Windsor University
School of Medicine
CHARACTERISTICS OF
METALS & NON-METALS
ACCEPT THE CHALLENGE, SO YOU MAY FEEL THE EXHILARATION OF VICTORY.
GEORGE S. PATTON
Ch 17. J.C. Rowe
Periodic Table/ Review



The periodic table contains chemical elements arranged by
order of atomic number allowing the periodic properties
(chemical periodicity) of the elements to be made clear.
The table includes Periods (horizontal sections) & Groups
(usually vertical). Elements in groups have similar
properties to each other.
The table is a masterpiece of organized chemical
information. The evolution of chemistry's periodic table
into the current form is an astonishing achievement with
major contributions from many famous chemists & other
eminent scientists
Metals & non-metals in the
Periodic Table.
Stable electronic configurations
Metals / cations



Metals tend to lose their
outer-shell electrons in
order to be stable.
Stable positively charged
ions (cations)
The ability to lose
electrons makes them
good reducing agents.
Non- metals/ anions



Non-metals tend to easily
gain electrons in order to
be stable.
Stable negatively charged
ions (anions)
The ability to gain
electrons makes them
good oxidising agents.
Valence electrons
Periodic Table
Chemical properties of selected
metals
Sodium (Na)
Calcium (Ca)
Aluminium (Al)
Iron (Fe)
Lead (Pb)

Potassium (K)
Magnesium (Mg)
Zinc (Zn)
Copper (Cu)
Chemical Properties of Sodium


Chemical properties are all those properties
that are visible only when any reaction is
taking place between sodium and any other
chemical substance.
As per the periodic table, sodium is more
reactive as compared to lithium and has less
reactive properties than potassium.
Reaction with water


Reaction of sodium with water results in the
formation of sodium hydroxide and hydrogen
gas. As heat is produced during this reaction,
it is called exothermic reaction. This released
heat often ignites the hydrogen gas and as a
result fire may break out. If large pieces of
sodium is put into water it can lead to loud
explosions.
2Na(s) + 2H2O(l)
2NaOH(aq) + H2(g)
Reaction with Oxygen



Sodium readily reacts with oxygen to form
sodium oxide.
When a fresh piece of sodium comes in contact
with air, it forms sodium oxide (Na2 O)
instantly and this oxide forms a white coating
and protects the underlying metal from any
further reaction.
4Na (s)+ O2 (g)
2Na2O (s)
Reaction with Oxygen cont’d.


When sodium is burned in air, it reacts with
atmospheric oxygen to form sodium peroxide
(Na2O2).
2Na (s)+ excess O2 (g)
Na2O2 (s)
Reaction with acids
Sodium reacts with acid to produce the
corresponding salt & hydrogen :

2Na (s)+ 2HCl(aq)

H2SO4 (aq)+ 2 Na (s)
2NaCl (aq)+ H2 (g)
Na2SO4 (aq)+ H2 (g)
Reaction with halogens



The reaction between sodium and chlorine is
very vigorous and the product is sodium
chloride. The following equation illustrates
this reaction.
sodium + chlorine = sodium chloride
2 Na(s) + Cl2(g)
2Na + Cl-(s)
Sodium + bromine = sodium bromide
2Na(s)+ Br2(g)
2NaBr(s)

Chemical Properties of Potassium

The reactions of potassium & sodium are
similar; however the reactions of potassium
are more violent.
Chemical properties of Calcium



Calcium is a moderately active element. It reacts readily
with oxygen to form calcium oxide (CaO):
Calcium reacts with the halogens— fluorine, chlorine,
bromine, iodine.
Calcium also reacts readily with cold water, most acids,
and most nonmetals, such as sulfur and phosphorus.
Reaction with water


When water combines with water, it forms slaked
lime, or calcium hydroxide (Ca(OH) 2 ):
Calcium + Water ---> calcium Hydroxide +
Hydrogen
Ca(s) + 2H2O(l) --->
Ca(OH)2(aq ) + H2(gas)
Reaction with Oxygen



Calcium burn vigorously with a brick-red
flame.
Calcium reacts with oxygen to make calcium
oxide, CaO.
2Ca (s) + O2(g)
2CaO (s)
Reaction with acids


when calcium reacts with hydrochloric acid it forms calcium chloride (an
alkali) and hydrogen gas as products :
Ca(s) + 2HCl(aq)
CaCl2(aq) + H2(g)
when calcium reacts with Sulphuric acid it forms calcium sulphate salts
and hydrogen gas as products:

