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  • Created on: 06-04-15 13:43

atomic structure

the proton and neutron are within the nucleus which is within the centre of the atom, the elctrons are on the edges of the atom (http://www.bbc.co.uk/staticarchive/e0d8ee053a94b0c6b5ef79fc787edf374ffe4b54.gif)nucleus - middle of atom, contains protons and neutrons, positive charge, almost whole mass of the atom

electrons - move around nucleus in electron shells, negative charge, tiny, virtually no mass

  • Protons : mass 1 charge +ve
  • Neutron : mass 1 charge 0
  • electron: Mass 0.0005 charge -ve

number of protons = number of electrons in an atom

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atomic number and mass number

  • (http://physicsnet.co.uk/wp-content/uploads/2010/06/atomic-number-and-mass-number.jpg)mass no. = total no. of protons and neutrons
  • atomic no. = number of protons

to get no. of neutrons, subtract atomic from mass number

mass is always the bigger number.

on a periodic table, mass is "relative atomic mass" (RAM)

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Development of the Atom

  • John Dalton 1803 - everything was made from tiny spheres (atoms)
    • all matter is made from atoms
    • atoms cannot be made or destroyed
    • all atoms of an element are identical
    • different elements contain different types of atoms
  • J J Thomson 1897 - atoms not solid spheres. atoms contain subatomic particles (electrons). "plum pudding model"
  • Ernest Rutherford 1911 - atom is mostly empty space with electrons arranged around nucleus
    • gold foil experiment - fired +ve particles at thin sheet of gold, if plum pudding model was accurate, most of particles should have been deflected by +ve "pudding". But most particles went straight through and few were deflected so pudding model was wrong
  • Niels Bohr 1913 - electrons moved in fixed orbits or shells around nucleus

theories have to be backed up by evidence

  • peer review - check for errors, see new ideas, develop own work
  • accepted as fitted scientific evidence to support theory
  • new evidence - new improved ideas, used to make predicitons
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isotopes are different forms of the same element, which have the same number of protons but different number of neutrons

if they had different atomic numbers, they would be different elements

  • e.g. carbon-12
    • Six protons (because its proton number, at the bottom, is 6)
    • Six electrons (because the number of protons and electrons in an atom is the same)
    • six neutrons
  • carbon-14
    • Six protons (because its proton number, at the bottom, is 6)
    • Six electrons (because the number of protons and electrons in an atom is the same)
    • 8 neutron
  • when an atom becomes an ion, its nucleus stays the same and so do the numbers of protons and neutrons.
  • an ion with single -ve charge has one more electron then original atom, and an ion with a single +ve charge has one less electron
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Hstory of Periodic Table

  • Döbereiner - triads
    • groups based on chemical propertie
    • groups of 3 - middle element was average of the two
  • Newland - Law of Octaves
    • all elements known at the time in order of RAM; each element was similar to the element eight places further on
    • BUT - groups contained elements without similar properties; mixed up metals and non-metals; didnt leave gaps for undiscovered elements
  • Mendeleev - in order of RAM but left gaps and predicted new elements
    • similar properties in vertical groups
    • later discoveries enforced this (e.g. atomic number exactly 1 more then previous element each time)
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Electron Shells

  • always occupy shells/energy levels
  • lowest energy levels are occupied first
  • 1st shell = 2 electrons, 2nd shell = 8 electrons, 3rd shell = 8 electrons (maximum)


Structure of a fluorine atom. A black dot represents the nucleus. The small circle around this has two red dots on it, representing the first energy level with two electrons. A larger outer circle has seven red dots on it, representing the second energy level with seven electrons (http://www.bbc.co.uk/staticarchive/9aeb6abcdb100140b628cbc96ce4d6fa9c614580.gif)

flourine, 2,7

group 7 (number in outter shell determines group);

9 electrons in total

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Ionic Bonding

  • ionic bonding atoms lose or gain electrons to form charged particles (and full outter shells) which are strongly attracted to one another (metal and non-metal)
    • Metal atoms lose the electron, or electrons, in their last shell and become positively charged ions
    • Non-metal atoms gain an electron, or electrons, from another atom to become negatively charged ions
      • shell with just one electron will lose it to get full outter shell e.g. sodium Na+
      • shell thats nearly full will gain electron to get full outter shell e.g. chlorine Cl-


