Section 2- Metals

Extracting Metals:

• are linked to the order of reactivity 

1. Only metals less reactive than carbon can be extracted by a reduction reaction with carbon ----------> heating the ore with carbon monoxide 
• more reactive elements form compounds more readily 
• reduction= loss of oxygen 
2. Hydrogen: metals above H react with acids to release H2 

Extracting Iron:

(Blast Furnace) 

Raw materials:

1. Iron ore -----> iron
2. Coke -----> almost pure carbon; reducing iron oxide to iron metal
3. Limestone -----> take away impurities from slag 
4. Air ------> allows the coke to burn 

Reducing Iron ore to Iron:

1. Hot air blasted into the furnace -----> coke burns faster than normal. Raises to 1500C.  
2. Coke burns and produces carbon dioxide:
• Carbon and Oxygen ------> Carbon Dioxide 
• C            +        O2    ------> CO2 
3. Carbon monoxide 
• Carbon dioxide and Carbon -----> Carbon monoxide 
•          CO2          +         C     ----->          2CO 
4. Carbon monoxide then reduces the iron ore to iron
• Carbon monoxide + Iron(II) oxide ------> Carbon dioxide + Iron
•          3CO             +       Fe2O3       ------>        3CO2        +  2Fe 
5. Iron is molten at temperature also dense so goes to the bottom where it's tapped off 

Removing impurities: 

Slag forms as CaO from limestone reacts with rocks (SiO2) 

CaCO3 -----> CaO + CO2 

CaO + SiO2 ------> CaSiO3 (molten slag) 

• Coolen to solid and used for:
• road building 
• fertiliser 

Uses of Iron:

• steel 
• buildings
• car radio 
• railings

REDOX -------> reduction and oxidation occur together 

Oxidising Agent --------> produces Oxygen (O2) in the process. reduced itself

Reducing Agent --------> gains Oxygen (O2) in the process. oxidising itself

Rusting of iron:

• needs water and oxygen (from air) 
• doesn't rust in boiled water or dried air 

How to prevent rusting?

Barrier methods: 

• paint 
• plastic coat 
• tin plate 
• grease 
• coat in zinc 

Sacrificial methods: 

• Attach above (more reactive) metal to the iron (e.g. zinc) 

2f) Reactivity Series

 

Potassium	K
Sodium	Na
Calcium	Ca
Magnesium	Mg
Aluminium	Al
Carbon	Ca
Zinc	Zn
Iron	Fe
Tin	Sn
Lead	Pb
Hydrogen	H
Copper	Cu
Silver	Ag
Gold	Au

2e) Hydrogen and water

Prepared by:
Zn + H2SO4 -----> ZnSO4 + H2                        displaces 

Burns in air:
2H2 + O2 -----> 2H2O 

Physical test for water boils at 100C
Chemical test turns anhydrous copper sulfate -----> blue 

2d) Oxygen and Oxides

How to make O2:
2H2O2 (aq) -----> 2H2O (l) + O2 (g) 

metal oxides- bases
non metal oxides- acidic
water- neutral

Carbon dioxide prepared by
CaCO3 + HCl ------> CaCl + CO2 + H2O 

Delivery of CO2
Either upwards so that it's lighter than air
Or       downwards so that it is heavier than air

• Under pressure- disowns 
• used in fire extinguishers

Thermal disposition of copper(II) carbonate:

CuCO3(s) -----> CuO(s) + CO2(g) 

2c) Group 7 elements

Chlorine, bromine and iodine
called Halogens

Properties:

Group VII elements	Atomic Number	Colour	At room temperature	Boiling Point
Chlorine	17	green	gas	-34C
Bromine	35	red-brown	liquid	59C
Iodine	53	dark grey	solid	185C

Bonding in molecules:

Displacement reactions:

• more reactive halogens displace less reactive ones 
• a solution of its salt 
• Cl2 + 2KBr -----> 2KCl + Br2 (orange-brown)
• Cl2 + 2KCl -----> 2KCl + I2 (dark brown)
• Br2 + 2KBr -----> 2KBr + I2 (dark brown) 
• REDOX Reactions 

Summary: Chlorine is smaller than bromine has less outer shells of electrons sheilding the nucleus 

attracts electrons more readily 

More reactive as the atomic number decreases 

HCl gas vs HCl (aq) 

• gas: covalently bonded 
• no ions so can't conduct electricity 
• not acidic 

HCl dissolved in water:

• forms H+ (aq) and Cl- (aq) 
• ions conduct electricity 
• solution acidic 

2b) Group 1 elements

The group 1 elements are lithium, sodium and potassium

• They all react a similar way in water:

metal and water -------> metal hydroxide and hydrogen
2M(s)       2H2O(l)            2MOH(aq)                H2(g)