Ca(s)+ H2SO4 (aq)
CaSO4 (s) + H2
Reaction with halogens


Calcium burns in the halogens to form the corresponding
calcium halide.
calcium (Ca) and the halogen bromine (Br) form the ionic
compound calcium bromide (CaBr2).
Ca (s) + Br2 (g)
 Calcium & the

Ca Br2
halogen fluorine form calcium fluoride
Ca(s) + F2(g)
Ca F2 (S)
Chemical Properties of Magnesium



Magnesium is a very reactive metal and does
not exist in a free state in nature.
It reacts with a slow pace with cold water and
at a very rapid pace with hot water.
The oxidation process of magnesium is very
rapid and if kept in open, a layer of oxidized
magnesium is formed on the surface of the
metal. Magnesium also burns very rapidly,
when it is at room temperature.
Reaction with water

Magnesium reacts slowly with cold water to
form a solution of magnesium hydroxide &
hydrogen:
Mg (s) + 2H2O(l) ->
Mg(OH)2(aq) + H2(g)
The reaction with steam occurs readily,
however, & the products are magnesium
oxide & hydrogen:
Mg (s) + H2O(g) ->
MgO(s) + H2(g)
Reaction with Oxygen

The magnesium is burned and reacts with the
oxygen in the air.
Magnesium + Oxygen gas → Magnesium
oxide
2Mg (s)+ O2 (g) →
2MgO (s)
Reaction with acids
Magnesium undergoes vigorous reactions with
dilute acids, to produce the corresponding salt
(magnesium cloride; magnesium sulphate) &
hydrogen
Mg (s) + 2HCl(aq) -->
MgCl 2 (aq) + H2(g)
Mg (s) + H2SO4 (aq) ->
MgSO4 (aq) + H2(g)
Reaction with halogens


Magnesium reacts readily with chlorine to
produce magnesium chloride
Mg (s) + Cl2(g) ->
MgCl2(s)
Magnesium reacts readily with bromine to
produce magnesium bromide
Mg (s) + Br2(g) ->
MgBr2(s)
Chemical Properties of Aluminium


The surface of aluminum metal is covered
with a thin layer of oxide that helps protect the
metal from attack by air.
So, normally, aluminum metal does not react
with air. If the oxide layer is damaged, the
aluminum metal is exposed to attack.
Reaction with water




The protective oxide layer prevents the
reaction with water. However, if the oxide
layer is removed then aluminium reacts wiyh
cold water to form aluminium hydroxide:
2 Al(s) + 6 H2O(l)
2Al(OH)3 (s)+ H2 (g)
Aluminium reacts readily with steam, but
aluminium oxide is formed:
2 Al(s) + 3 H2O(g)
Al2 O3(s) + 3H2(g)
Reaction with Oxygen


Aluminium burns with a characteristic
brilliant white flame & the oxide is formed:
2 Al(s) + 3 O2(g)
2Al2 O3(s)
Reaction with acids




Aluminium reacts slowly @ first & then speeds
up after the oxide layer has been removed &
the metal is exposed.
Aluminium reacts with dilute acids to produce
the corresponding (aluminium cloride;
aluminium sulphate) salt:
2Al(s) + 6HCl(aq)
2AlCl3(aq)+ 3H2(g)
2Al(s) + 3H2SO4(aq)
Al2(SO4)3 aq+ 3H2(g)
Reaction with halogens


Aluminium reacts vigorously with all
halogens to form the aluminium halide:
2 Al(s) + 3 Cl2 (g)
Al2 Cl3(s)
Reaction with alkali


Aluminium dissolves in sodium hydroxide
with the evolution of hydrogen gas, H2, and
the formation of aluminates of the type
[Al(OH)4]-.
2Al(s) + 2NaOH(aq) + 6H2O →
2Na+(aq) + 2[Al(OH)4]- + 3H2(g)
Reaction with alkali cont’d.


