Periodic Table

Periodic Table

1. The Atom

   • stages involved in developing current ideas
     • Democritus
       • suggestion that all things are made out of particles
       • 400 BC
     • Dalton
       • atomic theory
         • atoms of the same elements are all alike
         • combined to make compounds
         • named particles "atoms" as they could not be split
     • Thomson
       • discovery of the electron
       • "plum pudding theory"
         • atoms were balls of positive charge with tiny negative particles stuck inside
     • Rutherford
       • discovers proton and nucleus
     • Bohr
       • new model = valence shells, nucleus in the middle
     • Chadwick
       • discovers neutrons
   • Isotopes
     • must have same proton number but different mass numbers
       • e.g Carbon 12 and Carbon 14
         • both has 6 protons, yet C12 has 6 neutrons whereas C14 has 8 neutrons
     • if different proton numbers, different elements altogether
   • Relative Atomic Mass
     • average masses of all isotopes in elements
       • rel. mass of each isotope and its rel. abundance
         • re. abundance = how much there is of each isotope compared to total amount of its element in the world
     • A^r formula:
       • (rel. mass of iso. x rel. abundance) /relative abundance
         • e.g chlorine isotopes C35, rel. abundance = 3 and C37, rel. abundance = 1
           • Ar = (35x3)+(37x1)/(3+1) = 35.5
     • mass of an atom in atomic mass units is relative to carbon-12. Only C12 has a mass exactly equal to itself.
       • why atoms don't have whole numbered atomic mass
   • differences between elements, mixtures and compounds
     • elements
       • purest form of chemical substance, composed of one type of atom
       • groups of atoms are called molecules (H2, O2 etc.)
     • compounds
       • substances made of two or more different elements that are chemically joined
         • different types of atoms are allowed
         • difficult to separate the o.g elements again
         • properties of compound are different to the properties of o.g elements
     • mixtures
       • no chemical bound between mixtures
         • easily separable through physical methods
       • properties of mixture = properties of all sep. parts

2. Periodic Table

   • Groups
     • 1. alkali metals
       • most reactive group of metals
         • e.g Lithium, Sodium, Potassium etc.
       • low melting points
       • low densities
     • 2. alkali earth metals
       • compounds are common in rocks of Earth's crust
       • less reactive than group 1
         • e.g Magnesium, Beryllium, Calcium etc.
     • transition metals
       • 'typical' metals
         • hard, dense, shiny, good conductors of heat and electricity, malleable and ductile
       • less reactive than gp. 1 &2
       • forms colored compounds
     • 7. halogens
       • all form diatomic molecules
         • e.g F2, Cl2, Br2 etc.
       • becomes less reactive as you go down the group
     • 8. noble gases
       • very unreactive
         • stable = full valence shells
       • found as the atoms themselves
         • monatomic - 'one atom'
       • boiling points increase as you go down the group
         • forces between molecules get stronger as the molecules get larger
         • density also increases
   • Metals
     • properties
       • dense
       • high melting point
         • except alkali metals: low densities and melting points
       • ductile
       • malleable
       • shiny
       • good conductors
         • electricity
         • heat
       • sonorous
       • how it can be used to decide how to use it (examples)
         • e.g 1: copper
           • good conductor of electricity, is ductile and strong, can be bent, doesn't react to water = ideal to be drawn out into wires
         • e.g 2: aluminium
           • low density, malleable, corrosion resistant = ideal for aeroplanes
         • e.g 3: titanium
           • low density, very strong, corrosive resistant = ideal for replacement hips
       • can corrode when exposed to air and water
       • metal fatigue - breaking
     • metallic structures consist of meat ions and free moving valence electrons. As the metal ions are of positive charge, the opposite attract, creating strong metallic bonds - strong electrostatic attractions
       • free electrons allow carrying electricity and heat
         • also allows layers of atoms to slide over each other

3. Bonds

   • Ionic bondings
     • bonds between non metals and metals
     • used to ensure atoms have full outer shells
     • opposite ion charges strongly attracts - creates an ionic bond
     • ionic compounds - giant ionic structures
       • giant 3D lattices
       • positive and negative ions placed alternatively
         • strong electrostatic attractions is very strong
           • requires high levels of energy to break apart the strong bonds
           • high melting and boiling points
         • cations = positive
         • anions = negative
     • dissolves in water
   • Covalent Bonds
     • sharing of electrons to achieve a full outer shell
     • simple molecular structures
       • atoms form bonds to form molecules
         • forces between the molecules are very weak
           • weak intermolecular bonds
           • low melting and boiling points
             • breaks apart molecules, not the covalent bonds between atoms
     • giant covalent structures
       • similar to ionic lattices but there's no charged ions
         • only strong covalent bonds
         • do not conduct electricity
     • between non-metallic atoms

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