Inorganic Chemistry

Periodicity
Period 3
Properties of the Elements in Period 3

 Elements in the periodic table are arranged in order of increasing atomic

number and placed in vertical columns (groups) and horizontal rows
(periods)
 The elements across the periods show repeating patterns in chemical
and physical properties
 This is called periodicity
Atomic radius

 The atomic radius is the distance between the nucleus and the
outermost electron of an atom
 The atomic radius is measured by taking two atoms of the same
element, measuring the distance between their nuclei and then halving
this distance
 In metals this is also called the metallic radius and in non-metals,
the covalent radius

The atomic radius gives a measure of the size of atoms

Atomic radii of Period 3 elements table

 Across the period, the atomic radii decrease

 This is because the number of protons (the nuclear charge) and the
number of electrons increases by one every time you go an element to
the right
 The elements in a period all have the same number of shells (so
the shielding effect is the same)
 This means that as you go across the period the nucleus attracts the
electrons more strongly pulling them closer to the nucleus
 Because of this, the atomic radius (and thus the size of the
atoms) decreases across the period
Ionic radius

 The ionic radius is the distance between the nucleus and the outermost
electron of an ion
 Metals produce positively charged ions (cations) whereas nonmetals
produce negatively charged ions (anions)
 The cations have lost their valence electrons which causes them to be
much smaller than their parent atoms
 Because there are less electrons, this also means that there is
less shielding of the outer electrons
 Going across the period from Na+ to Si4+ the ions get smaller due to
the increasing nuclear charge attracting the outer electrons in
the second principal quantum shell nucleus (which has an increasing
atomic number)
 The anions are larger than their original parent atoms because each
atom has gained one or more electrons in their third principal
quantum shell
 This increases the repulsion between electrons, while the nuclear
charge is still the same, causing the electron cloud to spread out
 Going across P3- to Cl- the ionic radii decreases as the nuclear charge
increases across the period and fewer electrons are gained by the atoms
(P gains 3 electrons, S 2 electrons and Cl 1 electron)

Melting point

 A general increase in melting point for the Period 3 elements up to

silicon is observed
 Silicon has the highest melting point
 After the Si element the melting points of the
elements decreases significantly
Electrical conductivity

 Electrical conductivity refers to how well a substance can

conduct electricity
 Unlike the melting points, the electrical conductivity of the Period 3
elements shows a clear trend
 Going across the period, the electrical conductivity of the
elements decreases significantly

Bonding & structure of the elements table

 The table shows that Na, Mg and Al are metallic elements which form
positive ions arranged in a giant lattice in which the ions are held
together by a 'sea' of delocalised electrons around them

Metal cations form a giant lattice held together by electrons that can
freely move around

 The electrons in the ‘sea’ of delocalised electrons are those from

the valence shell of the atoms
 Na will donate one electron into the ‘sea’ of delocalised
electrons, Mg will donate two and Al three electrons
 As a result of this, the metallic bonding in Al is stronger than in Na
 This is because the electrostatic forces between a 3+ ion and the larger
number of negatively charged delocalised electrons is much larger
compared to a 1+ ion and the smaller number of delocalised electrons
in Na
 Because of this, the melting points increase going from Na to Al

 Si has the highest melting point due to its giant molecular structure in
which each Si atom is held to its neighbouring Si atoms by strong
covalent bonds
  P, S, Cl and Ar are non-metallic elements and exist as simple
molecules (P4, S8, Cl2 and Ar as a single atom)
 The covalent bonds within the molecules are strong,
however, between the molecules, there are only weak instantaneous
dipole-induced dipole forces
 It doesn’t take much energy to break these intermolecular forces
 Therefore, the melting points decrease going from P to Ar (note that
the melting point of S is higher than that of P as sulphur exists as larger
S8 molecules compared to the smaller P4 molecule)

Electrical conductivity

 The electrical conductivity decreases going across the Period 3

elements

Electrical conductivity decreases Period 3 elements table

 Going from Na to Al, there is an increase in the number of valence

electrons that are donated to the ‘sea’ of delocalised electrons
 Because of this, in Al there are more electrons available to move around
through the structure when it conducts electricity, making Al a better
electrical conductor than Na

 Due to the giant molecular structure of Si, there are no delocalised

electrons that can freely move around within the structure
 Si is therefore not a good electrical conductor and is classified as
a semimetal (metalloid)
 The lack of delocalised electrons is also why P and S cannot conduct
electricity
Reaction of Period 3 elements with oxygen table

Reaction of Period 3 elements with chlorine table

Reaction of sodium & magnesium with water

 Sodium reacts vigorously with cold water:

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

 The sodium melts into a ball and moves across the water surface until it
disappears
 Hydrogen gas is given off
 The solution formed is strongly alkaline (pH 14) due to the sodium
hydroxide which is formed
Reaction of Period 3 Oxides & Water

 Not all Period 3 oxides react with or are soluble in water

Reaction of Period 3 oxides with water table

Acid / Base Behaviour of Period 3 Oxides & Hydroxides

Period 3 oxides

 Aluminium oxide is amphoteric which means that it can act both as a

base (and react with an acid such as HCl) and an acid (and react with a
base such as NaOH)

