C4- The Periodic Table


The History of the Atom

  • At the start of the 19th century John Dalton described the atoms as spheres and said that diferent spheres made up of the different elements
  • JJ Thompson concluded from his experiments that atoms weren't solid spheres, his measurements of charge and mass showed that an atom must contain even smaller, negatively charged particles- electrons.
  • Ernest Rutherford fired positively charged particles at a thin piece of gold. They were expecting most of them to deflect back but most of them past throguh the gold and some deflected backwards. This made him come up with the theory of the nuclear atom. In this there is a tiny nucleus at the centre surrounded by a 'cloud' of negative electrons
  • Niels Bohr proposed a new mode of the atom where all the electrons were contained in shells. He suggested that electrons only exist in fixed orbits and each shell has a fixed energy. His theory was supported by many experiments
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  • The Nucleus: It's in the middle of the atom, it contains protons and neutrons, it has a positive charge and almost the whole of the mass of the atom is concentrated by the nucleus
  • The Electrons: Move around the nucleus in electron shells, negatively charged, they are tiny, the volume of their orbits determines the size of the atoms, have virtually no mass.

The number of protons = the number of electrons

  • The mass number is the total number of protons and neutrons (the big number)
  • The atomic number is the number of protons (the little number)
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Elements and Isotopes

  • There are around 100 elements
  • The modern periodic table shows the elements in order of ascending atomic number
  • It is laid out with elements with similar properties form columns
  • The vertical columns are called groups. Group 1 elements have one electron on their outer shell, etc....
  • The rows are called periods and each period represents another full shell of electrons
  • The period to which the element belongs corresponds to the number of shells of electrons it has


  • They have the same atomic number but different mass numbers
  • If they had different atomic numbers, they'd be different elements altogether
  • A very popular pair of isotopes are carbon-12 and carbon-14
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History of the Periodic Table

Dobereiner started to list elements into groups based on their propetes, he put them into groups of three which he called triads. The middle element of each had a relative atomic mass which was an average of the other two.

A new person called Newlands tried to arrange them more usefully and he noticed that every eigth element had similar properties. The pattern eventually broke down and this was because he left no gaps so his work was ignored.

Dmitri Mendelev arranged the table of elements with various gaps, he put them in order of atomic mass and he left gaps in the first two rows before the transition metals come in on the third row. The gaps then predicted the properties so far

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Electron Shells

  • Electrons always occupy shells
  • The lowest energy levels are always filled first
  • Only a certain number of electrons are allowed in each shell

Electron Configuration

  • The first shell has 2 electrons
  • Then each shell after that has up to 8 electrons
  • So the configuration for aluminium is 2,8,3 which adds to 13 and its atomic number
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Ionic Bonding

In ionic bonding atoms gain or lose electrons to form charged particles or ions which are then strongly attracted to one another.

If they have 1 or 2 electrons on the outer shell it is easier to lose electrons to make them positive. Elements with 6 or 7 electrons on the outer shell it is easier to gain electrons to fill the shell up they become negative.

Ionic bonds between metals and non-metals and always produce the giant ionic structures. The ions form a closely packed regular lattice arrangements and they aren't free to move though so they do not conduct electricity when solid. There are very strong chemical bonds between all the ions

If you dissolve certain elements like MgO or NaCl you make their ions free to move and they can conduct electricity. When dissolved the ions seperate and are free to move, they can carry electric current

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Ions and Ionic Compounds

  • Ions are charged particles- they can be single atoms or groups of atoms
  • When they lose or gain electrons to form ions all they are trying to do is get a full outer shell
  • When metals frm ions, they lose electrons to form positive ions
  • When non-metals form ions, they gain electrons to form negative ions
  • So when a metal and a non-metal combine they from ionic bonds
  • If two electrons are lost its charge is 2+, if three electrons are lost its charge is 3-
  • Dot and Cross diagrams show what happens to the electrons in ionic bonds
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Covalent Bonding

  • When non-metal atoms combine together they form covalent bonds by sharing pairs of electrons
  • This way both atoms feel that they have a full outer shell and that makes them happy
  • Each covalent bond provides on extra shared electron for each atom
  • Each atom involved has to make enough covalent bonds to fill up its outer shell
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Group 1- Alkali Metals

  • As you go DOWN group 1, the alkali metals become MORE reactive, the outer electron is more easily lost because it's further from the nucleus so less energy is needed to remove it. The alkali metals all have 1 outer electron which makes them very reactive.
  • They all have: a low melting point and boiing points, they have a low density and are very soft.
  • Group 1 metals are keen to lose an electron to form a 1+ ion with a stable electronic structure. The more reactive the metal the happier it is to lose an electron. This is called OXIDATION
  • When lithium, sodium or potassium are put in water they move around the surface fizzing furiously and produce hydrogen. The reactivity with water increases down the group- the reaction with potassium. Sodium and potassium melt in the heat of the reaction, an alkali forms which is the hydroxide

Lithium produces a red flame

Sodium produces a yellow flame

Potassium produces a lilac flame

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Group 7- Halogens

As you go DOWN group 7, the halogens become LESS reactive, there's less inclination to gain an electron to fill the outer shell when it is further out from the nucleus. As you go down group 7 the melting points and the boiling points of the halogens increase

