# Module C4 - The Periodic Table.

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## The History of the Atom

•  At the start of the 19th century John Dalton described atoms as solid spheres.
• In 1897 JJ Thomson said that an atom must contain negitively charged particles - electrons. This new theory is known as the 'plum pudding model'.
• In 1909 Ernest Rutherford, Hans Geiger and Ernest Marsden conducted the gold foil experiment where they fired positively charged particles at a thin sheet of gold.
• From the 'plum pudding model' they expected most of the particles to be deflected however most particles passed straight through.
• Therefore Rutherford came up with the nuclear atom theory, this theory said that in an atom there was a positively charged nucleus surrounded by a cloud of negative electrons.
• However scientists relised that the cloud would be attracted to the nucleus and the atom would collapse.
• Niels Bohr suggested a new model in which the atom contained shells of electrons. He said that electrons can only exist in fixed orbits and that each shell has a fixed energy. This theory was supported by scientists.
• Due to more experiments and scientific knowledge the Atom theory kept changing untill we got to the theory we have today which is similar to Bohr's model.
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## Atoms

• An atom consists of a nucleus and electrons.
• The nucleus, is in the middle of the atom, contains protons and neutrons, has a positive charge, contains almost the full mass of the atom.
• Electrons, move around the nucleus in electron shells, are negatively charged, have virtually no mass.
• Protons are heavy and positively charged.
• Neutrons are heavy and neutral.
• Electrons are tiny and negatively charged.
• The number of protons in a neutral atom equals the number of electrons. Therefore the atom has no charge because the charges cancel each other out.
• There are two numbers that allow you to know how many of each kind of particle an atom has.
• The mass number tells you the total number of protons and neutrons.
• The atomic number tells you the total number of protons.
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## Elements and Isotopes

• The modern periodic table shows the elements in order of ascending atomic number.
• Elements with similar properties form the following columns, reactive metals, transistion elements, post transition metals, non-metals, noble gases. These columns are called groups.
• The group an element is in depens on the number of electrons it has in its outer shell. For example group 1 elements have 1 electron in their outer shell.
• The rows in the periodic table are called periods each new period represents a new full shell of electrons.
• The period to whcih an element belongs depends on the number of elctron shells it has.
• Isotopes are different forms of the same element, which have the same number of protons but a different number of neutrons.
• Isotopes have the same atomic number but different mass numbers.
• If they had different atomic numbers they would be different elements.
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## History of the Periodic Table

• In the 1800's the only thing that could be measured was relative atomic mass.
• In 1828 Dobereiner started to put elements in groups of three based on their chemical properties. These were called triads, the middle element of each triad had a R.A.M that was the average of the other two.
• In 1864 Newlands noticed that every 8th element had similar properties so he listed the known elements in rows of seven. These sets were known as Newlands' Octaves.
• However there were several things wrong with his work.
• Groups contained elements that didn't have similar properties, e.g. oxygen and iron.
• He didn't leave any gaps for undiscovered elements.
• He mixed up metals and non-metals.
• In 1869, Dmitri Mendeleev arranged the 50 known elements into his Table of Elements.
• Mendeleev put the elements in order of atomic massbut he found he had to leave gaps to keep elements with similar properties in the same vertical groups.
• These gaps predicted the properties of undiscovered elements.
• Later discoveries showed that Dmitri's table made sense.
• Each element has an atomic number exactly one more than the previous element.
• The pattern in the periodic table matches the arrangement of of electrons in an atom.
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## Electron Shells

• Electrons always occupy shells.
• The first shell can only contain 2 electrons but the rest can contain up to 8 electrons.
• You can work out the number of electrons an element because it is equal to the number of protons the element has which you are told in the periodic table.
• Once you know how many electrons an element has you can work out its electronic configuration.
• For example...
• Argon has 18 protons so it must have 18 electrons. The first shell must have 2 electrons, the second shell must have 8 and the third shell will have 8 as well because that is the remaining amount of electrons.
•
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## Ionic Bonding

• In ionic bonding atoms lose or gain electrons to form charged particles, (ions).
• Atoms always want a full outer shell (8 electrons) so they lose or gain electrons to do this.
• If an atom loses an electron it becomes a positively charged ion and if an atom gains an electron it will become a positively charged ion.
• For example...
• Sodium (Na) has 1 electron in its outer shell and Chlorine (Cl) has 7 electrons in its outer shell. Na gives up its outer electron and becomes an Na+ ion. Cl picks up the spare electron and becomes a Cl- ion. They now both have full outer shells.
• Ionic bonds form between metals and non-metals and produce giant ionic structures.
• Ions form a closely packed lattice arrangement where the ions aren't free to move so when solid these compounds don't conduct electricity.
• There are very strong chemical bonds between all the ions.
• Giant ionic structures such as MgO and NaCl have high melting and boiling points because of the strong attraction between oppositely charged ions.
• When MgO and NaCl melt the ions are free to move therefore they will conduct electricity.
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## Ions and Ionic compounds

• Ions are charged particles that can be single atoms (Cl-) or groups of atoms (NO3-)
• When atoms lose or gain electrons to form ions they are trying to get a full outer shell. This is also called a stable electronic structure.
• When metals form ions they lose electrons and become positively charged ions.
• When non-metals form ions they gain electrons and become  negatively charged ions.
• When a metal and a non-metal combine they form ionic bonds.
• The number of electrons lost or gained by an atom is the same number as the charge on the ion.
• For example...
• If 2 electrons are lost the charge is 2+.
• If 3 electrons are gained the charge is 3-.
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## Covalent Bonding