Acidic & basic nature of the Period 3 oxides

Reactions of the Period 3 oxides with acid/base table

 The acidic and basic nature of the Period 3 elements can be explained
by looking at their structure, bonding and the Period 3
elements’ electronegativity
Structure, bonding & electronegativity of the Period 3 elements table

 The difference in electronegativity between oxygen and Na, Mg and Al

is the largest
 Electrons will therefore be transferred to oxygen when forming oxides
giving the oxide an ionic binding
 The oxides of Si, P and S will share the electrons with the oxygen to
form covalently bonded oxides
 The giant ionic and giant covalent structured oxides will have high
melting points as it is difficult to break the structures apart
Period 3 hydroxide

 NaOH is a strong base and will react with acids to form a salt and water:

NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

 Mg(OH)2 is also a basic compound which is often used in indigestion

remedies by neutralising the excess acid in the stomach to relieve pain:

Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + 2H2O(l)

 Al(OH)3 is amphoteric and can acts both as an acid and base:

Al(OH)3(s) + 3HCl(aq) → AlCl3(s) + 3H2O(l)

Al(OH)3(s) + NaOH(aq) → NaAl(OH)4(aq)

Reaction of Period 3 Chlorides & Water

 Chlorides of Period 3 elements show characteristic behaviour when

added to water which can be explained by looking at their chemical
bonding and structure

Chemical bonding & structure of Period 3 chlorides table

Sodium & magnesium chloride

 NaCl and MgCl2 do not react with water as the polar water molecules
are attracted to the ions dissolving the chlorides and breaking down
the giant ionic structures: the metal and chloride ions become
hydrated ions

The diagram shows water molecules breaking down the giant ionic
structure of NaCl and MgCl2 to form hydrated ions
Aluminium chloride
 Aluminium chloride exists in two forms:
o AlCl3 as a giant lattice and with ionic bonds
o Al2Cl6 as a dimer with covalent bonds

The two forms in which aluminium chloride exists

 When water is added to aluminium chloride the dimers are broken

down and Al3+ and Cl- ions enter the solution
 The highly charged Al3+ ion becomes hydrated and causes a water
molecule that is bonded to the Al3+ to lose an H+ ion which turns the
solution acidic
 The H+ and the Cl- form hydrogen chloride gas which is given off
as white fumes

Silicon chloride

 SiCl4 is hydrolysed in water, releasing white fumes of hydrogen

chloride gas in a rapid reaction

SiCl4(l) + 2H2O(l) → SiO2(s) + 4HCl(g)

 The SiO2 is seen as a white precipitate and some of the hydrogen

chloride gas produced dissolves in water to form an acidic solution
Phosphorus(V) chloride

 PCl5 also gets hydrolysed in water

PCl5(s) + 4H2O(l) → H3PO4(aq) + 5HCl(g)

 Both H3PO4 and dissolved HCl are highly acidic

Period 3: Trends in Electronegativity & Bonding

Electronegativity

 Electronegativity is the power of an element to draw the electrons

towards itself in a covalent bond
 Going across the period, the electronegativity of the elements
increases

 As the atomic number increases going across the period, there is

an increase in nuclear charge
 Across the period, there is an increase in the number of valence
shells however the shielding is still the same as each extra electrons
enters the same shell
 As a result of this, electrons will be more strongly attracted to the
nucleus causing an increase in electronegativity across the period

Bonding & structure of Period 3 elements

 The table shows that going from Al to S the bonding changes

from metallic to covalent and the structure changes
from giant to simple structure
 Na, Mg and Al are metallic elements which form positive ions arranged
in a giant lattice in which the ions are held together by a ‘sea’ of
delocalised electrons around them
 Since Al donates three electrons into the sea of delocalised electrons to
form an ion with +3 charge, the electrostatic forces between the
electrons and the aluminium ion will be very strong
 The electrons in the ‘sea’ of delocalised electrons are those from
the valence shell of the atoms
 Na will donate one electron into the ‘sea’ of delocalised
electrons, Mg will donate two and Al three electrons
  As a result of this, the metallic bonding in Al is stronger than in Na
 This is because the electrostatic forces between a 3+ ion and the larger
number of negatively charged delocalised electrons are much larger
compared to a 1+ ion and the smaller number of delocalised electrons
in Na
 Since there are more electrons in a metallic lattice of aluminium
compared to sodium and magnesium, aluminium is a better electrical
conductor

Metal cations form a giant lattice held together by electrons that can
freely move around

 Si is a non-metallic element and has a giant molecular structure in

which each Si atom is held to its neighbouring Si atoms by strong
covalent bonds
 There are no delocalised electrons in the structure of Si which is why
silicon cannot conduct electricity and is classified as a metalloid

The diagram shows the giant molecular structure of silicon where silicon
atoms are held together by strong covalent bonds

 Phosphorus, sulfur, chlorine argon re both non-metallic elements that

exist as simple molecules (P4 , S8 , Cl2 and Ar as single atoms)
 The covalent bonds within the molecules are strong,
however, between the molecules there are only weak instantaneous
dipole-induced dipole forces
 It doesn’t take much energy to break these intermolecular forces
 The lack of delocalised electrons means that these compounds cannot
conduct electricity
Period 3 chlorides