Aluminium reacts with sodium (or potassium)
hydroxide to form a complex salt (an
aluminate: [Al(OH)4]- ) & hydrogen:
2 Al(s) + 2NaOH + 6H2 O (l) ……
….
2NaAl(OH)4 (aq)+ 3H2 (g)
Chemical properties of Zinc





Reacts with water
Reacts with dilute acids
Reacts with oxygen
Reacts with halogens
Reacts with alkali
Reaction with water


Zinc does not react with water; however, redhot zinc will react with steam to produce zinc
oxide & hydrogen:
2Zn(s) + H2 O (g) →
ZnO (s) + H2 (g)
Reaction with Oxygen


Zinc metal tarnishes in moist air. Zinc metal
burns in air to form the white zinc(II) oxide, a
material that turns yellow on prolonged
heating.
2Zn(s) + O2(g) →
2ZnO(s) [white]
Reaction with acids

Zinc metal dissolves slowly in dilute acids to
form solutions containing the aquated Zn(II)
ion to produce the corresponding salt and
hydrogen gas, H2.

Zn(s) + H2SO4(aq) →

Zn(s) + 2HCl(aq) →
ZnSO4(aq)+H2(g)
ZnCl2(aq)+H2(g)
Reaction with halogens

Zinc dibromide, zinc(II) dibromide, ZnBr2, and
zinc diiodide, zinc(II) diiodide, NiI2, are
formed in the reactions of zinc metal and
bromine, Br2, or iodine, I2.

Zn(s) + Br2(g) →

Zn(s) + I2(g) →
ZnBr2(s) [white]
ZnI2(s) [white]
Reaction with alkali

Zinc metal dissolves in aqueous alkalis such as
potassium hydroxide, KOH, to form zincates
such as [Zn(OH)4]2-. The complex salts formed
is called potassium zincate.
Zn(s) + 2KOH(aq) + 2H2 O (l)
K2 Zn(OH)4
(aq) +
H2 (g)
Chemical properties of Iron





Reacts with steam water
Reacts with dilute acids
Reacts with oxygen
React with halogens
Does not react with alkali
Reaction with water



In the absence of air, cold water doesn’t react
with iron. In the presence of air , however,
rusting occurs.
Hot iron react with steam :
3Fe(s) + 4H2O(g) →
Fe3O4 (s) + 4H2(g)
Reaction with Oxygen




Iron metal reacts in moist air by oxidation to give
a hydrated iron oxide. This does not protect the
iron surface to further reaction since it flakes off,
exposing more iron metal to oxidation. This
process is called rusting and is familiar to any car
owner.
On heating with oxygen, O2, the result is
formation of the iron oxides Fe2O3 and Fe3O4.
4Fe(s) + 3O2(g) →
2Fe2O3(s)
3Fe(s) + 2O2(g) →
Fe3O4(s)
Reaction with acids



Iron metal dissolves readily in dilute sulphuric
acid in the absence of oxygen to form solutions
containing the aquated Fe(II) ion together with
hydrogen gas, H2. In practice, the Fe(II) is
present as the complex ion [Fe(OH2)6]2+.
Fe(s) + H2SO4(aq) →
Fe2+(aq) +
SO42- (aq) + H2(g)
If oxygen is present, some of the Fe(II)
oxidizes to Fe(III).
Reaction with halogens





Iron reacts with excess of the halogens F2, Cl2,
and Br2, to form ferric, that is, Fe(III), halides.
Iron(III) fluoride; Iron(III) chloride; Iron(III)
bromide.
2Fe(s) + 3F2(g) →
2Fe(s) + 3Cl2(g) →
2Fe(s) +3Br2(l) →
2FeF3(s) (white)
2FeCl3(s) (dark brown)
2FeBr3(s) (reddish brown)
Chemical properties of Copper





Does not react with water
Does not react with alkali
Reacts with dilute acids
Reacts with oxygen
Reacts with halogens
Reaction with water/alkali


Copper doesn’t react with cold water or with
steam.
Copper doesn’t react with alkali
Reaction with Oxygen


Copper metal is stable in air under normal
conditions. Copper does not burn in oxyen;
however, a black oxide layer (copper (II)oxide)
is formed on surface of the metal:
2Cu(s) + O2(g) →
2CuO(s)
Reaction with acids


Copper metal dissolves in hot concentrated
sulphuric acid to form solutions containing the
aquated Cu(II) ion together with hydrogen
gas, H2.
Cu(s) + H2SO4(aq) →
Cu2+(aq) + SO42-(aq)
+ H2(g)
Reaction with acids cont’d


