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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  • 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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  • 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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