metal reactivity

Structure of metals

All metals have a giant metallic lattice structure with strong electrostatic forces of attraction between the metal ions (positive) with the sea of delocalised and mobile electrons (negative)

Physical properties of metals

1. Good conductors of heat and electricity

2. Usually have high melting and boiling points (Exocet Grp 1 and mercury)

3. Have high densities

4. Malleable (can be hammered into different shapes)

5. Ductile (can be drawn into wires without breaking)

Reasons for good conductors of electricity and heat

The sea of delocalised electrons act as mobile charge carriers and conducts heat

Have high melting and boiling points (except group 1 and mercury )

They have giant metallic lattice structures with strong electrostatic forces of attraction between metal ions and sea of delocalised and mobile electrons which require large amounts of energy to overcome

High densities

Metal ions are closely packed together in the solid state

Malleable And ductile

Atoms are orderly arranged in layers and these layers of atoms can slide over each other easily when a force is applied.

Alloy

A mixture of a metal with other elements (non metal and metal). They are stronger and more corrosion resistant

Examples of alloy

1. Mild steel; —> mixture of iron and carbon (used to make car bodies and machinery)

2. Stainless steel—> mixture of iron chromium nickel and copper (used to make cutlery , surgical instruments)

Why are alloys stronger and harder than pure metals

Alloys are made up of atoms of different sizes which disrupts the orderly arrangement of atoms making it hard for the layers of atoms to slide over one another when a force is applied. On the other hand, atoms are orderly arranged in layers . The layers of the atoms can slide easily over on another when a force is applied .

Reactivity Series table

Potassium

Sodium

Calcium

Magnesium

Aluminium

Carbon

Zinc

Iron

Tin

Lead

Hydrogen

Copper

Silver

Gold

The reactivity series can be used to predict

1. Chemical behaviour of a metal

2. Position of an unfamiliar metal in the reactivity series

3. The more reactive the metal is the more electrons they lose easily to form cations

Reactivity of metals can be compared by these reactions

1. Metal +cold water —> metal hydroxide and hydrogen

2. Metal + steam —> metal oxide and hydrogen

3. Dilute hydrochloric acid + metal —> meta chloride + hydrogen

Reaction with cold water (potassium)

Reacts very violently . Potassium darts around and sizzles on the surface of the water . Effervescence of colourless , odour less gas that extinguishes lighted splint with a ‘pop’ sound.

Reaction with cold water (sodium)

Reacts violently . Sodium darts around and sizzles on the surface of water .Effervescence of colourless , odour less gas that extinguishes lighted splint with a ‘pop’ sound

Reaction with cold water (calcium)

Reacts readily .Effervescence of colourless , odour less gas that extinguishes lighted splint with a ‘pop’ sound

Reaction with cold water (magnesium )

Reacts very slowly. Effervescence of colourless , odour less gas that extinguishes lighted splint with a ‘pop’ sound

Aluminium when reacting with water , stem or hcl

No visible reaction because aluminium reacts readily with oxygen in the air to form an insoluble protective layer of aluminium oxide which prevents it from further reaction

Reaction with water from carbon onwards

No reaction

Reaction with steam (potassium-calcium)

Reacts explosively

Reaction with steam (magnesium)

Reacts violently. Bright white glow produced, grey solid turns white

Reaction with steam (zinc)

Reacts readily. Hot grey zinc forms a yellow solid that turns white when cooled.

Reaction with steam (iron)

Grey iron turns red before forming a black solid

Reaction with steam (lead-gold)

No reaction

Reaction with dilute hydrochloric acid (Potassium , sodium , calcium, magnesium )

Potassium +sodium = reacts explosively

Calcium = violently

Magnesium = rapidly

All produce Effervescence of colourless , odour less gas that extinguishes lighted splint with a ‘pop’ sound

Reaction with hcl (zinc, iron)

Zinc= moderately fast

Iron = slowly

All produce Effervescence of colourless , odour less gas that extinguishes lighted splint with a ‘pop’ sound

Reaction with dilute hcl (lead-gold)

No reaction

Displacement reactions Definition

A more reactive metal can displace a less reactive metal from its salt solution

Eg . Iron + copper sulfate what is the displacement reaction

Iron sulfate + copper. Because iron is more reactive and loses electrons more readily. Hence, iron will displace copper from copper sulfate solution to form iron sulfate and copper

Colour of group 1 and 2 metals

Colourless

Colour of salts of copper and iron

Copper —> blue (green)

Iron —> Fe(ii)) green , Fe(iii) yellow/ brown

Action of heat on carbonates

The more reactive the metal, the more thermally stable is it carbonate . Hence, more difficult to decompose the carbonate by heat

When potassium and sodium carbonate are heated what happened

They do not decompose because they are very reactive metals and hence their carbonates are very thermally stable

Calcium , magnesium Zin , iron , lead , copper carbonate upon heating what happens

Decomposes into metal oxide and carbon dioxide on heating

Zin carbonate —> zinc oxide colors

Carbonate → white solid

Zinc oxide __> yellow when hot , white when cold

Copper carbonate —> copper oxide colour

Carbonate —> green

Oxide—> black

Silver carbonate Upon heat

Decomposes into silver and carbon dioxide on heating

Reaction between a metal and the oxide of another metal

A more reactive metal can reduce a less reactive metal from its oxide

Magnesium ribbon function in reduction of metal oxides

Acts as a fuse to start the reaction as hen magnesium burns it produces a lot of heat which will provide the necessary activation energy to start the reaction

Reduction of metal oxides with carbon and hydrogen (reducing agents)

The more reactive a metal is , the more difficult it is to reduce the metal from its oxide by carbon and hydrogen

Reduction with carbon (Potassium to aluminium oxide)

Cannot be reduced

Reduction with carbon (zinc to copper oxide)

Reduction by carbon to form carbon dioxide and metal

Reduction with carbon and hydrogen on silver oxide

It can be decomposed by heating alone without any reducing agent

Reduction with hydrogen (Which cannot be reduced)

Potassium , sodium , calcium , magnesium , aluminium, zinc

Reduction with hydrogen (which can be reduced)

Reduced by hydrogen to form metal + steam