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Giant Ionic Lattices

    • ionic bonds form between metals and non-metals and always produce giant ionic structures
    • form a closely packed rgular lattice arrangement. ions are not free to move so do not conduct electricty when solid
    • very strong chemical bonds between all ions
  • example:
    • MgO and NaCl - high melting & boiling point as of very strong attraction between opposite charged ions in giant structure. need a lot of energy to break these bonds
      • MgO made of Mg2+ and O2- ions and NaCl made of Na+ and Cl- ions so MgO has higher melting point as it has double the charge so attraction is harder to overcome
      • O2- are smaller then Cl- so ions in MgO can pack closer together so attraction is harder to overcome so has higher melting and boiling point
      • when MgO and NaCl melt, ions are free to move so they'll conduct electricity
      • NaCl dissolves to form solution that conducts electricty. when dissolved the ions separate and are all free to move in solution so can carry electric current
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Covalent Bonding

  • covalent bonds happen with atoms sharing a pair of electrons to form full outer between two non-metals
  • Simple molecules consist of a small number of atoms joined by covalent bonds. For example, water and carbon dioxide exist as simple molecules.
    • atoms within molecules have strong covalent bonds
    • forces of attraction between molecules are weak
    • low melting and boiling points as they are easily parted
    • usually gas/iquid at room temp
    • dont conduct electricty as there are no free electrons or ions
  • e.g. hydrogen, chlorine

Shows a dot and cross model of a hydrogen electrons. The circle on the left has 1 red dot and the circle on the right has 1 blue cross. They overlap, and the cross and the dot are in the same area, representing a covalent bond. (http://www.bbc.co.uk/staticarchive/90791006e8e16edd68ac5e04bf4adf2142c4d7b6.gif)shows a dot and cross model - two circles overlap. The left circle has 5 red dots and the right has 5 blue croses. The area where they overlap has one red dot and one blue cross in it. This represents a covalent bond (http://www.bbc.co.uk/staticarchive/d31e49d68c17f088998f07aa0ec57fb24c200251.gif)

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Group 1 - Alkali Metals

  • lithium, sodium, potassium, (rubidium, caesium, francium)
  • as you go down, becomes more reactive - outer electron more easily lost as its further from nucleus so less energy needed to remove it
    • oxidation - loss of electrons; losses electrons to form full shell
    • M ® M+ + e- (m = metal)

    • low melting/boiling point compared to other metals
    • low density - float on water
    • very soft - cut with knife

reaction with water

  • react vigorously in water
  • move around fizzing as hydrogen is produced
  • potassium gets hot enough to ignite
  • Na and K melt in heat of reaction
    • metal + water → metal hydroxide + hydrogen
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Flame Tests

  • Lithium - red flame
  • Sodium - yellow/orange flame
  • potassium - lilac flame


  • dip wire into HCl to clean and mositen it
  • put loop into powdered sample of compound to be tested, then place end in a blue bunsen flame
  • alkali metal ions will give coloured flame
    • colour of flame tells you which alkali metal is present

~ cannot use Rubidium or ceasium in school as EXTREMELY reactive. when wet they explode ~

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Group 7 - Halogens (salt makers)

  • flourine, chlorine, bromine, iodine, astatine
  • as you go down, they become less reactive - less inclination to gain extra electron to fill outer shell as its further from nucleus
    • as you go down, melting/boiling points increase so at room temp:
    • chlorine - fairly reactive, poisonous, dense green gas (low boil point) - sterilise water
    • bromine - dense, poisonus, orange liquid - making pesticides
    • iodine - dark grey crystalline solid (high boiling point), when warmed it changes to purple vapour - sterilising wounds
  • reduction - gain electrons, the more reactive more likely to gain electron
    • Ha2 + 2e- - 2Ha-

  • halogens with alkali metals to form salts
    • react to form salts called metal halides
    • e.g. sodium + chlorine - soium chloride
    • general balanced symbol equation:  M=alkali metal Ha =halide
      • 2M + Ha2 -  2MHa
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Displacement reactions

more reactive halogens will displace less reactive ones

  • Flourine more reactive then chlorine
  • chlorine more reactive then bromine
  • bromine more reactive then iodine
    • e.g. chlorine + sodium bromide - bromide + sodium chloride (chlorine pushes out (displaces) bromine from sodium bromide solution.
    • chlorine will displace bromine from metal bromides and iodides from metal iodides
    • bromine will displace iodine from metal iodides
  • reaction will not work other way around (e.g. bromine added to sodium cloride)
  • flourine most reactive so will displace all grp7 elements
  • astatine least reactive so cannot displace any of grp7 elements (everything displaces astatine)
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Transition Elements

  • The transition elements are those in the middle section of the Periodic Table.
  • All the transition elements are metals and so they have typical metallic properties
    • they conduct heat and electricity
    • they are malleable and ductile 
    • form positive ions when they react with non-metals.

The compounds of transition metals are often coloured.