 The bonding and structure of the Period 3 elements are summarised in

the table below:

 The table shows that Na, Mg and Al are metallic elements which form
positive ions arranged in a giant lattice in which the ions are held
together by a ‘sea’ of delocalised electrons around them
 The electrons in the ‘sea’ of delocalised electrons are those from
the valence shell of the atoms
 Na will donate one electron into the ‘sea’ of delocalised
electrons, Mg will donate two and Al three electrons
 As a result of this, the metallic bonding in Al is stronger than in Na
 This is because the electrostatic forces between a 3+ ion and the larger
number of negatively charged delocalised electrons is much larger
compared to a 1+ ion and the smaller number of delocalised electrons
in Na
 Because of this, the melting points increase going from Na to Al

 Si has the highest melting point due to its giant molecular structure in
which each Si atom is held to its neighbouring Si atoms by strong
covalent bonds

 P, S, Cl and Ar are non-metallic elements and exist as simple

molecules (P4, S8, Cl2 and Ar as single atom)
 The covalent bonds within the molecules are strong,
however between the molecules there are only weak instantaneous
dipole-induced dipole forces
  It doesn’t take much energy to break these intermolecular forces
 Therefore, the melting points decrease going from P to Ar (note that
the melting point of S is higher than that of P as sulphur exists as larger
S8 molecules compared to the smaller P4 molecule)
 The presence of a ‘sea’ of delocalised electrons also determines whether
the element is a good conductor or not
 Going across the period the electrical conductivity of the elements
decreases due to lack of delocalised electrons

 The electronegativities of the Period 3 elements therefore determines

the chemical bonding and structure of their chlorides and oxides

 Going across Period 3, their chlorides and oxidised become

more covalent and their structure shifts from a giant ionic to a simple
molecular structure
 Their reactions with water become more vigorous as a result of this as it
becomes easier to hydrolyse the chlorides and oxides

Periodicity: Predicting Position & Properties

 If the chemical and physical properties of an element are known,
the position of that element in the Periodic Table can be predicted
 Similarly, predictions can be made about
the physical and chemical properties of elements if the position of the
element in the Periodic Table is known

Break the question down and systematically approach the question

 Step 1: ‘Element X forms a chloride, which reacts with water to form a

solution of pH 1’

The low pH of the solution formed suggests that the chloride is a non-
metallic chloride (group 13 to 17)
 Step 2: ‘The oxide does not dissolve in or react with aqueous sodium
hydroxide’

Since aluminium oxide does reaction with sodium hydroxide, element G

cannot be Group 13

 Step 3: It forms an oxide which has a melting point of 1610 °C’

This suggests a giant molecular (covalent) structure which corresponds

to Group 14

 Step 4: Element X cannot be carbon (which is in Group 14) as carbon

dioxide is a gas whereas the element X oxide is a solid (with a melting
point of 1610 °C)

 Step 5: Element X is therefore a Group 14 element in Period 3 or lower

 Note that this is an example of predicting the position of an element

based on its physical and chemical properties
Group 2
 The Group 2 elements react with oxygen, water and dilute acids

Group 2 reactions table

Group 2 reactions with oxygen & water chemical equations

Group 2 reactions with dilute hydrochloric acid & dilute sulfuric acid chemical equations

 The reaction of all metals with oxygen follows the following general equation:

2M (s) + O2 (g) → 2MO (s)

Where M is any metal in Group 2

Remember that Sr and Ba also form MO2

 The reaction of all metals with water follows the following general equation:
 M (s) + 2H2O (l) → M(OH)2 (s) + H2 (g)

The exceptions to this general equation are:

 Be which does not react with water

 Mg which forms MgO (s) and H2 (g)

 The reaction of all metals with dilute HCl follows the following general equation:

M(s) + 2HCl(aq) → MCl2(aq) + H2(g)

 The reaction of all metals with dilute H2SO4 follows the following general equation:

M(s) + H2SO4(aq) → MSO4(aq) + H2(g)

Remember that CaSO4, SrSO4 and BaSO4 are insoluble

Reactions of Group 2 Oxides, Hydroxides & Carbonates

Reactions of Group 2 oxides with water

 All Group 2 oxides are basic, except for BeO which is amphoteric (it can
act both as an acid and base)
 Group 2 oxides react with water to form alkaline solutions which
generally get more alkaline going down the group
o This happens because the hydroxides that form become more
soluble as you move down the group
o This means that more hydroxide ions, OH–, dissociate into the
solution causing the pH to increase

Group 2 oxide reactions with water table

 Remember that:
 oxide + water → hydroxide

And that calcium hydroxide is also called limewater

Reactions of Group 2 metals with acid

 The Group 2 metals will react with dilute acids to form colourless
solutions of metal salts
o For example, they will form colourless solutions of metal chlorides
if reacted with hydrochloric acid
 When metals react with an acid, the by-product of this reaction is
hydrogen gas

