Elements of Life

1. Bohr Model
3. Isoptopes

   Annotations:

   • Isotopes are where you get the same element with the same number of protons but a different number of neutrons 
   1. Relative Atomic Mass = sum of (isotopic mass x Isotopic abundance)/100
4. Electron configuration

   Attachments:

   • Basic Electron Configuration (EL)
5. History
   1. Ancient Greece

      Annotations:

      •    Ancient Greeks hypothesised that matter was made of indivisible particles called atoms   
   2. JJ Thompson

      Annotations:

      • JJ Thompson theorised that atoms contained tiny charged particles called electrons that spread out in a positive ‘dough’ which he called the plumb pudding model
   3. Ernest Rutherford

      Annotations:

      •    Rutherford used the gold leaf experiment to disprove the plumb pudding model. This proved the atom contained a small dense positively charged nucleus surrounded by a cloud of electrons   
   4. James Chadwick

      Annotations:

      •    James Chadwick discovered the existence of the neutron which inspired a whole new branch of nuclear physics   
   5. Neils Bohr

      Annotations:

      •    Neils Bohr developed the idea that the electrons exist in fixed energy levels called shells   
   6. John Dalton

      Annotations:

      •    John Dalton thought of atoms as solid spheres with elements been one type of atom and compounds as a mixture of elements bonded together   

Attachments:

• GCSE Bonding Diagrams

1. Ionic Bonding

   Annotations:

   • Metals lose electrons to get a stable full outer shell becoming positive ions Non-metals gain electrons to get a stable full outer shell becoming negative ions Positive and negative ions attract by an electrostatic attraction
   1. Very high melting point

      Annotations:

      • A lot of strong ionic bonds in a giant ionic lattice. This requires a lot of energy to break these strong bonds
   2. Often soluble in water

      Annotations:

      • Polar water molecules can dissociate the ions causing them to dissolve
   3. Conduct electricity as a liquid or dissolved in water

      Annotations:

      • When molten or aqueous the ions are free to move and can carry charge As a solid the ions are fixed in position and can not carry charge
2. Covalent Bonding

   Annotations:

   • 2 non-metals share electrons to achieve a stable full outershell
   1. Simple molecular
      1. Non conductor of electricity

         Annotations:

         • No free electrons or ions to carry the charge
      2. Low melting point

         Annotations:

         • Only need to break weak intermolecular forces  to move the molecules apart which doesn't require a lot of energy
   2. Giant Covalent
      1. Very High melting point

         Annotations:

         • Have lots of strong covalent bonds holding the giant structure together. This requires a lot of energy to overcome these strong bonds
      2. Variable electrical conductivity

         Annotations:

         • Graphite and fullerenes conduct electricity as carbons 4th outer shell electron is delocalised and free to mover and carry charge
3. Metallic bonding

   Annotations:

   • Metal atoms lose electrons to gain a full outer shell, becoming positive ions. The positive ions are arranged in rows with a sea of negative delocalised electrons surrounding them. The positive ions are held together by negative electrons by an electrostatic force (positive attracts to negative)
   1. Conduct electricity

      Annotations:

      • Delocalised electrons are free to move and carry charge
   2. Conducts heat

      Annotations:

      • Delocalised electrons are free to move and carry thermal energy (vibrations)
   3. High melting point

      Annotations:

      • Ions held together by strong electrostatic forces which require energy to overcome

1. The Mole

   Annotations:

   • One mole of anything contains the same number of particles (atoms, molecules, ions, electrons etc.) than there is in 12g of the isotope \(^{12}C\)
   1. Avogadros constant

      Annotations:

      • The number of particles (atoms, molecules, ions, electrons etc.) in 1 mole = \(6.02 X 10^{23}\)
   2. moles = mass/molar mass

      Annotations:

      • molar mass can be relative atomic mass or molecular mass
   3. Gas volumes

      Annotations:

      • One mole of an ideal gas occupies \(24dm^3\) at standard conditions
2. Empirical Formula

   Annotations:

   • Worked Example A phosphorous oxide contains 6.2g of phosphorus and 8g of oxygen, calculate the empirical formula of the oxide.  Step 1: Divide by the atomic mass to calculate the moles  P = 6.2/31 = 0.2           O = 8/16 = 0.5 Step 2: Divide by the smallest number P = 0.2/0.2 = 1              O  = 0.5/0.2 = 2.5 Step 3: Either round the numbers or multiply to get whole numbers  P = 1 x 2 = 2                       O  = 2.5 x 2 = 5  Step 4: Write the empirical formula  \(P_2O_5\)   
3. %Composition

   Annotations:

   • % mass composition of an element is the total mass of an element in a compound divided by the total molecular mass of the compound times by 100
4. Chemical Equations

   Annotations:

   •    Key skills include:   - Balancing symbol equations   - using state symbols (g,s,l & aq)   - Use mole ratios in calculations   

1. Table
2. Nuclear Decay Equations

   Annotations:

   • Make sure the nuclear decay equations balance for both atomic number and atomic mass.
3. Half Lives

   Annotations:

   • The time it takes for the activity or mass of a radioactive isotope to halve
   1. Tracers

      Annotations:

      •    Medial tracers don't want a half-life too small so that it would decay and be undetectable. They also shouldn’t have a half-life too long as it would remain in the body and potentially ionise DNA increasing the risk of cancer.   
   2. Dating

      Annotations:

      •    When using radioactive dating the half-life shouldn’t be too long, so that less than one half-life has elapsed, as any decay would be undetectable. Also if the half-life is too short the isotope would decay too much and any activity would be undetectable.   

1. Emission
2. Absorption

Attachments:

• Shapes of Molecules

1. Electron repulsion theory

   Annotations:

   •    Model answer There are ____ sets of electrons around the central ___ atom. Electrons are negative and repel and try and get as far apart from them as possible. There are ____ bonding pairs and ____ lone pairs (only if necessary). The shape of the molecule is __________ and the bond angle is ________.   

1. History of the periodic table
   1. Newlands

      Annotations:

      • Arranged elements in order of mass and grouped as octaves as every eighth element had similar properties. However after the first couple of rows the pattern of reactivity didn't work
   2. Mendeleev

      Annotations:

      • Also arranged table by mass but forced elements into groups with similar reactivity. He did this by leaving gaps for undiscovered elements and swapping the elements iodine and tellurium which didn't match reactivity with others in the group if ordered by mass.
   3. Modern

      Annotations:

      • The modern periodic table is ordered by proton number and includes the noble gases which Newland and Mendeleev didn't know existed. The table also matches electronic structure.
2. Group 2
   1. Solubility

      Annotations:

      • As you go down the group the hydroxides become more soluble and hence more alkali. As you go down the groups the carbonates become less soluble
   2. Reactivity with water

      Annotations:

      • Group 2 metals react with water to form hydroxides and hydrogen gas. \(Ca + 2H_2O => Ca(OH)_2 + H_2\) Group 2 increases in reactivity as you go down the group as the outershell is further from the positive nucleus with more shielding so the outershell electrons are easier to lose
   3. Basic Oxides

      Annotations:

      • Group 2 oxides and hydroxides act as bases which will neutralise acids \(CaO + 2HCl => H_2O + CaCl_2\) \(Ca(OH)_2 + 2HCl => 2H_2O + CaCl_2\) (a base is a proton/hydrogen ion acceptor)
   4. Thermal decomposition

      Annotations:

      • A group 2 carbonate will thermally decompose forming an oxide and carbon dioxide \(CaCO_3 => CaO + CO_2\) Thermal stability of the carbonates increases as you go down the group. This is because as you go up the group  as charge density on the cation increases which polarises the carbonate anion more making it less stable.
3. Periodicity: Trend in boiling points across a period

   Annotations:

   • As you look across the group the trend matches the bonding physical properties. Firstly metallic bonding with bp increasing with the amount of electrons lost to form the positive ion Next the graph peaks with a giant covalent structure Finally the bps drop as elements form simple molecular covalent bonding