Untitled Lesson

CHEMICAL properties of period
2 & 3 oxides & hydroxides

CHEMICAL properties of
substances period 2 & 3
oxides & hydroxides:

•amphoteric character of alumina
•basic character of metal oxides and hydroxides
•ease of electrolysis.

Industrial processes for school & colleges

http://www.rsc.org/Education/Teachers/Resources/Alchemy/index2.htm

Reactions of Period 3 Elements with oxygen

Starter activity:

• Copy and complete the following reactions; identify the
different type of oxides – basic, acidic & amphoteric.

1. Sodium: Na + O2 →

2. Magnesium:

Mg + O2 →

3. Aluminium:

Al + 3O2 →

4. Silicon:

Si + O2 →

5. Phosphorus:

P +5O2 →

6. Sulfur:

S + O2 →

Reactions of Period 3 Elements with oxygen

Starter activity:
• Copy and complete the following reactions; identify the
different type of oxides – basic, acidic & amphoteric.

1. Sodium: 4Na + O2 → 2Na2O (basic).

2. Magnesium: 2Mg + O2 → 2MgO (basic).

3. Aluminium: 4Al + 3O2 → 2Al2O3 (amphoteric).

4. Silicon: Si + O2 → SiO2 (acidic).

5. Phosphorus: 4P +5O2 →2P2O5 (acidic)

4P +5O2 → P4O10 (acidic).

6. Sulfur: S + O2 → SO2 & 2S + 3O2 → 2SO3 (acidic).

Reactions of Period 3 Elements with oxygen

• Most metal oxides act as bases because of the oxide ion,

O2-

•

Reaction with water to form
hydroxide ions
O2-+ H2O →

2OH-

Neutralisation of an acid

O2-+ 2H+→ H2O

Aluminium oxide

–

amphoteric properties

• Aluminium oxide (Al2O3) also called alumina is described

as amphoteric oxide.

Why?

• Because the oxide can behave as an acid or base and can

therefore react with both acids and bases, neutralising the
other and producing a salt.

• Reaction with:

Acid:
Al2O3(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2O(l)
Base:
Al2O3(s) + 2NaOH(aq) → 2NaAl(OH)4(s)

• alumina has a strong ionic bond with very high melting

point. In a later section, we’ll be looking at how the metal
can be extracted from its ores.

Chemical properties of substances: (Period 3 OXIDES & HYDROXIDES)

Acidic, basic & amphoteric oxides across the PT

Aluminium oxide

–

Properties

Physical

Chemical

High melting and boiling point

does not dissolve in water

thermal conductivity

it is highly reactive to chlorine trifluoride and ethylene
oxide. Mixing aluminium oxide with either of these
chemicals causes a fire

electrical conductivity

Mechanical strength

Density

Compressive strength

Aluminium oxide

–

Uses

Uses of aluminium oxide
•Most of the aluminium oxide produced is used to form aluminium metal.

•Used in refractories: it has excellent wear resistance properties. It has high corrosion endurance and high temperature
stability, low thermal expansion and a favourable stiffness-to-weight ratio.

•Due to its excellent mechanical, chemical and thermal qualities, alumina stands out from many comparable materials by
delivering equal or better solutions for low-cost production and manufacturing.

•Medical industry: Due to aluminium oxide’s hardness, bio-inertness and chemical properties, it is a preferred material
for bearings in hip replacements, as prostheses, bionic implants, prosthetic eye substitutes, tissue reinforcements, dental
crowns, abutments, bridges, and other dental implants. It is also used in lab equipment and tools like crucibles, furnaces
and other labware

•Military and protective equipment: Aluminium oxide’s strength and lightweight qualities contribute to enhancing body
armours, like breastplates, as well as vehicle and aircraft armour, which is its biggest market. Aluminium oxide is also
used in synthetic-sapphire bulletproof windows and ballistics.

•Electrical and electronics industry: Its high melting and boiling points, in addition to its excellent thermal resistive
properties, make aluminium oxide desirable in the manufacture of high-temperature furnace insulations and electrical
insulators. Alumina films are also vital components in the microchip industry. Some of its other uses include spark plug
insulators, micro-electric substrates and insulating heatsinks.

•Industrial applications: Since alumina is chemically inert, it is utilised as a filler in plastics, bricks, and other heavy
clayware, like kilns. Due to its extreme strength and hardness, it is often used as an abrasive for sandpaper. It is also an
economical substitute for industrial diamonds

•Gem industry: Aluminium oxide is a valuable component in the formation of rubies and sapphire or precious gems. Its
crystalline form, corundum, is the base element for these precious gems.

Uses of Period 3 Oxides & Hydroxides

Substances

Uses

Calcium oxide
(CaO) - quicklime
•Reacts with water to form Ca(OH)2 (lime).
•Used by farmers to raise the pH of acidic soil (treatment of
acidic effluent in furnaces).