Halogens are keen to gain an electron to form a 1- ion with a stable eletronic structure. The more reactive the halogen the happier it is to gain an electron. This is called REDUCTION

Halogens react vigorously with alkali metals to form salt called 'metal halides'

2Na + Cl2 -> 2NaCl

Sodium + Chlorine -> Sodium Chloride

Chlorine can displace bromine and iodine from a solution of bromide or iodide. Bromine will also displace iodine, you could be asked to predict the results of displacement reactions using other halogens- just remember more reactive halogens displace less reactive

Chlorine + Potassium Iodine -> Iodine + Potassium Chloride

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  • Metals have a crystal structure and they all have the same basic properties. These are due to a special type of bonding that exists. They are held together with metallic bonds and these allow the outer electrons to move freely. This creates delocalised electrons throughout the metal which is what gives rise to many of the properties of the metal.
  • Metals are hard, dense and lustrous. They have a strong attraction between them and are packed with positive ions. They have high melting and boiling points because of these strong bonds. You need alot of energy to break them. They are strong and hard to break but they can be hammered into a different shape.
  • SAUCEPANS: Good conductor of heat, doesn't rust easily. It is make of stainless steel (cheap)
  • ELECTRICAL WIRING: Good conductor of electricity, easily bent. It is made of copper
  • AEROPLANES: Low density, strong and doesn't corrode. It is made of aluminium, titanium is sometimes used but it is much more expensive
  • BRIDGES: They are strong, made of steel which is mostly iron but it has carbon in it which makes it alot less brittle
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Superconductors and Transition Metals

Normally all metals have resistance, even good ones like copper. That resistance means that whenever electricity flows through them, they heat up and some is wasted as heat. If you make some cold enough their resistance disappears completely and these are called superconductors. Without resistance none is turned into heat so none is wasted.

With these you can make...... Power Cables, Strong Electromagnets, Electronic Circuits. Metals only start superconducting at less than -265'c and getting things that cold is really hard

Transition Metals that are everyday metals are: copper, iron, zinc, gold, silver and platinum. They all have metallic properties

Iron is a catalyst used in the Haber Process ad Nickel is useful for the hydrogenation of alkenes.

The compounds of transition elements are colourful due to the metal ion they contain. Iron (II) Compounds are light green, Iron (III) Compounds are orange/brown and copper compounds are often blue

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Thermal Decomposition and Precipitation

Thermal decomposition is when a substance breaks down into at least two other substances when it is heated. Transition metal carbonates break down. They break down into metal oxide and carbon dioxide. Copper (II) Carbonated -> Copper Oxide + Carbon Dioxide

Precipitation reaction is where two soultions react and an insoluble solid forms in the solution. The solid is said to precipitate out and the solid is called a precipitate.

  • copper (II) sulfate + sodium hydroxide -> copper (II) hydroxide + sodium sulfate
  • iron (II) sulfate + sodium hydroxide -> iron (II) hydroxide + sodium sulfate
  • iron (III) sulfate + sodium hydroxide -> iron (III) hydroxide + sodium sulfate

Copper (II) Hydroxide is a blue solid

Iron (II) Hydroxide is a grey/green solid

Iron (III) Hydroxide is an orange/brown solid

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Water Purity

In the UK we get our water from: SURFACE WATER.... Lakes, rivers, resevoirs or GROUNDWATER.... aquifers.

Water is purified in water treatment plants:

  • Filtration: a wire mesh screens out large twigs and then gravel and sand beds filter out any other solid bits
  • Sedimentation: iron sulfate or aluminium sulfate is added to the water, which makes fine particles clump together and settle at the bottom
  • Chlorination: chlorine gas is bubbled through to kill harmful bacteria and other microbes

Tap water can still contain impurities:

  • Nitrate Residues: from excess fertiliser 'run off' into rivers and lakes. Nitrates can prevent the blood carrying oxygen properly
  • Lead Compounds: from old lead pipes, lead is very poisonous
  • Pesiticide Residue: from spraying too near rivers and lakes
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Testing Water Purity

Water companies have to test their water regularly to make sure that the pollutant levels don't exceed strict limits. You can test for dissolved ions very easily using precipitation reactions.

You test for Sulfate Ions using Barium Chloride: 1) add some dilute hydrochloric acid to the test sample. 2) then add 10 drops of barium chloride solution. 3) if you see a white precipitate there are sulfate ions in the sample

Test for halide ions using silver nitrate. 1) add some dilute nitric acid to the test sample. 2) then add 10 drops of silver nitrate solution. 3) if halide ions are present a precipitate will form.

  • Chloride ions will produce a white precipitate
  • Bromine ions will produce a cream precipitate
  • Iodide ions will produce a pale yellow precipitate
  • Silver nitrate + sodium chloride -> silver chloride + sodium nitrate
  • Silver nitrate + sodium bromide -> silver bromide + sodium nitrate
  • Silver nitrate + Sodium iodide -> silver iodide + sodium nitrate
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U should include all the topics, but so far it has been helpfull;





otherwise good!

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