• alkali metals 
• reactive metals 
• shiny silver when freshly cut
• low density 
• low melting point 

Metal	Observations
Lithium (Li)	Fizzes, floats and moves on the water surface
Sodium (Na)	Fizzes, floats, forms a ball and moves around surface of water (may ignite)
Potassium (K)	Fizzes, reacts vigorously and burns with a lilac flame (sometimes explodes)

Summary: Group 1 elements get more reactive as the atomic number increases 

1. Atoms increase in size 

2. Outer electrons further from nucleus 

3. Therefore they are easily lost 

2a) The Periodic Table

Origins of the periodic table

1. Newland's Octaves:
• elements in groups have similar properties 
• didn't leave gaps for undiscovered elements 
• metals and non metals are mixed up
2. Mendeleev 
• arranged by mass
• left gaps 
• some elements in wrong order 
• now arranged in order of atomic number 

LHS metals                                                                                                                  RHS Non-metals 

Groups- elements that have the same number of outer shells are part of the same group

              They have similar chemical properties and physical appearance is similar 

Periods- number of electrons or protons increases across period recurring patterns 

Transition Elements

• metals 
• hard/strong
• high melting points 
• unreactive
• act as a catalyst 
• form a coloured compound 
• 3rd shelll being filled across period 

Section 2

Chemistry of the elements

1f) Ionic compounds + 1g) covalent compounds

Ionic bonding:

• strong electrostatic attraction 
• opposite charged ions 
• throughout lattice 

Properties of ionic compounds:

• high melting point 
• strong electrostatic attraction 
• giant ionic lattice 
• soluble in water 
• Ions attracted to polar water molecules 
• conduct electricity 
• ions can move and carry charge

A Lattice is a 3-D network of ions in fixed positions bonded together 

OIL RIG 

Oxidation- loses electrons                      Reduction- gains electrons 

Group	1	2	3	4	5	6	7	0/8
Charge	1+	2+	3+	X	3-	2-	1-	X

Covalent bonding:

• non metal with non metal 
• outer shells overlap 
• electron shared in pairs 
• atoms achieve outer shells (full)
• attracts both nuclei 

Properties of simple covalent molecules:

• very strong bonds
• low melting and boiling points 
• insoluble in water 
• no free electrons 
• Don't conduct electricity 
• soluble in organ solvents 

Intermolecular forces 

• weak bonds
• easily broken 

Giant covalent bonds 

• strong bonds 
• no free electrons 
• rigid 
• all atoms joined together 
• covalent unreactive 

Diamonds 

• rigid 
• hard 
• sparkly
• four covalent bonds  

Graphite

• layers three covalent bonds 
• slide over each other 
• free electrons 
• good conductor of electricity 
• lubricant 

1c) Atomic structure

Nucleus- contains neutrons and protons
Protons- has a positive charge due to neutrons
Electrons- move around the nucleus in shells; negatively charged
2,8,8,2: Elements in the same group have the same number of outer shell electrons

Atoms consist of a nucleus in the centre, which has neutrons and protons. Protons have a positive charge due to neutrons. There are also electrons which move around the nucleus in shells and are negatively charged. Elements in the same group have the same amount of outer shell electrons.

Particle	Relative Mass	Relative Charge
Proton	1	+1
Neutron	1	0
Electron	1/2000	-1

• Atomic Number tells you how many protons there are
• Mass number is the total number of protons and neutrons in an atom
• Isotopes have the same chemical reactivity but different number of neutrons they also have the same molecular formula but different structural formula.
• The Relative atomic mass (Ar) is the average mass of all the isotopes of an element. It allows for the relative mass of each isotope and its relative abundance.
• Relative Abundance is how much there is of each isotope compared to the total amount of the element in the world. This can be represented as a ratio, fraction or percentage.

The Periodic Table

Is an arrangement of elements in order of their atomic number

Electronic configurations of the 1^st twenty elements

Element	Electronic Configuration
Hydrogen	1
Helium	2
Lithium	2,1
Berylium	2,2
Boron	2,3
Carbon	2,4
Nitrogen	2,5
Oxygen	2,6
Fluorine	2,7
Neon	2,8
Sodium	2,8,1
Magnesium	2,8,2
Aluminium	2,8,3
Silicon	2,8,4
Phosphorus	2,8,5
Sulphur	2,8,6
Chlorine	2,8,7
Argon	2,8,8
Potassium	2,8,8,1
Calcium	2,8,8,2