Group 2 metal element reactions with dilute acids table

 When some of Group 2 metals react with sulfuric acid rather than
hydrochloric, an insoluble sulfate forms
 Going down the group, the Group 2 sulfates become less and less
soluble
o Calcium sulfate is sparingly soluble, but strontium sulfate and
barium sulfate are insoluble
o
Reactions of Group 2 oxides with acid

 Group 2 sulfates also form when a Group 2 oxide is reacted with an acid
 The insoluble sulfates form at the surface of the oxide, which means
that the solid oxide beneath it can’t react with the acid
 This can be prevented to an extent by using the oxide in powder form
and stirring, in which case neutralisation can take place
 Remember that:

oxide + dilute hydrochloric acid → salt + water

oxide + dilute sulfuric acid → sulfate + water

Reactions of group 2 hydroxides

 The Group 2 metal hydroxides form colourless solutions of metal

chlorides when they react with a dilute acid
 The sulfates decrease in solubility going down the group (barium
sulfate is an insoluble white precipitate)

Group 2 hydroxide reactions with dilute acids table

 Remember that:

hydroxide + dilute acid → salt + water

hydroxide + dilute sulfuric acid → sulfate + water

Reactions of group 2 carbonates

 All Group 2 carbonates (except for BeCO3) are insoluble in water

 All Group 2 carbonates will form soluble chloride salts, water and
carbon dioxide gas when reacted with dilute hydrochloric acid
 When reacted with sulfuric acid, the carbonates of Ca, Sr and Ba form
an insoluble sulfate layer on their surface which stops any further
reaction after the initial bubbling (effervescence) of carbon dioxide
gas is seen

Group 2 carbonate reactions with dilute acids

 Remember that:

carbonate + dilute hydrochloric acid → salt + water + carbon dioxide

carbonate + dilute sulfuric acid → sulfate + water + carbon dioxide

Group 2: Thermal Decomposition of Nitrates & Carbonates

Thermal decomposition is the breakdown of a compound into two or more

different substances using heat

Thermal decomposition of carbonates

 The Group 2 carbonates break down (decompose) when they are

heated to form the metal oxide and give off carbon dioxide gas
 The general equation for the decomposition of Group 2 carbonates is:

XCO3 (s) XO (s) + CO2 (g)

X = Group 2 element

 Going down the group, more heat is needed to break down the
carbonates

MgCO3 (s) MgO (s) + CO2 (s)

Thermal decomposition of nitrates

 Group 2 nitrates also undergo thermal decomposition

 Group 2 nitrates decompose to form the metal oxide, nitrogen dioxide
gas and oxygen gas
 The general equation for the decomposition of Group 2 nitrates is:

X(NO3)2 (s) XO (s) + 2NO2 (g) + ½O2 (g)

OR

2X(NO3)2 (s) 2XO (s) + 4NO2 (g) + O2 (g)

X = Group 2 element

 Nitrogen dioxide gas is observed as brown fumes and is toxic

 An example of this reaction is:

2Ca(NO3)2 (s) 2CaO (s) + 4NO2 (g) + O2 (g)

Trend in thermal stabilities

 Going down Group 2, more heat is needed to break down the

carbonate and nitrate ions
 The thermal stability of the Group 2 carbonates and nitrates
therefore increases down the group

 The smaller positive ions at the top of the groups will polarise the
anions more than the larger ions at the bottom of the group

 The small positive ion attracts the delocalised electrons in

the carbonate ion towards itself
 The higher the charge and the smaller the ion the higher the
polarising power

 The more polarised they are, the more likely they are to thermally
decompose as the bonds in the carbonate and nitrate ions
become weaker
Group 2: Physical & Chemical Trends

Chemical trends

 All elements in Group 2 (also called alkali earth metals) have the two
electrons in their outermost principal quantum shell
 All Group 2 metals can form ionic compounds in which they donate
these two outermost electrons (so they act as reducing agents) to
become an ion with +2 charge (so they themselves become oxidised)
 Going down the group, the metals become more reactive
 This can be explained by looking at the Group 2 ionisation energies:

 The first ionisation energy is the energy needed to remove the first
outer electron of an atom
 The second ionisation energy is the energy needed to remove the
second outer electron of an atom
 The graph above shows that going down the group, it becomes easier
to remove the outer two electrons of the metals
 Though the nuclear charge on the nucleus increases going down the
group (because there are more protons), factors such as an increased
shielding effect and a larger distance between the outermost
electrons and nucleus outweigh the attraction of the higher nuclear
charge
 As a result of this, the elements become more reactive going down the
group as it gets easier for the atoms to lose two electrons and become
2+ ions
 This trend is shown by looking at reactions of the Group 2 metals:

 With dilute hydrochloric acid: bubbles of hydrogen gas are given

off much faster indicating that the reactions become more
vigorous
 With oxygen hydrochloric acid: the metals get more reactive with
oxygen down the group (Ba is so reactive, that it must be stored in
oil to prevent it from reacting with oxygen in air)
Physical trends

 Going down the group, the elements become larger as the outer two
electrons occupy a new principal quantum shell which is further away
from the nucleus

