Metals:
Giant metallic structures with metal ions that
a. are closely packed together in a regular 3-D pattern/lattice, with valence electrons from each atom free to move at random in space between ions, forming "sea" of electrons
b. have as metallic bonds (see Bonding and structure) binding atoms together
c. lose electrons to form positive ions

Physical Properties (influenced by structure and thus bonding between atoms)

Chemical properties (due to position in periodic table/ metallic character)

Alloys: Pure metals are soft and may react with air and water, thus corroding easily. Alloys are thus used in place of pure metals.


Is a mixture of a metal with one or more other elements

Non-metals

Metals

Appearance: Usually dull and cannot be polished
M.p and b.p: Mostly liquids or gases with low m.p and b.p (except for diamond, silicon, graphite that have high m.p due to giant molecular structure )
Heat and electrical conductivity: Poor conductors
Density: Relatively low densities
Malleability: Brittle, breaks without bending and strectching
Ductility: Not ductile, cannot be stretched

Appearance: Usually lustrous and can be polished
M.p and b.p: Solids with high m.p and b.p (except for mercury which is a liquid at room temp )
Heat and electrical conductivity: Good conductors
Density: High density
Malleability: Can be bent and beaten into different shapes
Ductility: Can be stretched and pulled into wires

High m.p and b.p
Atoms in a metal are packed tightly in layers and held together by strong metallic bonds. Large amount of energy is needed to overcome the strong electrostatic forces of attraction between positive metal ions and "sea" of delocalised electrons.

Ductile and malleable
Due to the giant metallic structure, pure metals' ions are arranged regularly and orderly in layers. The atoms are of the same size. When a force is applied, layers of metal ions of the same size can easily slide over one another.

Heat and electrical conductivity
In the giant metallic structure, the outermost electrons of the atoms can break away easily from the valence shell. Positive metal ions are thus surrounded by a "sea" of mobile electrons that are free to move throughout the metal to conduct electricity.


Movement of mobile electrons within metal lattice allows heat to be transferred easily.




High densities
Metals have giant metallic structures, metal atoms are arranged in a regular and closely packed manner.

Metals

Non-metals

Formation of ions: Form positive ions by losing electrons → are good REDUCING AGENTS
Formation of oxides: Form basic and amphoteric oxides by burning/ reacting with oxygen


some basic oxides dissolve in water to form alkalis

Formation of ions: Form negative ions by gaining electrons → are OXIDISING AGENTS
Formation of oxides: Form acidic and neutral oxides by burning/ reacting with oxygen


some basic oxides dissolve in water to form acids

Physical properties

a. are harder and stronger than pure metals
b. improves the appearance of metals
c. lower the melting point of metals (for industrial purposes)
d. make metals more resistant to corrosion

Corrosion
Metals in pure form are chemically reactive and can be easily corroded by the surrounding atmospheric gases and moisture. Alloying a metal increases the inertness of the metal, which, in turn, increases corrosion resistance.

Stronger and harder
In alloys, atoms of added element have a different size from atoms in the pure metal. This disrupts the orderly arrangement of atoms in pure metals. Hence it is much harder for atoms to slide over each other when a force is applied.

Lower m.p
Impurities disrupt the regular arrangement of atoms and hence the bonding in between atoms. Less energy is needed to break the weakened bonds and melting point decreases

.
(Higher b.p?)
When boiling, particles are in liquid state and can rearrange themselves to bond with ions that allows them to achieve maximum attraction and hence stability. More energy is needed to break the strengthened bonds and b.p increases.

Reaction (depends on position in reactivity series)



Lithium Li
Potassium K
Strontium Sr
Calcium Ca
Sodium Na
Magnesium Mg
Aluminum Al
Zinc Zn
Chromium Cr
Iron Fe
Cadmium Cd
Cobalt Co
Nickel Ni
Tin Sn
Lead Pb
Hydrogen gas H2
Antimony Sb
Arsenic Ar
Bismuth Bi
Copper Cu
Mercury Hg
Silver Ag

Extraction

Low down:
opposite of "high up"

High up: Metals have a greater tendency to lose electrons to form positive ions then a metal less reactive


  • reacts vigorously and quickly with chemicals
  • corrodes easily
  • readily gives up electrons in reactions to form positive ions

.

.

With Dilute ACIDS

Displacement

With COLD WATER/HOT STEAM

Reduction metal oxides with non-metals

Thermal decomposition (of carbonates)
The more reactive the metal, the more thermally stable is the carbonate


Ag2CO3 decomposes to give silver, O2 and CO2
No such thing as Al2(CO3)3 and Fe3(CO3)2

non-metal 1: carbon (up till zinc)
the more reactive, the harder to take

non-metal 2: hydrogen (up till iron)
the more reactive, the harder to take

of metals from solutions:

  • Reactive metal, greater tendency to lose electrons to form +ve ions (electropositive)
  • Ions of less reactive metals can more readily accept electrons from more reactive metal

  • Further apart, easier/faster displacement reaction


  • Iron, being more reactive than copper, loses electrons more readily than copper

  • Iron displaces Cu2+ from blue CuSO42- to form pale green FeSO4 and reddish brown Cu

of metal oxides with more reactive metal:
Metal higher in reactivity series can take O2 away from oxide of a metal lower in reactivity series.


e.g Aluminium
Reaction VERY EXOTHERMIC with a lot of heat and light energy given out (thermit)


Aluminium being more reactive than iron in iron(III) oxide will hence take O2 away from iron(III) oxide to form aluminium oxide


USED TO REPAIR BROKEN PARTS OF RAIL, MOLTEN IRON (reduced and forms iron) FLOWS INTO BREAK TO WELD BROKEN PARTS

Step 1: Getting rid of impurities
Step 2: Obtaining pure metal

Methods depends on reactivity

Case study: Extraction of Iron

What is iron needed for?

  • to be converted into steel, an alloy

Reaction (corrosion/ oxidation)

Rusting

Rust prevention:

  1. Use protective layer by placing barrier around metal to prevent oxygen in air and water from reaching surface of iron
  • Painting, Oiling and greasing, Plastic coating, Tin plating (if scratched, iron rusts quickly when O2 and water reach iron as iron being more reactive than tin, loses electrons more readily and in preference to tin, tin not sacrificial) , Chrome plating (chromium attractive appearance, once exposed to air and moisture, hard non-porous coating of chromium(III) oxide formed on surface of iron
  1. Sacrificial protection:
  • Zinc plating (galvanising)
    Zinc is cheap, has fairly low m.p and can thus be easily applied by dipping iron/steel in molten zinc)

Zinc being more reactive than iron loses electrons more readily than iron and corrodes in place of iron.


  • Attaching metal blocks such as Zn/ Mg
  1. Using alloys (rust resistant)

-Stainless steel, alloy of iron, carbon and large amnt of chromium and nickel (once exposed to air and moisture, hard non-porous coating of chromium(III) oxide formed on surface of iron and prevents rusting)

← For

Recycling because continuous extraction is not possible as metals are finite resources

Advantages

Issues