Copper metal also dissolves in dilute or
concentrated nitric acid, HNO3.
3Cu(s) + 8HNO3(aq) →
3Cu(NO3)2 (aq) + 2NO (g) +4 H2O(l)
Reaction with halogens




The reaction between copper metal and the
halogens fluorine, F2, chlorine, Cl2, or bromine,
Br2, affords the corresponding dihalides
copper(II) fluoride, CuF2, copper(II) chloride,
CuCl2, or copper(II) bromide,
CuBr2 respectively.
Cu(s) + F2(g) →
CuF2(s) [white]
Cu(s) + Cl2(g) →
CuCl2(s) [yellow-brown]
Cu(s) + Br2(g) →
CuBr2(s) [black]
Chemical Properties of Lead

Doesn’t react with water
Doesn’t react with dilute acids
Doesn’t react with alkali
Doesn’t react readily with halogens

Reacts with oxygen



No Reactions with:






Water
The surface of metallic lead is protected by a thin layer of
lead oxide, PbO. It does not react with water under normal
conditions.
Dilute acids
The surface of metallic lead is protected by a thin layer of
lead oxide, PbO. This renders the lead essentially insoluble in
sulphuric acid, and so, in the past, a useful container of this
acid. Lead reacts slowly with hydrochloric acid and nitric
acid, HNO3. In the latter case, nitrogen oxides are formed
together with lead(II) nitrate, Pb(NO3)2.
Halogens
Alkali
Reaction with Oxygen


The surface of metallic lead is protected by a
thin layer of lead oxide, PbO. Only upon
heating lead to 600-800°C does lead react with
oxygen in air to from lead (II) oxide, PbO
(a yellow solid).
2Pb(s) + O2(g) →
2PbO(s)
The reactivity of metals

Reactivity series

What makes a metal reactive ?

Displacement of metals
The Reactivity Series of Metals



Although most metals are usually electropositive in nature
and lose electrons in a chemical reaction they do not react
with the same vigor or speed.
Metals display different reactions towards different
substances. The greater the ease with which an element
loses its electrons and acquires a positive charge, the greater
is its reactivity. Further, the greater the number of shells and
lesser the number of valence electrons, the greater is the
reactivity of the metal.
The activity series of metals, arranges all metals in order of
their decreasing chemical activity. As we go down the
activity series from potassium to gold the ease with which a
metal loses electrons, and forms positive ions in solutions,
decreases.
The Most Reactive and Least Reactive
Metals

The most active metal,
potassium, is at the top
of the list and the least
reactive metal, gold, is
at the bottom of the list.
Although hydrogen is a
non-metal it is included
in the activity series
due to the fact that it
behaves like a metal in
most chemical reactions
i.e., the hydrogen ion
has a positive charge
[H+] like other metals.
Cont’d.

Two non-metals, carbon and hydrogen are
important chemical reference points with
regard to the method of metal extraction and
reactivity towards acids. Metals above carbon
cannot be extracted by carbon reduction and
are usually extracted by electrolysis. Metals
below hydrogen will not displace hydrogen
from acids.
The Most Reactive and Least Reactive
Metals Cont’d.



The higher the metal in the series, the more
reactive it is i.e., its reaction is fast and more
exothermic.
This also implies that the reverse reaction
becomes more difficult i.e., the more reactive a
metal, the more difficult it is to extract from its
ore. The metal is also more susceptible to
corrosion with oxygen and water.
The reactivity series can be established by
observation of the reaction of metals with water,
oxygen or acids.
Within the general reactivity or activity series there
are some periodic table trends:



Down Group 1 (I) the "Alkali Metals" the activity
increases Cs > Rb > K > Na > Li. Down Group 2
(II) the activity increases e.g., Ca > Mg.
On the same period, the Group 1 metal is more
reactive than the group II metal, and the group II
metal is more reactive than the Group III metal,
and all three are more reactive than the
"Transition Metals".
e.g., Na > Mg > Al (on Period 3) and K > Ca > Ga
> Fe/Cu/Zn etc. (on Period 4)
The chemical properties of
compounds of metals.
The action of heat on the hydroxides; on
the nitrates; on the carbonates
The action of water on the oxides; on the
hydroxides; on the carbonates
Reduction of metal oxides
Reaction of metal oxides & hydroxides
with dilute acids.

