Group 1 elements
Group 1- The Alkali Metals Li, Na, K, Rb, Cs, Fr
As you go DOWN the group the alkali metals become:
- Bigger atoms-one extra full shell of electrons for each row down
- More reactive-the outer electron is easily lost (further from nucleus)
- Higher density- atoms have more mass.
- Lower melting point
- Lower boiling point
Properties of Alkali metals:
- Very reactive- stored in oil and handled with forceps
- All have one outer electron
- All form 1+ ions- lose their one outer electron
- Always form ionic compunds
- Reaction with water produces hydrogen gas
Group 7- The Halogens
Group 7- The Halogens F, Cl, Br, I, At
As you go DOWN the group the halogens become:
- Less reactive
- Higher melting point
- Higher boling point
Properties of Halogens:
- All non-metals with coloured vapours
- All form molecules which are pairs of atoms
- React with metals to form salts (or metal halides)
- More reactive halogens will displace less reactive ones
Chlorine can displace bromine and iodine from a solution of bromide or iodide.
Alkali Metals form a hydroxide in solution: aqueou
sodium + water sodium hydroxide + hydrogen
2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)
The metals react vigorously when placed in water and fizz furiously, producing hydrogen.
To test for hydrogen, do the squeaky pop test. A lighted splint will indicate hydrogen by producing this squeaky pop.
Coloured vapours of halogens
Fluorine is a very reactive, poisonous yellow gas.
Chlorine is a fairly reactive, poisonous dense green gas.
Bromine is a dense, poisonous, red-brown volatile liquid.
Iodine is a dark grey crystour.alline solid or purple vapour.
Transition Metals- In the middle of Group 2 and Group 3
Transition metals are 'typical' metals, with properties you would expect of a metal:
- Good conductors of heat and electricity
- Dense, strong and shiny
- Much less reactive than Group 1 metals- do not really react with water or oxygen
- Denser, stronger and harder than Group 1 metals, and have a much higher melting point (excluding mercury which is liquid at room temperature
- They often have more than one ion e.g. Fe2+ and Fe3+. Different ions usually form different coloured compounds: Fe2+ ions give green compounds, Fe3+ ions form red/brown compounds.
- The compounds are very colourful (see back of card)
- Transition metals and their compounds make good catalysts- Iron is the catalyst used in the Haber process, Nickel is used to turn oils into fats.
- Their properties are due to the way their electron shells fill (see next card).
The colours of compounds
The compounds are colourful due to the transition metal ion they contain e.g.
- Potassium chromate (VI) is yellow
- Potassium manganate (VII) is purple
- Copper sulphate (II) is blue
The colours of hair, gemstones and pottery glazes are all due to transition metals.
Transition metal electron shells
The further you get from the nucleus, the closer the energy levels get until they start to overlap. This first happens between energy levels 3 and 4. It affects the way the shells fill. Potassium has 19 electrons, but the 19th electron goes into the 4th shell, so the electron arrangement is 2,8,8,1.
However, with the transition metals, the electrons are put into the overlapping 3rd energy level until it's full.
- Sc 2,8,9,2
- Ti 2,8,10,2
- V 2,8,11,2
- Cr 2,8,13,1 (Cr and Cu fill up a bit differently. Just know the numbers, not why.)
- Mn 2,8,13,2
- Fe 2,8,14,2
- Co 2,8,15,2
- Ni 2,8,16,2
- Cu 2,8,18,1
- Zn 2,8,18,2
Acids and Alkalis
Lowry and Bronsted said acids are proton donors
Acids release H+ ions- i.e. they're proton donors
Bases accept H+ ions- i.e. they're proton acceptors
Protons are hydrated in water.
In acidic solutions, the acid molecules dissociate, releasing lots of H+ ions. These become hydrated and are called hydrated protons, represented by H+ (aq). These make acids acidic.
In basic solutions, water molecules can dissociate into H+ and OH- ions, although not really in pure water. Some base molecules e.g. amonia can take hydrogen ions from water, causing more molecules to dissociate and leaving an excess of OH- ions behind. Others release hydroxide ions straight into the solution.