OCR Gateway B Module C4: The Periodic Table

Atoms:

• A nucleus of an atom is made up of protons and neutrons
• Relative charges, neutron: 0, electron: -1, proton: +1. Therefore, the nucleus is positively charged and electrons are negatively charged, so the atom is neutral
• Relative masses, neutron: 1, electron: 0.0005 (generally considered as zero since it's so small), proton: 1. So the mass number of an atom is the number of protons and neutrons
• An atom is neutral because it has an equal number of electrons and protons. Positive charges cancel out the negative charges
• Radius of an atom is 1 x 10-10 m and a mass of 1 x 10-23 g
• The atomic number of an atom is the position in the periodic table (the bottom number) example: 199F the 9 is the atomic number.
• Mass number is the biggest number on the atom, example: 199F, the 19 is the mass number
• Number of protons = atomic number
• Number of electrons = atomic number
• Number of neutrons = mass number - atomic number, example:199F, 19 - 9 = 10, so number of neutrons = 10

Isotopes:

• An element with the same atomic number but different mass numbers are isotopes
• Carbon-12 is an example of an isotope (126C)

Electron arrangement:

• Elements are written in the periodic table with increasing atomic numbers
• Electrons occupy shells, with 2 being the maximum in the inner shell and 8 being the maximum in the next two. The innermost shells are filled first. So, for example, a lithium atom (3 electrons), has an electron structure of 2.1 (2 in the inner shell, 1 in the outer)
• The row that an element is in can be determined from the number of electron shells so sulphur 2.8.6 is in the third row as it has 3 shells
• The atomic number can be deduced from this by adding the numbers of electrons in each shell
• Elements in the same group are arranged vertically (have the same number of electrons in the outer shell)
• Elements in the same period are arranged horizontally (have the same number of shells)

Historical context of atoms:

• Early theory of atoms was developed by John Dalton, which was confirmed with more evidence later
• J.J Thompson discovered the electron
• Rutherford discovered the nucleus
• Bohr studied electron orbits
• Geiger and Marsden conducted an experiment which involved alpha particles being beamed at a gold leaf. Some particles were deflected by the positive nucleus but some were not, which showed that an atom contains a nucleus. They had unexpected results which made a huge contribution to Rutherford and Bohr's ideas

Ionic bonding:

• Atoms that have an outer shell of 8 electrons have a stable electronic structure as the shells are full
• Atoms that are not stable can be made stable by transferring electrons, this is called ionic bonding
• Positively charged atoms attract negatively charged atoms to create a strong bond. Ionic bonding is when a non metal atom forms a negative ion and a metal atom forms a positive ion, forming a strong attraction - an ionic bond
• Metal atoms lose electrons to get a stable electronic structure, when an atom loses electrons a positive ion is formed
• Non-metal atoms gain electrons to get a stable electronic structure, when an atom gains electrons a negative ion is formed
• Dot and cross diagrams are used to describe ionic bonding. You can either draw all of the shells or just the outer one.
• During the bonding of magnesium oxide, the magnesium atoms (originally 2 electrons in the outer shell) lose two electrons to form positive magnesium ions. Oxygen atoms (originally 6 electrons in the outer shell) gain two electrons to form negative oxide ions - if an atom loses electrons it becomes positive and if it gains electrons it becomes negatively charged

Covalent bonding:

• Non-metals share electron pairs between atoms, which is known as covalent bonding
• This can be shown as a dot and cross diagram which only show the outer shell arrangement

Sodium chloride and magnesium oxide - giant ionic lattices:

• Sodium chloride (NaCl) and magnesium oxide (MgO) structure is a giant ionic lattice where the positive ions have a strong electrostatic attraction to negative ions
• Therefore, they both have high melting points
• They cannot conduct electricity when solid as the ions can't move
• When they are in solution or molten they can conduct as they have free ions
• Magnesium oxide is made from bonded Mg2+ and O2- ions
• Sodium chloride is made from bonded Na+ and Cl- ions
• Magnesium oxide has a higher melting point than sodium chloride because Mg2+ and O2- ions have more charges than Na+ and Cl- so there are stronger electrostatic attractions between positive and negative electrons
• Each Mg atom donates two electrons to the oxygen atom, which makes a stronger bond when sodium atoms transfer one electron to chlorine atoms.
• Mg ions are very small in radius so the magnesium can get much closer to oxygen which makes the bond stronger, so more energy is required to separate them

Covalent bonding:

• Whereas ionic bonding is between a metal and non-metal, covalent bonds are between non-metals. A covalent bond is a shared pair of electrons

