C2

C2

1. Periodic Table

   • Early ideas
     • Early 1800s - elements arranged by atomic mass (scientists had no idea about atomic structure, protons, neutrons or electrons so they didn't have an atomic number at this point)
       • Early tables were incomplete and some elements were in the wrong place as they were placed based on their mass instead of their properties
         • 1869 - Mendeleev overcame some of these problems by taking the 50 known elements and arranging them in his table of elements & left gaps
           • He put elements mainly in order of atomic mass, but did switch that order if the properties meant it should be changed
             • Gaps were left as not all of the elements had been discovered - this made sure that elements with similar properties stayed in the same groups. when new elements were discovered and they fitted the pattern, it helped confirm Mendeleev's ideas
   • Modern Periodic Table
     • Helps you see patterns in properties
       • About 100 elements - each arranged with increasing atomic (proton) number
         • The table makes it easier to remember what elements have what properties as they are based in groups and periods. Also makes it easier to establish metals and non-metals
           • Metals = left
           • Non-metals = right
           • Elements with similar properties form columns = groups
             • Group number tells us how many electrons are in the outer shell. Eg. Group 1's elements have 1 e- in outer shell
               • Reactivity of element is based on about of e- in outer shell therefore explaining why all elements in the groups react in a similar way
           • Rows in the periodic table = periods. Each new period represents how many shells the elements have
   • Groups of the Periodic Table
     • Group 0
       • Noble Gases. All inert. Colourless gases
         • Have full outer shells = unreactive
           • BPT increases as move down group - due to greater intermolecular forces between increasing amount of electrons
             • He, Ne, Ar, Kr, Xe, Rn
     • Group 7
       • Halogens - Non-metals, coloured vapours
         • F - very reactive, poisonous yellow gass
           • Cl - fairly reactive, poisonous dense green gas
             • Br - dense, poisonous, red-brown volatile liquid
               • I - dark grey, crystalline solid or purple vapour
                 • All exist as molecules which are pairs of atoms
                   • As you go down the group: become less reactive, have higher MPT & BPT, have higher atomic masses
                     • Can form Molecular Compounds with simple molecular structures
                       • Can share e- via covalent bonding with other non-metals to achieve a full outer shell
     • Group 1
       • Alkali metals - Reactive, soft metals, low density
         • As you go down the group: increasing reactivity - the lone e- in the outer shell doesn't have a strong force of attraction to the nucleus as it is further away, therefore is easily lost
           • Lower MPT & BPT and higher atomic masses
             • Form Ionic Compounds with Non-metals
               • Usually compounds are white solids that dissolve into water to form colourless solutions

2. Discovery of the Atom

   • Initial ideas
     • 19th C, Dalton described atoms as solid spheres, and the different spheres made up the different elements
       • 1897 - JJ Thompson concluded the atoms were not solid spheres, his measurements of charge and mass shows the atom must contain even smaller, negatively charged particles = electrons
         • New theory = Plum Pudding Model -> + charged ball with - charged e- stuck in it
           • Rutherford showed this model idea to be WRONG
             • 1909 - alpha particle scattering experiment - + charged alpha particles fired at thin gold foil
               • Prediction: particles pass straight through sheet or be slightly deflected at most
                 • Results: Most particles did go straight through, BUT some were deflected more than expected, and a small number were deflected backwards - plum pudding model couldn't be right
                   • Rutherford came up both the nuclear model of the atom -> + charged nucleus where most of atoms mass was concentrated, electrons in a 'cloud' around the nucleus = most of atom = empty space
                     • + charged nucleus and + charged alpha particles = repel = deflection. if alpha particles were fired directly at the nucleus, this is why they deflected backwards
                       • Bohr' explained that if there was a 'cloud' of electrons, it would be attracted to the nucleus causing the atom to collapse
                         • Bohr's model suggested that all electrons were contained in shells and that they orbit the nucleus in fixed shells - each shell a fixed distance from the nucleus
                           • 20th C - Discovery of neutrons by James Chadwick, resulting in a model of the atom which was very close to the modern day version known as the nuclear model
                     • Further experimentation from Rutherford helped conclude that the nucleus could be divided into smaller particles - each of which has the same charge as a hydrogen nucleus = protons
   • The Atom
     • atomic number = number of protons in element / number of electrons in element
       • p = e
         • SMALLER number on periodic table
     • Mass number of an atom = protons + neutrons
       • p + n
         • BIGGER number on periodic table
     • Isotopes
       • Atoms of the same element with same number of protons but a different number of neutrons = different mass number
         • Eg. Cl 35 17 Cl 37 17

3. Transition metals

   • Properties
     • Strong, hard, high densities, high MPT & BPT (except for mercury = liquid @ room temp)
       • Less reactive than alkali metals... copper, silver and gold do NOT react with O2 or H2O @ room temp = good for pipes and jewellery
         • Transition metals = CATALYSTS very important as they speed up reactions
     • Often coloured so the compounds that contain them are also coloured
       • Good conductors of heat and electricity as have the properties of a 'proper metal'

4. Metal + Non-metal Compounds

   • When a metal loses an e- in the outer shell and a non-metal gains this e- (ionic bonding) this gives both a full outer shell
     • The metal becomes a + ion as there are now more protons than electrons
     • The non-metal becomes a - ion as there are now more electrons than protons
     • Ionic compounds and their properties linked to their structure
       • form giant ionic lattice structures, strong bonds due to strong forces of attraction between + and - ions, high MPT & BPT
         • When in solid state, DO NOT conduct electricity as no electrons are free to move so can't carry charge
           • However, when dissolved / in liquid state, CAN conduct electricity as electrons are now free to move so can carry charge
             • All ionic compounds dissolve easily in water