group due to the increase in occupied electron shells
        BP, MP & Density
        MP & BP increase across period (groups I to V) & then drops
        MP & BP decrease down group I (decreasing attraction between atoms) but increase down group 8 as van der Waals attractions increase with size
        density increases down all groups, increases across periods until group 3 then decreases
        Ionisation Energy
        the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms
        across a period increases due to increasing nuclear charge
        down a group, decreases due to decreasing attraction between outer electrons & nucleus
        Electronegativity
        measure of attraction an atom involved in a covalent bond has for the shared electron in that bond
        increases, period: decreases group
Enthalpy
1. every substance contains energy known as enthalpy
  1. during reactions, reactants are used up & products made - change in enthalpy going from reactants to products
    1. the enthalpy change (/\H) is the energy difference between reactants & products
2. Enthalpy Changes
  1. energy is required to break the bonds in reactant molecules (activation energy)
    1. when reactant molecules collide they forma highly energetic arrangement called an activated complex
      1. the activation energy is required by molecules to form an activated complex
        this highly energetic complex loses energy & can either go on to from products or fall back into the original reactnats
        energy is released when new bonds are formed in the products
3. Potential Energy Diagrams
  1. Endothermic; energy of products greater than reactants, therefore enthalpy change is positive (+ve)
    1. Exothermic: energy of products less than reactants, therefore, enthalpy change is negative (-ve)
4. Using Potential Energy Diagrams
  1. forward reaction: Ea = highest energy - starting energy, /\H = end energy - starting energy
    1. the same but everything from other side of graph e.g. starting energy fro forwards is ending energy for backwards
5. Catalysts
  1. lowers activation energy
    1. therefore more particles have energy required to overcome activation barrier so more product is formed
      1. no effect on enthalpy change
6. Equations
  1. Combustion
    1. 1 mole of fuel must be balanced
  2. Solution
    1. 1 mole of solute balanced
  3. Neutralisation
    1. 1 mole of water balanced
Bonding,
1. Metals
  1. Metallic Bonding
    1. outer electrons move freely between neighbouring atoms
      1. these electrons are said to be delocalised
        this leaves behind a positive core
        the metal is held together by the strong forces of attraction between positive cores & negatively charged delocalised electrons
  2. Structure
    1. giant lattice
      1. high melting & boiling point due to strength of the metallic bond
        strength of bond varies from metal to metal depending on the number of electrons that each atom 'delocalises' & also on how the atoms pack together
  3. Ionic Bonding
    1. occurs between metals & non-metals
      1. electrons are transferred from the metal to the non-metal in order for them to each have a full outer shell
    2. the ionic bond Is the electrostatic force of attraction between the positively charged metal ion & the negatively charged non-metal ion
  4. Ionic structure
    1. giant lattice of oppositely charged ions
    2. high melting & boiling points because it requires a lot of heat to break millions of strong ionic bonds
2. Non-Metals
  1. Covalent Bonding
    1. sharing electrons in order to achieve full outer shells
    2. the electrostatic attraction between both positively charged nuclei & the negatively charged shared pair of electrons
    3. Pure Covalent Bond
      1. electrons shared equally between atoms
        Electronegativity value will give a measure of the attraction an atom has for the electrons in the bond
        the higher the electronegativity value the more the atom will pull electrons towards itself
        in a pure covalent bond the electronegativity values are identical
    4. Polar Covalent Bond
      1. electrons are pulled closer to one atom giving it a slightly negative charge
        the other atom then has a slightly positive charge
        electronegativity values are different
3. Polar Molecules
  1. symmetrical - nonpolar
  2. non-symmetrical - polar
4. Intermolecular Forces
  1. Van der Waals
    1. v. weak, noble gases are monoamatic so Van der Waals is the only type of bonding they have
    2. (induced) temporary dipoles
      1. electrons spin around the nucleus - sides can have slgithly positive/negative charges
  2. Permanent Dipoles
    1. different electronegativity values - 1 much higher than the other
  3. Hydrogen Bonding
    1. strongest type of intermolecular attraction
    2. occurs when hydrogen is bonded with nitrogen, oxygen or fluorine
      1. these bonds are highly polar due to the different electronegativity values
        hydrogen bonds are electrostatic forces of attraction between molecules which contain these highly polar bonds
        v. high temp. needed to break bonds
Structure
1. Covalent Network Elements
  1. giant lattice of covalently bonded atoms
  2. Diamond
    1. v. high melting point - strong carbon to carbon covalent bonds
      1. v. hard - v. strong bonds
        non-conductor - electrons held tightly between atoms - not free to move
  3. Graphite
    1. high melting point - covalent bonding
      1. conducts electricity - delocalised electrons free to move between the sheets
        van der Waals between sheets - used a s a dry lubricant
  4. carbon, silicon & boron
2. Covalent Molecular Elements
  1. consist of discrete molecules held together by weak intermolecular forces
    1. e.g. fullerene
3. Covalent Network Compounds
  1. 2 main ones: silicon dioxide & silicon carbide
    1. silicon carbide is used as an abrasive
    2. silicon dioxide
Properties
1. Solubility
  1. ionic compounds & polar molecules tend to be soluble in polar solvents (e.g. water) & insoluble in non-polar solvents (e.g. hexane)
  2. non-polar substances tend to be soluble in non-polar solvents & insoluble in polar solvents
  3. like dissolves like
2. Melting & Boiling Points
  1. bond strength decreases down these lists
    1. intramolecular
      1. covalent network (diamond)
        ionic (sodium chloride)
    2. intermolecular
      1. hydrogen bonding (water)
        polar attractions (hydrogen chloride)
        van der Waals (neon)
        water freezes, hydrogen bonding causes molecules to form an open hexagonal shape which results in ice being less dense than water causing it to float on water
The Mole
1. the mass of 1 mole of a substance is its relative formula expressed in grams (GFM)
2. Avogadro's Constant
  1. 6.02 x 10^23

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