Magnesium oxide
(MgO)
•Used as desiccants when preserving books in the library

Sodium hydroxide
(NaOH)
•Used in some processes to make plastics and soaps.
•Used in food processing
•Used as drain cleaner or in oven cleaner

Magnesium
hydroxide
(Mg(OH)2

•Also known as ‘milk of magnesia’ and is used for treating
indigestion

Calcium hydroxide
(CaOH)2
•Used in the treatment of acidic effluent

Aluminium oxide
(Al2O3)
Refractory material in furnaces

Electrolysis

• What is electrolysis?
o Electrolysis is the chemical separation of ionic compound using a direct

electrical current (electricity).
o It is used to extract reactive metals from their ores.
o Ionic substances are decomposed (broken down) by an electrical charge being
passed through them during electrolysis
o The ions must be free to move for this to work, so the compound must either
be molten (reduced to liquid by form by heating) or in solution

• Electrolysis needs:
o a dc electrical supply
o a negative electrode, called a cathode
o a positive electrode, called an anode

• During electrolysis:
o positive ions, cations, move to the cathode
o negative ions, anions, move to the anode

o Sodium, chlorine, magnesium, and aluminium are four elements produced

commercially by electrolysis.

Electrolysis

o What is an anode?
o An anode is an electrode
through which the
conventional current enters
into an electrical device.
o It is the positively charged
electrode of an electrical
device
o Events happening at the anode:

o What is a cathode?
o a cathode, an electrode
through which conventional
current leaves an electrical
device.
o It is the negatively charged
electrode of an electrical
device

• Electrolysis set-up and components

•What is an electrolyte?
oIonic compounds conduct electricity when their ions are free to move.

oIons can move in the liquid state (after melting) or in aqueous solution (after dissolving in water).

oThe molten or dissolved substance is called the electrolyte.

oAn electrolyte is the chemical separation of ionic compound using a direct electrical current (electricity).

Electrolysis
• During electrolysis of
a molten ionic compound:

o positive metal ions are attracted to
the negative electrode (cathode),
where they gain electrons and form
metal atoms

o negative non-metal ions are
attracted to the positive electrode
(anode), where they lose electrons
and form non-metal atoms

• Non-metal atoms formed at the
positive electrode, the anode,
usually join together
by covalent bonding to form
simple molecules.

Example
•During electrolysis of molten lead(II) bromide, PbBr2(l):

•Pb2+ions gain electrons at the cathode and become Pb
atoms

•Br-ions lose electrons at the anode and become Br
atoms, which pair up to form Br2 molecules
•Overall, lead forms at the negative electrode and bromine
forms at the positive electrode.

Activity

Complete these tasks

1) Predict which products form at the anode and the
cathode during the electrolysis of molten calcium
chloride. Explain your answer.

2) Predict which products form at each electrode during
the electrolysis of molten aluminium oxide.

Activity_Ans

Complete these tasks

1)

Predict which products form at the anode and the cathode during
the electrolysis of molten calcium chloride. Explain your answer.
Answer:
Calcium will form at the cathode and chlorine will form at the anode.
Cathode: Ca2++ 2e-

Ca

Anode: 2Cl-

Cl2 + 2e-
This is because positive calcium ions are attracted to the negative electrode
(cathode), where they gain electrons to form calcium atoms.
At the same time, negative chloride ions are attracted to the positive electrode
(anode). They lose electrons to form chlorine atoms, which join by covalent
bonds to form chlorine molecules.

2) Predict which products form at each electrode during the electrolysis
of molten aluminium oxide.

Answer:
Aluminium will form at the negative electrode (cathode) and oxygen will form at
the positive electrode (anode).

Oxidation and reduction in electrolysis

• Oxidation and reduction can be described in terms
of electrons:

oxidation is the loss of electrons
reduction is the gain of electrons

• Reduction happens at the negative cathode because this
is where positive ions gain electrons.

• Oxidation happens at the positive anode because this is
where negative ions lose electrons.

• This tells us of the events happening at the anode and
cathode.
• A redox reaction takes place

• A half equation is used to model the reaction that
happens at an electrode during electrolysis.

• It shows what happens when ions gain or lose electrons.
• In half equations:
oelectrons are shown as e-
oAlways write +e- and never –e-
o+e written on left hand side of half equation for reduction

(cathode)

o+e- is written on the right hand side of half-equation for

oxidation (anode)

othe numbers of atoms of each element must be the same on
both sides
othe total charge on each side must be the same (usually zero)

Electrolysis

–

writing half equations

Electrolysis

–

writing half equations

Cathode reactions

• Positively charged ions gain electrons at the cathode. These
are half equations for some reactions at the cathode:

Na++ e-→ Na

Pb2++ 2e-→ Pb

Example

• Balance the half equation for the formation of aluminium
during electrolysis: Al3++ e-→ Al

• The balanced half equation is: Al3++ 3e-→ Al

(because three negatively charged electrons are needed to balance the
three positive charges on the aluminium ion).

Electrolysis

–

writing half equations

Anode reactions
• Negatively charged ions lose electrons at the anode.
These are half equations for some reactions at the anode:

2Cl-→ Cl2 + 2e-
2O2-→ O2 + 4e-