• When non-metal atoms combine together they form covalent bonds by sharing pairs of electrons.
• Substances formed from covalent bonds usually have simple molecular structures.
• The atoms within the molecules are held together by strong covalent bonds.
• The result of weak intermolecular forces is that the melting and boiling points are very low.
• Molecular substances don't conduct electricity because they're are no ions.
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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 is further from the nucleus (it has a larger atomic radius) so less energy is needed to remove it.All alkali metals only have one outer electron so they are very reactive.
• They have low melting points and boiling points, low densities and are very soft.
• Alkali metals always form ionic compounds because they want to lose their outer electron.
• This is becuase they want to form a 1+ ion with a stable electronic structure.
• Loss of electrons is called OXIDATION.
• When alkali metals i.e. lithium, sodium and potassium are put in water they have a strong reaction. They move and fizz on the waters surface while producing hydrogen.
• Some of the more reactive metal such as potassium get hot enough to ignite and sodium and potassium melt in the reaction.
• An alkali forms which is the hydroxide of the metal for example sodium hydroxide NaOH.
• Alkali metals burn with specific colours, Li-red, Na-yellow/orange,  K-lilac.
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## Group 7 - Halogens

• All group 7 elements have 7 electrons in their outer shell so they react by gaining 1 electron in their outer shell to form a negative ion.
• As you go down group 7 the halogens become less reactive because there is less energy to gain an extra electron because the electrons are further away from the nucleus.
• As you go down group 7 the melting and boiling points increase.
• At room temperature...
• Chlorine is a fairly reactive, poisonous, dense green gas with a low boiling point.
• Bromine is a dense, poisonous orange liquid.
• Iodine is a dark grey crystalline solid with a high boiling point.
• Halogens want to gain an elctron to form an 1- ion with a stable electronic structure.
• Gain of electrons is called REDUCTION.
• Halogens react with alkali metals to form salts called metal halides e.g. Sodium Chloride.
• More reactive halogens will displace less reactive halogens.
• Chlorine + Potassium iodide ---> Iodine + Potassium chloride.
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## Metals

• All metals have the same basic properties. They're hard, dense, shiny and malleable.
• They're held together by metallic bonds that allow the outer electron of each atom to move freely. This creates a sea of delocalised electrons throughout metals which give them many of their properties.
• There's a strong attraction between the delocalised electrons and the positive ions.
• Metals have high melting and boiling points because of their strong metallic bonds.
• The strength of the metallic bond decreases as the the atomic radius increases.
• Metals are good conductors of heat and electricity because of the delocalised electrons. They are free to move so can carry the electrical current.
• Stainless steel is used to make saucepans because it's a good conductor of heat, doesn't rust easily and is cheap.
• Copper's used to make electrical wires because it's a good conductor of electricity and is bendy.
• Aluminium's used to make aeroplanes because it has a low density, is strong and doesn't corrode. Titanium is sometimes used but it is more expensive.
• Steel is used to make bridges because it's strong. Steel is made up of mainly iron but also some carbon which makes it less brittle.
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## Superconductors and Transition Metals

• All metals have some electrical resistance, this means whenever electricity flows through them they heat up.
• If a metal is cold enough the resistance disappears an the metal becomes a superconductor.
• This means no electrical energy is lost as heat, so if you started a current flowing through a superconductor circuit and took the battery out the current would carry on flowing.
• Using superconductor wires you can make power cables for loss-free power transmission, strong electromagnets and fast electrical circuits.
• Metals become superconductors at -265*C so the use of superconductors is limited.
• Some metal oxides can superconduct at -135*C which is cheaper.
• Metals in the middle of the periodic table such as zinc and copper are transistion metals.
• Transistion metals have typical metallic properties and some make good catalysts.
• Iron is used in the Haber process and nickel is used for the hydrogenation of alkenes.
• The compounds of transition elements are colourful because of the transiton metal ion they contain. Iron(II) compounds are green, iron(III) compounds are orange/brown and copper compounds are blue.
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## Thermal Decomposition and Precipitation

• Thermal Decomposition is when a substance breaks down when it's heated.
• Transition metal carbonates such as copper(II)carbonate break down into a metal oxide and carbon dioxide. This usually results in a colour change.
• Copper(II) Carbonate (green) ---> Copper Oxide (black) + Carbon Dioxide.
• A precipitation reaction is where two soulutions react to form an insouluble solid in solution. It has precipitated out.
• Some soluble transition metals react with sodium hydroxide to form an insoluble hydroxide.
• Copper(II) Sulfate + Sodium Hydroxide ---> Copper(II) Hydroxide + Sodium Sulfate.
• Some insoluble transistion metal hydroxides have distinctive colours, copper(II) hydroxide is blue, iron(II) hydroxide is grey/green and iron(III) hydroxide is orange/brown.
• You can use this to test which transistion metal a solution contains.
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## Water Purity

• In the UK we get our water from, surface water (e.g. lakes, rivers, reservoirs) and ground water (e.g. aquifers -rocks that trap water underground).
• All these resources are limited by annual rainfall.
• How much purification water needs depends on the source, ground water is usually cleaner than surface water. The process of water purification includes...
• Filteration - a wire mesh removes large twigs the gravel and sand beds filter out other solids.
• Sedimentation - iron sulfate or aluminium sulfate is added to the water. This makes small particles cluster together at the bottom.
• Chlorination - chlorine gas is bubbled through to kill harmful bacteria and other microbes.
• Tap water still contains impurities such as nitrate residues from fertilisers, theses are dangerous because they stop blood from carrying oxygen properly.
• Other impurities are lead and pesticide residues.
• In some dry countries sea water is distilled to produce drinking water.
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