Chemical properties – water and carbon dioxide:

• Water (H2O), and carbon dioxide (CO2) cannot conduct electricity
• Forces between molecules are called intermolecular forces
• Water configuration: oxygen has 6 electrons in the outer shell so it combines with two hydrogen atoms which have 2 electrons in the outer shell, and complete the shell to have 8 electrons
• Carbon dioxide configuration: carbon has four electrons in its outer shell so it needs four more to complete it, oxygen has 6 electrons in the outer shell so they need two more to complete it, so each oxygen outer shell is shared with two of the electrons of the carbon outer shell
• The weak intermolecular forces give substances a simple molecular structure so they can easily be broken apart and have low melting points

Group number and elements in periods:

• The group number shows how many electrons are in the outer shell
• The number of shells that are full can be seen from the period the atom is in

Theories of the periodic table:

• 1817: Döbereiner put forward his law of triads in 1817. Each of Döbereiner's triads was a group of three elements, the appearance and reactions of the elements in a triad were similar to each other
• 1865: Newland put 56 elements into groups and saw that every 8th element behaved in a similar way, this was later accepted (50 years on) with more evidence
• 1869: Mendeleev arranged elements into order in a table, and predicted that there would be new elements found. He saw that there were gaps in the pattern where these elements would fit
• 1891: his idea of missing elements was confirmed

Group 1 elements (the alkali metals):

• Group 1 elements are very reactive and the most reactive is the element at the bottom of group 1 (Francium)
• They react vigorously with water, hydrogen gas is given off, and the metal reacts with the water to give an alkali - the hydroxide of the metal example: sodium + water → sodium hydroxide + hydrogen (`2Na + 2H2O → 2NaOH + H2`)
• They also react with air and must therefore be stored under oil
• As the reactivity increases down the group, the melting and boiling points decrease which helps us predict melting points for other metals in group 1
• Atoms of group 1 have similar properties because they have one electron in their outer shell
• Atoms of alkali metals react losing one electron and forming an outer shell which is full and creates a stable electronic structure, an ion which has more positive charge in its nucleus than negative electrons in the shells it becomes a positive ion
• This can be shown by half equations which are oxidation as they are losing electrons (`Na - e- → Na+`). Remember - oxidation is the loss of electrons and reduction is the gain of electrons in a reaction. This is often remembered as OILRIG - Oxidation is loss, reduction is gain

Flame tests:

• Flame tests can be used to determine if lithium, potassium or sodium are present in a compound
• To carry this out first a flame test wire is moistened in dilute hydrochloric acid, then the flame wire is dipped into the solid chemical, this is the put on a blue Bunsen burner flame and the colours are recorded
• If potassium is present the flame will be lilac, lithium presence is shown by a red flame, and a yellow flame indicates sodium

Group 7 elements - the halogens:

• At room temperature: chlorine is a green gas, bromine is an orange liquid and iodine is a grey solid
• Chlorine is used to sterilise water and to make pesticides and plastics, iodine is used to sterilise wounds
• Group 7 elements have similar properties as they have seven electrons in their outer shell
• It is possible to predict the properties of the halogens, as they all follow a trend
• Halogens have similar properties because when they react an atom gains one electron to form a negative ion with a stable electronic structure
• The nearer the nucleus to the outer shell, the easier it is for it to gain an electron, the more reactive the halogen

Halogen reactions:

• When a halogen reacts with an alkali metal, a metal halide is made
• Example: potassium with iodine forms potassium iodide, `potassium + iodine → potassium iodide`, `2K + I2 → 2KI`
• This is a reaction between group 1 and 7 elements which is very reactive
• They can be shown by half equations which are reduction as electrons are being gained `Br2 + 2e- → 2Br-`

Displacement reactions:

• When halogens are bubbled through solutions of metal halides there are two possible outcomes, either a displacement reaction or no reaction
• Chlorine displaces the bromide to form a bromine solution
• `Chlorine + potassium bromide → potassium chloride + bromine` this is an orange solution, `Cl2 + 2KBr → 2KCl + Br2`
• Bromine also displaces iodides from solutions, `Br2 + 2KI → 2KBr + I2`
• The reactivity of halogens increases further up the group allowing the reactions between halogens and metal halides to be predicted

Transition metals:

• The elements between group 2 and 3 are known as the transition elements
• Copper and iron are transition metals which are normally coloured. Copper: blue, iron(II): pale green, iron(III): orange/brown
• Transition metals and their compounds are often used as catalysts, iron for the Haber process - to make fertilisers, and nickel is used in the production of margarine to harden the oils
• Thermal decomposition is a reaction in which a substance is broken down into 2 or more different substances when heated
• When a transition metal carbonate is heated it undergoes thermal decomposition to form a metal oxide and carbon dioxide
• `FeCO3 → iron oxide + carbon dioxide`,
`CuCO3 → copper oxide + carbon dioxide`,
`MnCO3 → manganese oxide + carbon dioxide`,
`ZnCO3 → zinc oxide + carbon dioxide`
• `FeCO3 → FeO + CO2`,
`CuCO3 → CUO + CO2`,
`MnCO3 → MnO + CO2`,
`ZnCO3 → ZnO + CO2`
You may be asked to write the symbol or word equation for the thermal decomposition of a transition metal in an exam, but only for these four metals
• Sodium hydroxide solution reacts with compounds of each transition metal to make a solid of a particular colour, sodium hydroxide helps identify the presence of transition metal ions in the solution
• Cu2+ gives a blue solid, Fe2+ gives a grey/green solid, Fe3+ gives an orange/brown solid
• Precipitation reactions are reactions between solutions that make an insoluble solid, e.g.:
`Cu2+ + 2OH- → Cu(OH)2`,
`Fe2+ + 2OH- → Fe(OH)2`,
`Fe3+ + 3OH- → Fe(OH)3`

Metal structure and bonding:

• Metals have specific properties that make them good for different uses
• Physical properties: lustrous, hard and high density, high tensile strength, high melting and boiling points, good conductors of heat and electricity
• Aluminium has low density so is good in the use of cars and aircraft
• A chemical property of a metal is the resistance it has to oxygen and acids, which can be shown by gold. Copper is also resistant, hence why it is used in saucepans

Metallic bonding:

• Metals have high boiling points and melting points due to the strong metallic bonds, a lot of energy is required to break these bonds
• A metallic bond is a strong electrostatic force of attraction between close packed positive metal ions and a sea of delocalised electrons
• It can conduct electricity as the delocalised electrons within its structure can move easily
• Copper is used in wiring because it is a good conductor of electricity. Metal particles are held together by strong metallic bonds, which is why they have high melting and boiling points. The free electrons in metals can move through the metal, allowing metals to conduct electricity

Superconductors:

• Superconductors are metals that conduct electricity with little or no resistance
• The electrical resistance of mercury suddenly drops to zero at -268.8°c, this is called super conductivity
• When a substance goes to a superconducting state it no longer has magnetic fields inside it, when a small magnet is brought near the superconductor it is repelled and if a small permanent magnet is placed above the superconductor it levitates
• Benefits with superconductors: loss-free power transmission, super-fast electronic circuits, and powerful electromagnets
• However the difficulties with them is that they only work at low temperatures which limits their use and superconductors which function at 20°C need to be developed

Purifying and testing water:

• In the UK, water can be found at lakes, rivers, aquifers and reservoirs
• Water pollutants are nitrates from fertilisers, pesticides from crop spraying and lead from old water pipes which slowly dissolves into the water
• Before purification, the water usually contains dissolved salts and minerals, microbes, pollutants and insoluble materials
• Three main stages for water purification are:
- sedimentation: chemicals are added to make the solid particles and bacteria settle out
- filtration: water is passed through specially prepared layers of sand and gravel, as it passes through it removes fine particles also some sand filters kill the microbes
- chlorination: chlorine is added to kill microbes
• Sea water has so many dissolved substances, it requires huge amounts of energy and is very expensive and only used when there is no fresh water

Water tests:

• Water can be tested with precipitation reactions using aqueous silver nitrate and barium chloride solution to test for the presence of certain ions
• The equations:
`barium chloride + magnesium sulfate → barium sulfate (white precipitate) + magnesium chloride`
`silver nitrate + sodium bromide → silver bromide (cream precipitate) + sodium nitrate`
`silver nitrate + sodium iodide → silver iodide (yellow precipitate) + sodium nitrate`
`silver nitrate + sodium chloride → silver chloride (white precipitate) + sodium nitrate`
So for the first equation here, a white precipitate is formed if barium chloride is added to a solution which contains magnesium sulfate
• Balanced symbol equations:
`BaCl2 + MgSO4 → BaSO4 + MgCl2`
`AgNo3 + NaBr → AgBr + NaNO3`
`AgNo3 + Nal → Agl + NaNO3`
`AgNO3 + NaCl → AgCl + NaNO3`