The action of heat on the hydroxides




Potassium & sodium are stable; there is no
decomposition.
The metal hydroxide of Calcium, Magnesium,
Aluminium, Zinc, Iron, Lead, Copper, Silver
decomposes to form the corresponding oxide
& water.
Mg(OH)2(s) ->heat
MgO (s) + H2O(g)
Cu(OH)2(s) ->heat
CuO (s) + 2H2O(g)
The action of heat on the nitrates

Potassium nitrate & sodium nitrate
decompose to form the nitrite & oxygen gas:

2KNO3(s) -> heat
2KNO2 (s) + O2 (g)

2NaNO3(s) ->heat
2NaNO2 (s) + O2 (g)
The action of heat on the nitrates Cont’d





Nitrates of calcium, magnesium, aluminium,
zinc, iron, lead , copper decompose to form
the oxide, nitrogen dioxide (poisonous gas) &
oxygen.
heat 2ZnO (s) +4NO (g) +O (g)
2Zn(NO3) 2(s) ->
2
2
2Pb(NO3) 2(s) ->
2PbO (s)+4NO2 (g) +O2 (g)
heat
Silver nitrate decomposes to give silver metal
2AgNO3(s) ->
2Ag (s) + 2NO2 (g) + O2 (g)
heat
The action of heat on the carbonates





Potassium carbonate & sodium carbonate are
stable. There is no decomposition.
Carbonates of Ca, Mg, Al, Zn, Fe, Pb, Cu, Ag
decompose to form the metal oxide & carbon
dioxide.
CaCO3(s) -> heat
CaO (s) + CO2 (g)
ZnCO3(s) -> heat
ZnO (s) + CO2 (g)
CuCO3(s) ->heat
CuO (s) + CO2 (g)
The action of water on the oxides

Oxides of potassium, sodium, calcium,
magnesium are soluble in water forming a
solution of corresponding hydroxide (an alkali).

K2O (s) + H2O(l) ->
2KOH (aq)

CaO (s) + H2O(l) ->
Ca(OH) 2 (s)

Oxides of Al, Zn, Fe, Pb, Cu, Ag do not react
with water.
The action of water on the hydroxides






Potassium, sodium :
Very soluble in water
Calcium, magnesium :
slightly soluble in water
Aluminium, zinc, iron, lead, copper, silver:
Insoluble in water
The action of water on the carbonates




Potassium, sodium :
Soluble in water
Calcium, magnesium, aluminum, zinc, iron,
lead, copper, silver :
Insoluble in water
Reduction of metal oxides


Metal oxides can be reduced by hydrogen or
carbon to the corresponding metal; however,
this reaction occurs particularly with the
oxides of the less reactive metals.
CuO (s) + H2 (g) ->
Cu (s) + H2O (l)
Reduction of metal oxides cont’d.


Carbon can be used to reduce the oxides of
zinc & those metals that are lower in the
reactivity series, while hydrogen wil reduce
the oxides of iron, lead, copper & silver to the
corresponding metal.
Fe2 O3(s) + 3C (s) ->
2Fe (s) + 3CO(g)
Reaction of metal oxides &
hydroxides with dilute acids



Metal oxides & hydroxides are classified as
bases, so they will react with acids to produce
a salt & water only. This is called a neutralisation reactions.
NaOH(aq) +HCl (aq)
Ca(OH)2(s) + H2SO4 (aq) ->
+ 2H2O(l)
NaCl (aq) + H2O (l)
CaSO4 (s)
Reaction of metal oxides &
hydroxides with dilute acids


CuO(s) + H2SO4 (aq) ->
+ H2O(l)
MgO(s) + 2HCl (aq) ->
+ H2O(l)
cont’d.
CuSO4 (aq)
MgCl2 (aq)
Chemical reactions of some
non-metals
Hydrogen
Carbon
 Nitrogen
Oxygen
 Silicon
Sulfur
Chlorine

Chemical properties of H2
Reaction of hydrogen with air
 Hydrogen is a colourless gas, H2, that is lighter
than air. Mixtures of hydrogen gas and air do
not react unless ignited with a flame or spark,
in which case the result is a fire or explosion
with a characteristic reddish flame whose only
products are water, H2O.
 2H2(g) + O2(g) → 2H2O(l)
Reaction of hydrogen with the halogens