 The melting point of the elements decreases going down the group as
the outer electrons get further away from the nucleus
 This means that the attraction between the nucleus and
the bonding electrons decreases causing a decrease in melting point
 As you go down the group, the density of the alkali earth metals
increases
Trends in Solubility in Group 2 Hydroxides & Sulfates

Group 2 hydroxides

 Going down the group, the solutions formed from the reaction of Group
2 oxides with water become more alkaline
 When the oxides are dissolved in water, the following ionic reaction
takes place:

O2- (aq) + H2O(l) → 2OH- (aq)

 The higher the concentration of OH- ions formed, the more alkaline the
solution
 The alkalinity of the formed solution can therefore be explained by
the solubility of the Group 2 hydroxides

 he hydroxides dissolve in water as follows:

X(OH)2 (aq) → X(aq) + 2OH- (aq)

Where X is the Group 2 element

 When the metal oxides react with water, a Group 2 hydroxide is formed
 Going down the group, the solubility of these hydroxides increases
 This means that the concentration of OH- ions increases, increasing the
pH of the solution
 As a result, going down the group, the alkalinity of the solution formed
increases when Group 2 oxides react with water
Group 2 sulfates

 The solubility of the Group 2 sulfates decreasing going down the group
Group 17
 The group 17 elements are called halogens
 The halogens have uses in water purification and as bleaches agents (chlorine), as flame-
retardants and fire extinguishers (bromine) and as antiseptic and disinfectant agents (iodine)

Colours

 All halogens have distinct colours which get darker going down the group

The colours of the Group 17 elements get darker going down the group

Volatility

 Volatility refers to how easily a substance can evaporate

o A volatile substance will have a low melting and boiling point

The melting & boiling points of the Group 17 elements increase going down the group which
indicates that the elements become less volatile

 Going down the group, the boiling point of the elements increases which means that
the volatility of the halogens decreases

 This means that fluorine is the most volatile and iodine the least volatile
Group 17: Trends in Bond Strength

 Halogens are diatomic molecules in which covalent bonds are formed by overlapping their
orbitals
 In a covalent bond, the bonding pair of electrons is attracted to the nuclei on either side and it is
this attraction that holds the molecule together
 Going down the group, the atomic size of the halogens increases
 The bonding pair of electrons get further away from the halogen nucleus and are therefore less
strongly attracted towards in

A covalent bond is formed by the orbital overlap of two atoms and the attraction of electrons
towards the nuclei; The bigger the atom, the weaker the covalent bond

 The bond strength of the halogen molecules therefore decreases going down the group

The bond enthalpies decrease indicating that the bond strengths decrease going down the
group

 Bond enthalpy is the heat needed to break one mole of a covalent bond
 The higher the bond enthalpy, the stronger the bond
 An exception to this is fluorine which has a smaller bond enthalpy than chlorine and bromine
 Fluorine is so small that when two atoms of fluorine get together their lone pairs get so close
that they cause significant repulsion counteracting the attracting between the bonding pair of
electrons and two nuclei

The lone pairs on fluorine get so close to each other in a fluorine molecule that they cause
repulsion which decreases the bond strength
Group 17: Dipole Forces & Volatility

 Halogens are non-metals and are diatomic molecules at room temperature

o This means that they exist as molecules which are made up of two similar atoms, such
as F2
 The halogens are simple molecular structures with weak van der Waals’ forces between the
diatomic molecules caused by instantaneous dipole-induced dipole forces

The diagram shows that a sudden distribution of electrons in a nonpolar molecule can cause
an instantaneous dipole. When this molecule gets close to another non-polar molecule it can
induce a dipole as the cloud of electrons repel the electrons in the neighbouring molecule to
the other side

 The more electrons there are in a molecule, the greater the instantaneous dipole-induced
dipole forces
 Therefore, the larger the molecule the stronger the van der Waals’ forces between molecules
 This is why as you go down the group, it gets more difficult to separate the molecules and
the melting and boiling points increase
 As it gets more difficult to separate the molecules, the volatility of the halogens decreases going
down the group

Going down the group, the van der Waals’ forces increase due to an increased number of
electrons in the molecules which means that the volatility decreases
Group 17: Oxidising Agents

 Halogens react with metals by accepting an electron from the metal atom to become an ion
with 1- charge

Eg. Ca(s) + Cl2(g) → Ca2+(Cl-)2(s)

 Halogens are therefore oxidising agents:

o Halogens oxidise the metal by removing an electron from the metal (the oxidation
number of the metal increases)
o Halogens become reduced as they gain an extra electron from the metal atom (the
oxidation number of the halogen decreases)
 The oxidising power of the halogens decreases going down the group (the halogens get less
reactive)
 This can be explained by looking at their electronegativities:

The electronegativity of the halogens decreases going down the group

 The electronegativity of an atom refers to how strongly it attracts electrons towards itself in a
covalent bond
 The decrease in electronegativity is linked to the size of the halogens
 Going down the group, the atomic radii of the elements increase which means that the outer
shells get further away from the nucleus
 An ‘incoming’ electron will therefore experience more shielding from the attraction of the
positive nuclear charge
 The halogens’ ability to accept an electron (their oxidising power) therefore decreases going
down the group
With increasing atomic size of the halogens (going down the group) their electronegativity,
and therefore oxidising power, decreases

