``` Shiny Hard Malleable Ductile Sonorous High melting&boiling points Good conductors of heat & electricity ```

Metals Bonding And Reactivity Flashcards by Anais Tokley

Properties of metals

Shiny
Hard
Malleable
Ductile
Sonorous
High melting&amp;boiling points
Good conductors of heat &amp; electricity

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Metallic bonding

A LATTICE of positive ions held by a SEA OF DELOCALISED ELECTRONS. The forces between the ions and electrons is electrostatic force.

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Delocalised

Electrons are free to move because they are not attached to a particular ion.

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Describe electrical conductivity of metals

Metals conduct electricity because they have DELOCALISED eletrons (free) who can carry the charge (electricity) around the lattice.

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Describe malleability of metals

Metals are malleable because they have delocalised electrons that are free to move and so enable the metal ions to ROLL PAST each other when a stress is applied to the metal.

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Ionic bonding

Metal + non - metal

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What proves that both ionic + metallic bonding must consist of particles held by strong forces in a neat regular lattice?

Both
Are crystalline
Conduct electricity when molten
Have high melting points

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Differences between ionic and metallic bonding

Malleability (metallic)
Ductility (metallic)
Conductivity when solid(metallic)

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Why do ionic substances only conduct when molten?

Because although they too are made of charged particles, these are not mobile

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Definition of an alloy

Mixtures of metals with other elements, most commonly other metals or carbon e.g. Brass

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Why are alloys used?

Give particular properties
E.g. Flexibility
Strength

For example adding steel to iron makes it harder and resistant to corrosion

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Why are alloys stronger?

They contain ions of different sizes which distort the regular arrangements of ions in the metal structure. This makes it more difficult for the layers to slide past each other so alloys are harder than pure metals

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Chemical properties of metals

React with dilute acids
'' oxygen
Form BASIC oxides
Have good reducing agents
Lose their valence electrons easily
Have 1-3 outer electrons
React with water

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Indication of speed of reaction

Rate at which bubbles are given off

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Metal + acid

—> salt + hydrogen

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Metal + oxygen

—-> metal oxide

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Metal + water

—-> metal hydroxide + hydrogen

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Describe reaction of Magnesium with HCl acid

Immediate response :
Fizzing
Hydrogen produced (+ magnesium chloride)
Mg (s) + 2HCl (aq) –> MgCl2 (aq) + H2 (g)

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Describe reaction of oxygen with sodium

Melts in a small ball
While burning, starts to glow
Turns orange
White smoke produced
ALKALINE solution

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Desribe reaction of water and calcium

Hydrogen produced
White precipitate formed
ALKALINE solution

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Reactivity serie

Potassium
Sodium
Calcium
Magnesium
Zinc
Iron
Hydrogen
Copper

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Why is Potassium very reactive?

It has only 1 valence electron so it is very easy to lose it forming a positive ion.

Describe reactivity series

Going from bottom of reactivity series, they lose electrons more readily to form positive ions (easier for first one to lose one)

The more reactive a metal, the more ……. To split up its oxides

Difficult

Explain reduction by carbon

Metals below carbon in the reactivity series can be reduced by carbon. Ex : Copper oxide + carbon --> copper + carbon dioxide

Definition of displacement

Atoms of the most reactive metals will push atoms of less reactive metals out of the compound

Explain the competition of metals for oxygen

When a metal is heated with the oxide of a less reactive metal, it will remove the oxygen from it Ex : iron + copper(2) oxide ---> iron (2) oxide + copper

Practical example of redox reaction (= displacement reaction = reduction/ oxidation reaction)

Thermit reaction (to make liquid iron for welding railway tracks together)

Experiment to see carbon is in the reactivity series

Heat carbon with metal oxides Carbon reacts with copper and iron but not with magnesium so in between 2

Why is alumium apparently unreactive?

It is actually reactive but has and OXIDE layer that protects the metal. It therefore seems to have no reaction with water, steam and dilute acids. However the coating of aluminum oxide is invisible to naked eye and protcts it from contact with reagents.

Definition of thermal decomposition

The breaking down of a substance using heat

Metal hydroxide + heat

----> metal oxide + water

Do group 1 metals decompose with heat?

Metal nitrate + heat

------> 1) alkali metals (group 1) Nitrite + oxygen Except lithium -------> 2) nitrates of other metals Oxide + nitrogen dioxide + oxygen

Metal carbonate + heat

-----> metal oxide + carbon dioxide NOT GROUP 1

Quicklime

Calcium oxide

Limestone

Calcium carbonate

Calcium carbonate + heat

------> calcium oxide + carbon dioxide

Uses of quicklime

Neutralise acid Make slaked lime Neutralise industrial waste products

Quicklime + a little bit of water

--------> slaked lime (calcium hydroxide)

Slaked lime + a lot of water

------> lime water (calcium hydroxide solution)

Limewater + bubble in CO2

--------> calcium carbonate (limestone)

Uses of slake lime (calcium hydroxide)

Neutralise acid in soil | '' industrial waste

Uses of limestone (calcium carbonate

Manufacture iron and cement Making glass Neutralise acid in soil and lakes Neutralise industrial waste products

Thermal stability

How resistant a metal is to thermal decomposition

Definition of oxidation

Addition of oxygen to a substanec

Definition of reduction

Removal of oxygen from a subtsance

Definition of redox reaction

During redox reactions, electrons are transferred from one substance to another. Oxidation gains electrons, reduction loses electrons

Definition of oxidation agent

A substance whihc oxidises antoher subatcne during a redox reaction

Definition of reducing agent

A substance which reduces another during a redox reaction

Metal atom + oxidation ->

Positive ions

Metal ions in compounds + reduction ->

Metal atoms

Uses of aluminium

``` Manufacture of aircraft (because of strenght and density) Food containers (resistance to corrosion) ```