Hydrogen gas, H2, reacts with fluorine, F2, in
the dark to form hydrogen fluoride.
H2(g) + F2(g) → 2HF(g)
Hydrogen gas reacts with chlorine, Cl2, to
form hydrogen chloride .
H2(g) + Cl2(g) → 2HCl(g)
Reaction of hydrogen with nitrogen
& sulphur





Hydrogen reacts with nitrogen & sulphur to
give respectively ammonia & hydrogen
sulfide
H2(g) + N2(g) → 2NH3 (g)
H2(g) + S(s) → H2S (g)
Hydrogen acts as reducing agent since it
removes oxygen from the oxides of metal:
H2(g) + CuO (s) → Cu (s) + H2O (l)
Chemical properties of C




Carbon, as graphite, burns to form gaseous
carbon dioxide, CO2. Diamond is a form of
carbon and also burns in air when heated to
600-800°C - an expensive way to make carbon
dioxide!
C(s) + O2(g) → CO2(g)
When the air or oxygen supply is restricted,
incomplete combustion to carbon monoxide,
CO, occurs.
2C(s) + O2(g) → 2CO(g)
Chemical properties of C
cont’d.



Carbon is a good reducing agent since it
removes the oxygen from metal oxides to form
the metal & carbon dioxide:
2PbO (s) + C (s)……..2Pb (s) + CO2(g)
Carbon does not react with dilute acids &
alkalis
Chemical properties of N2

Nitrogen combine with oxygen & hydrogen
to form nitrogen monoxide & ammonia,
respectively

N2(g) + O2(g) → 2NO (g)

N2(g) + 3H2(g) → 2NH3 (g)
Chemical properties of O2

Oxygen is needed for combustion & therefore
reacts with metals & non-metals to produce
oxides.

O2(g) + Cu(s) → CuO2(g)

O2(g) + 2Ca(s) → 2CaO(s)
Chemical properties of Si





Silicon reacts vigorously with all the halogens to
form silicon tetrahalides. So, it reacts with
fluorine, F2, chlorine, Cl2, bromine, I2, and iodine,
I2, to form respectively silicon(IV) fluoride, SiF4,
silicon(IV) chloride, SiCl4, silicon(IV) bromide,
SiBr4, and silicon(IV) iodide, SiI4. The reaction
with fluorine takes palce at room temperature but
the others requiring warming over 300°C.
Si(s) + 2F2(l) → SiF4(g)
Si(s) + 2Cl2(l) → SiCl4(g)
Si(s) + 2Br2(l) → SiBr4(l)
Si(s) + 2I2(l) → SiI4(s)
Reaction of silicon with air




The surface of lumps of silicon is protected by a
very thin layer of silicon dioxide, SiO2. This
renders silicon more or less inert to further
oxidation by air even up to about 900°C. After
this, reaction with oxygen in the air gives silicon
dioxide. At temperatures above about 1400°C,
silicon reacts with nitrogen, N2, in the air as well
as oxygen, to form the silicon nitrides SiN and
Si3N4.
Si(s) + O2(g) → SiO2(s)
2Si(s) + N2(g) → 2SiN(s)
3Si(s) + 2N2(g) → Si3N4(s)
Reaction of silicon with bases


Silicon is attacked by bases such as aqueous
sodium hydroxide to give silicates, highly
complex species containing the anion [SiO3]2-.
Si(s) + 2NaOH(aq) +HO (l)→ Na2 SiO3(aq) +
2H2(g)
Chemical properties of S



Sulphur burns in air to form the gaseous
sulphur dioxide, SO2.
S(s) + O2(g) → SO2(g)
Sulphur is also oxidised to sulphur dioxide in
the presence of oxidising agents such as
sulfuric acid.
Chemical properties of S
Cont’d.




Sulphur combines with most metals to form
the metal sulphide :
S(s) + Fe (s)………. FeS (s)
Hydrogen reacts with sulphur to produce
hydrogen sulpide. Hydrogen sulphide has a
characteristic odour loke that of rotten eggs.
Sulphur does not react with either dilute
acids or alkalis.
S.E. Whitnall, 1933
The primary duty of the University to a student is
to provide him with such instructors as will make
him realize that the responsibility for progress is
his own and no one else's.