 The reactivity of halogens is also shown by their displacement reactions with other halide ions
in solutions
 A more reactive halogen can displace a less reactive halogen from a halide solution of the less
reactive halogen

 Eg. The addition of chlorine water to a solution of bromine water:

Cl2(aq) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)

 The chlorine has displaced the bromine from solution as it is more reactive which can be
summarised in the following ionic equation:

Cl2(aq) + 2Br-(aq) → 2Cl-(aq) + Br2(aq)

Group 17: Reaction with Hydrogen

 Halogens react with hydrogen gas to form hydrogen halides

 Due to the decrease in reactivity of the halogens going down the group, the reactions between
halogen and hydrogen gas become less vigorous
 The table below shows a summary of the reaction between halogen and hydrogen gas

Reaction between halogen & hydrogen gas

Thermal Stability of the Hydrogen Halides

 Thermal stability refers to how well a substance can resist breaking down when heated
o A substance that is thermally stable will break down at high temperatures
 The hydrogen halides formed from the reaction of halogen and hydrogen gas decrease
in thermal stability going down the group
 The decrease in thermal stability can be explained by looking at the bond energies of the
hydrogen-halogen bond
o Going down the group, the atomic radius of the halogens increases
o The overlap of its outer shell with a hydrogen atom therefore gives a longer bond length
o The longer the bond, the weaker it is, and the less energy required to break it
 As the bonds get weaker, the hydrogen halogens become less stable to heat going down the
group
 The thermal stability of the hydrogen halide decreases going down the group as their bonds
become weaker due to the increased atomic radius of the halogens

Halide Ions: Reducing Agents

 Halide ions can also act as reducing agents and donate electrons to another atom
 The halide ions themselves get oxidised and lose electrons
 The reducing power of the halide ions increases going down the group
 This trend can be explained by looking at the ionic radii of the halides’ ions

The diagram shows that going down the group the ionic radii of the halogens increases

 Going down the group, the halide ions become larger

 The outermost electrons get further away from the nucleus
 The outermost electrons also experience more shielding by inner electrons
 As a result of this, the outermost electrons are held less tightly to the positively charged nucleus
 Therefore, the halide ions lose electrons more easily going down the group and their reducing
power increases
 The reducing power of the halide ions increases going down the group

Exam Tip
The ionic radius is a measure of the size of an atom’s ion.
Reactions of Halide Ions

Silver ions & ammonia

 Halide ions can be identified in an unknown solution by dissolving the solution in nitric acid and
then adding a silver nitrate solution followed by ammonia solution
 The halide ions will react with the silver nitrate solution as follows:

AgNO3(aq) + X-(aq) → AgX(s) + NO3-(aq)

(general equation)

Ag+(aq) + X-(aq) → AgX(s)

(ionic equation)


o X- is the halide ion in both equations
 If the unknown solution contains halide ions, then a precipitate of the silver halide will be
formed (AgX)

A silver halide precipitate is formed upon addition of silver nitrate solution to halide ion
solution

 Dilute followed by concentrated ammonia is added to the silver halide solution to identify the
halide ion
 If the precipitate dissolves in dilute ammonia the unknown halide is chloride
 If the precipitate does not dissolve in dilute but in concentrated ammonia the unknown halide
is bromide
 If the precipitate does not dissolve in dilute nor concentrated ammonia the unknown halide is
iodide

Silver chloride and silver bromide precipitates dissolve on addition of ammonia solution
whereas silver iodide doesn’t
Reaction of halide ions with silver nitrate & ammonia solutions table

Concentrated sulfuric acid

 Chloride, bromide and iodide ions react with concentrated sulfuric acid to produce toxic gases
 These reactions should therefore be carried out in a fume cupboard
 The general reaction of the halide ions with concentrated sulfuric acid is:

H2SO4(l) + X-(aq) → HX(g) + HSO4-(aq)

(general equation)

Where X- is the halide ion

Reaction of chloride ions with concentrated sulfuric Acid

 Concentrated sulfuric acid is dropwise added to sodium chloride crystals to produce hydrogen
chloride gas

Apparatus set up for the reaction of sodium chloride with concentrated sulfuric acid

 The reaction that takes place is:

H2SO4(l) + NaCl(s) → HCl(g) + NaHSO4(s)

 The HCl gas produces is seen as white fumes

Reaction of bromide ions with concentrated sulfuric acid

 The thermal stability of the hydrogen halides decreases down the group
 The reaction of sodium bromide and concentrated sulfuric acid is:

H2SO4(l) + NaBr(s) → HBr(g) + NaHSO4(s)

 The concentrated sulfuric acid oxidises HBr which decomposes into bromine and hydrogen
gas and sulfuric acid itself is reduced to sulfur dioxide gas:

2HBr(g) + H2SO4(l) → Br2(g) + SO2(g) + 2H2O(l)

 The bromine is seen as a reddish-brown gas

Reaction of iodide ions with concentrated sulfuric acid

 The reaction of sodium iodide and concentrated sulfuric acid is:

H2SO4 (l) + NaI (s) → HI (g) + NaHSO4 (s)

 Hydrogen iodide decomposes the easiest

 Sulfuric acid oxidises the hydrogen iodide to several extents:
 The concentrated sulfuric acid oxidises HI and is itself reduced to sulfur dioxide gas:

2HI (g) + H2SO4 (l) → I2 (g) + SO2 (g) + 2H2O (l)

 Iodine is seen as a violet/purple vapour

 The concentrated sulfuric acid oxidises HI and is itself reduced to sulfur:

6HI (g) + H2SO4 (l) → 3I2 (g) + S (s) + 4H2O (l)

 Sulfur is seen as a yellow solid

 The concentrated sulfuric acid oxidises HI and is itself reduced to hydrogen sulfide:

8HI (g) + H2SO4 (l) → 4I2 (g) + H2S (s) + 4H2O (l)

 Hydrogen sulfide has a strong smell of bad eggs

Halide ion reactions with concentrated sulfuric acid table
Reaction of Chlorine

 A disproportionation reaction is a reaction in which the same species is both oxidised and
reduced
 The reaction of chlorine with dilute alkali is an example of a disproportionation reaction
 In these reactions, the chlorine gets oxidised and reduced at the same time
 Different reactions take place at different temperatures of the dilute alkali

Chlorine in cold alkali (15 oC)

 The reaction that takes place is:

 The ionic equation is:

 The ionic equation shows that the chlorine gets both oxidised and reduced
 Chlorine gets oxidised as there is an increase in ox. no. from 0 to +1 in ClO-(aq)
o The half-equation for the oxidation reaction is:

 Chlorine gets reduced as there is a decrease in ox. no. from 0 to -1 in Cl-

(aq)
Chlorine in hot alkali (70 oC)

 The reaction that takes place is:

 The ionic equation is:

 The ionic equation shows that the chlorine gets both oxidised and
reduced
 Chlorine gets oxidised as there is an increase in ox. no. from 0 to +5 in
ClO3-(aq)

Chlorine gets reduced as there is a decrease in ox. no. from 0 to -1 in Cl-(aq)

  Chlorine can be used to clean water and make it drinkable
 The reaction of chlorine in water is a disproportionation reaction in
which the chlorine gets both oxidised and reduced

The disproportionation reaction of chlorine with water in which chlorine

gets reduced to HCl and oxidised to HClO

 Chloric(I) acid (HClO) sterilises water by killing bacteria

 Chloric acid can further dissociate in water to form ClO-(aq):

HClO(aq) → H+(aq) + ClO-(aq)

 ClO-(aq) also acts as a sterilising agent cleaning the water

Nitrogen and Sulfur
 Nitrogen is a diatomic molecule and the main unreactive gas in air
 78% of air is nitrogen gas
 The lack of reactivity of nitrogen gas can be explained by looking at its intramolecular bonds
 Intramolecular bonds are the bonds within a molecule

Bonding in nitrogen

 The electron configuration of a nitrogen atom is 1s2 2s2 2p3

 To achieve a full outer shell of electrons, it needs to gain three electrons
 Nitrogen atoms therefore form a triple covalent bond between two nitrogen atoms in which
they share three electrons with each other

 The bond enthalpy of the nitrogen triple bond is 1000 kJ mol-1

 This means that 1000 kJ of energy is needed to break one mole of N2 triple bond
 As it is so difficult to break the nitrogen triple bond, nitrogen and oxygen gas in air will not react
with each other
 Only under extreme conditions will nitrogen gas react (eg. during a thunderstorm)

Polarity of nitrogen

 The electrons in a nitrogen molecule are shared equally between the two nitrogen atoms
 Therefore, nitrogen molecules are nonpolar molecules

Since the electronegativity of the two nitrogen atoms is the same, the will pull the electrons
towards them equally so overall the molecule is nonpolar
  Due to the lack of polarity, nitrogen gas is not attracted to or likely to react with other
molecules the way polar molecules would

Nitrogen is very unreactive due to the lack of polarity and strength of its triple bond.

Properties of Ammonia

 Ammonia is a compound of nitrogen and will turn damp red litmus paper blue as it is
an alkaline gas
 Ammonia is made on a large scale in industry using the Haber process:

N2(g) + 3H2(g) ⇌ 2NH3(g)

Basicity of ammonia

 Ammonia can act as a Brønsted–Lowry base by accepting a proton (H+) using the lone pair of
electrons on the nitrogen atom to form an ammonium ion:

NH3(aq) + H+(aq) → NH4+(aq)

 In an aqueous solution of ammonia, an equilibrium mixture is established

NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)

 Since the position of the equilibrium lies well over to the left the ammonia solution is
only weakly alkaline
 There is a higher concentration of ammonia molecules than hydroxide ions in solution
 Ammonia is therefore a weak base

Structure & formation of ammonium ion

 The ammonium ion is formed by an acid-base reaction of ammonia with water:

NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)

 The nitrogen in ammonia is covalently bonded to three hydrogen atoms and has one lone pair of
electrons causing the ammonia molecule to have a pyramidal shape