  • Copper compounds are blue
  • Iron(II) compounds are light green
  • Iron(III) compounds are orange/brown
    • Iron is a catalyst in the Haber process
    • Nickel is a catalyst used in the manufacture of margarine (hydrogenation)
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Thermal Decomposition

  • thermal decomposition - one substance breaks down to form 2 or more other things when heated
    • general word equation : metal carbonate - metal oxide + carbon dioxide
  • transition metal carbonates break down on heating e.g. anything with CO3 at end
  • usually there is a colour change
    • e.g. copper (II) carbonate
    • copper(II) carbonate - copper oxide + carbon dioxide
    • goes from green to black
      • e.g. iron (II) carbonate to iron oxide, manganese carbonate to manganese oxide, zinc carbonate to zinc oxide are the same but with different colour change
    • If limewater is shaken with a sample of the gas produced, the limewater turns milky. This shows that the gas is carbon dioxide.
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Precipitation Reactions

  • where 2 solutions react to form an insoluble solid in the solution (the precipitate)
  • some soluble transition metal compounds react with sodium hydroxide to form an insoluble hydroxide
    • Copper hydroxide - blue solid Cu2+ + 2OH- → Cu(OH)2
    • Iron (II) hydroxide - grey/green solid Fe2+ + 2OH- → Fe(OH)2
    • Iron (III) hydroxide - orange/brown solid Fe3+ + 3OH- → Fe(OH)3
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Metallic Structure

  • in general metals are hard, shiny, good conductors or heat/electricity
  • many metals have a high tensile strength so they resist being stretched. also have high melting/boiling points
    • saucepans - conductor of heat, stainless steel
    • electrical wiring - conductor of electricity, bendable, copper
    • aeroplanes - low density, strong - aluminium/titanium(more expensive)
    • bridges - strong, steel - mostly iron but also has carbon making less brittle
  • metals are held togther by metallic bonds
    • contain +ve metal ions packed closely together which form when electrons leave outer shell
    • electrons become free to move within structure
    • form a "sea" of delocalised electrons
    • metallic bonding is the strong force of attraction between sea of delocalised electrons and closely packed +ve metal ions
    • these are strong so need high melting boiling point
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  • some metals become superconductors at very low temperatures
  • these have little reistance to the flow of electricity or even no resistance at all
  • little or no loss of energy from superconducters as they have no resistance
    • used to make powerful electromagnets - MRI scanners
    • could be used for super-fast electronic circuits for powerful computers
    • could replacemetal cables that carry electricty from power stations to homes/offic/factories


  • HOWEVER - only work and low temperatures
    • difficult to get to low temperatures and very expnsive limiting use
    • scientists trying to develop superconductors at room temp (20oC)
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  • demand increases every year. some worry we will run out of water to supply everybody, so we must conserve it.
  • surface water - lakes, rivers, reservoirs. start to run dry in summer months
  • groundwater - aquifiers (rocks that trap water underground)

used as a solvent to dissolve substances, coolant to stop overheating, cheap raw material

  • water contains insoluble materials (leaves), dissolved salts and minerals, microbes and pollutants e.g.:
    • nitrates from fertiliser that runs of fields
    • lead compounds from lead pipes
    • pesticides from spraying crops in field too close to water resources.
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Purifying water

Water must be purified before use:

  • Filtration - The water is sprayed onto specially prepared layers of sand and gravel. As it trickles through, different sized insoluble solids are removed. The filter beds are cleaned periodically by pumping clean water backwards through the filter.
  • Sedimentation - A chemical is added which causes tiny solid particles (which would pass through a filter) to clump together into larger particles. These can then be allowed to settle out or may be filtered
  • Chlorination - Chlorine gas, injected into the water, kills microbes.
  • some soluble substances are not removed during purificaion. these could be poisonous, so strict limits are placed on how much of them can be in water
  • in dry countries sea water is distilled to produce drinking water
    • needs a lot of energy so is expensive
    • therefore not prectical to make in lare quantities
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Testing Water

  • Sulfate ions using Barium Chloride
    • add dilute HCl to test sample, then 10 drops of barium chloride
    • if white precipitate, there are sulfate ions
  • Barium chloride solution + sodium sulfate solution → sodium chloride solution + (solid) barium sulfate
    • BaCl2(aq) + Na2SO4(aq) → 2NaCl(aq) + BaSO4(s)
  • Halide Ions using Silver Nitrate
    • add dilute nitric acid to sample, then 10 drops of silver nitrate
  • Silver nitrate solution + sodium bromide solution → sodium nitrate solution + solid silver bromide
    • chloride ions - white precipitate
    • AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s)
    • bromide ions - cream precipitate
    • AgNO3(aq) + NaBr(aq) → NaNO3(aq) + AgBr(s)
    • iodide ions - pale yellow precipitate
    • AgNO3(aq) + NaI(aq) → NaNO3(aq) + AgI(s)
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