Ammonia has a pyramidal shape due to its lone pair of electrons

 The nitrogen atom in ammonia uses its lone pair of electrons to form a dative bond with a
proton to form the ammonium ion
 The ammonium ion has a tetrahedral shape in which all bonds have the same length
Preparation of ammonia gas from an ammonium salt

 Ammonia gas can be prepared from an ammonium salt and a base in an acid-base reaction:

 Ammonium chloride (NH4Cl) and calcium hydroxide (Ca(OH)2) are mixed together and then
heated
 NH4+ acts as an acid (proton donor) and OH- acts as a base (proton acceptor)
 This acid-base reaction can be used to test if an unknown solution contains ammonium ions
 If the unknown solution does contain ammonium ions, it will react with calcium hydroxide to
form ammonia gas
 This ammonia gas will turn damp red litmus paper blue

Natural occurrence of nitrogen oxides

 Due to its lack of reactivity, only under extreme conditions will nitrogen react with oxygen to
form gaseous nitrogen oxides
 An example of these extreme conditions is lightning which can trigger the formation of
nitrogen(II) and nitrogen(IV) oxides (NO and NO2 respectively)
 The chemical equations for these reactions are:

N2(g) + O2(g) → 2NO(g)

N2(g) + 2O2(g) → 2NO2(g)

Man-made occurrence of nitrogen oxides

 In the engine of a car, a mixture of air and fuel is compressed and ignited by a spark
 Air consists of 78% of nitrogen and 21% of oxygen
 Under the high pressure and temperature inside a car engine, nitrogen can react with oxygen to
form nitrogen oxides
 These nitrogen oxides are released into the atmosphere through the car’s exhaust fumes

Catalytic removal of nitrogen oxides

 The nitrogen oxides released through cars’ exhaust fumes pollute the atmosphere
 Many car exhaust systems are therefore fitted with catalytic converters to reduce
the pollutants from motor vehicles
 The nitrogen oxides are reduced on the surface of the hot catalyst (e.g. platinum) to form the
unreactive and harmless nitrogen gas which is then released from the vehicle’s exhaust pipe into
the atmosphere
  The chemical reaction for the reduction of nitrogen oxide to nitrogen gas by the catalyst is as
follows:

2CO(g) + 2NO(g) → 2CO2(g) + N2(g)

A catalytic converter helps reduce the pollutants from motor vehicles

Oxides of Nitrogen & Photochemical Smog

 Nitrogen oxides are examples of primary pollutants because they are given off directly into the
air from the source of pollution
 Examples of pollution sources are car exhausts and power plants
 Nitrogen oxides are extra dangerous as they can react with substances in the air to
make secondary pollutants
 These are pollutants that are not given off directly into the air from human activity
 Exhaust fumes contain another primary pollutant called volatile organic compound (VOCs)
 These are unburnt hydrocarbons from fuel and their oxidised products
 VOCs react with nitrogen oxides in air to form peroxyacetyl nitrate (PAN, CH3CO3NO2)
 Sunlight provides the energy needed to start off the reactions of VOCs and nitrogen oxides in
air, so they are also called photochemical reactions
 PAN is one of the harmful pollutants found in photochemical smog
 ‘Smog’ is derived from ‘smoke’ and ‘fog’
 PAN affects the lungs and eyes and in high concentrations plant-life
Primary & secondary pollutant types & their pollution source table

The diagram shows the formation of PAN from the photochemical reaction between VOCs
and nitrogen oxide

Oxides of Nitrogen & Acid Rain

Formation of acid rain by nitrogen oxides

 As mentioned earlier, lightning strikes trigger the formation of nitrogen(II) and nitrogen (IV)
oxides in air:

2NO(g) + O2(g) ⇌ 2NO2(g)

 The air also contains oxygen and tiny droplets of water that make up clouds
 The nitrogen(IV) oxide (NO2) dissolves and reacts in water with oxygen as follows:

NO2(aq) + H2O(l) + 1½O2(g) → 2HNO3(aq)

 When the clouds rise, the temperature decreases, and the droplets get larger
 When the droplet containing dilute nitric acid are heavy enough, they will fall down as acid rain
 The diagram shows the formation of acid rain by the oxidation of nitrogen(IV) oxide

Nitrogen oxide as a catalyst

 Acid rain also contains dilute sulfuric acid (H2SO4)

 Sulfur(IV) oxide (SO2) is another pollutant found in the atmosphere
 When SO2 is oxidised, it forms SO3 which reacts with rainwater to form dilute sulfuric acid as
follows:

SO3(g) + H2O(l) → H2SO4(aq)

 Nitrogen oxides can directly cause acid rain but can also act as catalysts in the formation of acid
rain
 NO2 catalyses the oxidation of SO2 to SO3:
 NO2(g) + SO2(g) → SO3(g) + NO(g)


o The formed NO gets oxidised to regenerate NO2:

NO(g) + ½ O2(g) → NO2(g)

 The regenerated NO2 molecule can get again oxidise another SO2 molecule to SO3 which will
react with rainwater to form H2SO4

The formation of dilute sulfuric acid is catalysed by the nitrogen oxides