Acid-base Behavior of the Oxides Of The Period 3 elements
Metals react
with oxygen in the air to produce metal oxides, like magnesium oxide.
Non-metals
react with oxygen in the air to produce non-metal oxides.
Sulphur a non-metal burns in air to form sulphur dioxide: sulphur +
oxygen → sulphur dioxide
Here we will
discusse the reactions of the oxides of Period 3 elements (sodium to chlorine)
with water, and with acids or bases where relevant (as before, argon is omitted
because it does not form an oxide).
The trend in
acid-base behavior can be summarized as follows:
Acidity
increases from left to right, ranging from strongly basic oxides on the left to
strongly acidic ones on the right, with an amphoteric oxide (aluminum oxide) in
the middle. An amphoteric oxide is one which shows both acidic and basic
properties.
1.Sodium Oxide
Sodium oxide
is a simple strongly basic oxide. It is basic because it contains the oxide
ion, O2-, which is a very strong base with a high tendency to combine with
hydrogen ions.
Reaction
with water: Sodium oxide reacts exothermically with cold water to produce
sodium hydroxide solution. A concentrated solution of sodium oxide in water
will have pH 14. Na2O+H2O→2NaOH
Reaction
with acids: As a strong base, sodium oxide also reacts with acids. For example,
it reacts with dilute hydrochloric acid to produce sodium chloride solution.
2.Magnesium
oxide
Magnesium
oxide is another simple basic oxide, which also contains oxide ions. However,
it is not as strongly basic as sodium oxide because the oxide ions are not as
weakly-bound. In the sodium oxide, the solid is held together by attractions
between 1+ and 2- ions. In magnesium oxide, the attractions are between 2+ and
2- ions. Because of the higher charge on the metal, more energy is required to
break this association. Even considering other factors (such as the energy
released from ion-dipole interactions between the cations and water), the net
effect is that reactions involving magnesium oxide will always be less
exothermic than those of sodium oxide.
Reaction
with water: At first glance, magnesium oxide powder does not appear to react
with water. However, the pH of the resulting solution is about 9, indicating
that hydroxide ions have been produced. In fact, some magnesium hydroxide is
formed in the reaction, but as the species is almost insoluble, few hydroxide
ions actually dissolve. The reaction is as follow: MgO+H2O→Mg(OH)2
Reaction
with acids: Magnesium oxide reacts with acids as predicted for a simple metal
oxide. For example, it reacts with warm dilute hydrochloric acid to give
magnesium chloride solution. MgO+2HCl→MgCl2+H2O
3.Aluminum
Oxide
Sometime the properties of aluminum oxide can be confusing because it exists in a number
of different forms. One of those forms is very unreactive (known chemically as
alpha-Al2O3) and is produced at high temperatures. The following reactions
concern the more reactive forms of the molecule. Aluminium oxide is amphoteric.
It has reactions as both a base and an acid.
Reaction
with water: Aluminum oxide is insoluble
in water and does not react like sodium oxide and magnesium oxide. The oxide
ions are held too strongly in the solid lattice to react with the water.
Reaction
with acids: Aluminum oxide contains oxide ions, and thus reacts with acids in
the same way sodium or magnesium oxides do.
Aluminum oxide reacts with hot
dilute hydrochloric acid to give aluminum chloride solution. Al2O3+6HCl→2AlCl3+3H2O
This
reaction and others display the amphoteric nature of aluminum oxide.
Reaction
with bases: Aluminum oxide also displays acidic properties, as shown in its
reactions with bases such as sodium hydroxide. Various aluminates (compounds in
which the aluminum is a component in a negative ion) exist, which is possible
because aluminum can form covalent bonds with oxygen.
This is possible because
the electronegativity difference between aluminum and oxygen is small, unlike
the difference between sodium and oxygen, for example (electronegativity
increases across a period)
Aluminum
oxide reacts with hot, concentrated sodium hydroxide solution to produce a
colorless solution of sodium tetrahydroxoaluminate:
Al2O3+2NaOH+3H2O→2NaAl(OH)4
4.Silicon
dioxide (silicon(IV) oxide)
Silicon is
too similar in electronegativity to oxygen to form ionic bonds. Therefore,
because silicon dioxide does not contain oxide ions, it has no basic
properties. In fact, it is very weakly acidic, reacting with strong bases.
Reaction
with water: Silicon dioxide does not react with water, due to the thermodynamic
difficulty of breaking up its network covalent structure.
Reaction
with bases: Silicon dioxide reacts with hot, concentrated sodium hydroxide
solution, forming a colorless solution of sodium silicate:
SiO2+2NaOH→Na2SiO3+H2O
In another
example of acidic silicon dioxide reacting with a base, the Blast Furnace
extraction of iron, calcium oxide from limestone reacts with silicon dioxide to
produce a liquid slag, calcium silicate:
SiO2+CaO→CaSiO3(1.8)
Phosphorus
Oxides
Two
phosphorus oxides, phosphorus(III) oxide, P4O6, and phosphorus(V) oxide, P4O10,
are treated in this article.
Phosphorus(III)
oxide: Phosphorus(III) oxide reacts with cold water to produce a solution of
the weak acid, H3PO3—known as
phosphorous acid, orthophosphorous acid or phosphonic acid:
P4O6+6H2O→4H3PO3
Below is the image of the fully-protonated acid structure :
The protons
remain associated until water is added; even then, because phosphorous acid is
a weak acid, few acid molecules are deprotonated. Phosphorous acid has a pKa of
2.00, which is more acidic than common organic acids like ethanoic acid
(pKa =
4.76).
Phosphorus(III)
oxide is unlikely to be reacted directly with a base. In phosphorous acid, the
two hydrogen atoms in the -OH groups are acidic, but the third hydrogen atom is
not. Therefore, there are two possible reactions with a base like sodium
hydroxide, depending on the amount of base added:
NaOH+H3PO3→NaH2PO3+H2O
2NaOH+H3PO3→Na2HPO3+2H2O
In the first
reaction, only one of the protons reacts with the hydroxide ions from the base.
In the second case (using twice as much sodium hydroxide), both protons react.
If instead
phosphorus(III) oxide is reacted directly with sodium hydroxide solution, the
same salts are possible:
4NaOH+P4O6+2H2O→4NaH2PO3
9NaOH+P4O6→4Na2HPO3+2H2O
Phosphorus(V)
oxide: Phosphorus(V) oxide reacts violently with water to give a solution
containing a mixture of acids, the nature of which depends on the reaction
conditions. Only one acid is commonly considered, phosphoric(V) acid, H3PO4
(also known as phosphoric acid or as orthophosphoric acid).
P4O10+6H2O→4H3PO4
The fully protonated acid has the following structure:
Phosphoric(V)
acid is another weak acid with a pKa of 2.15, marginally weaker than
phosphorous acid. Solutions of each of these acids with concentrations around 1
mol dm-3 have a pH of about 1.
Phosphoric
(V) oxide is also unlikely to be reacted directly with a base, but the
hypothetical reactions are considered. In its acid form, molecule has three
acidic -OH groups, which can cause a three-stage reaction with sodium
hydroxide:
NaOH+H3PO4→NaH2PO4+H2O
2NaOH+H3PO4→Na2HPO4+2H2O
3NaOH+H3PO4→Na3PO4+3H2O
Similar to
phosphorus (III) oxide, if phosphorus(V) oxide reacts directly with sodium
hydroxide solution, the same possible salt as in the third step (and only this
salt) is formed:
12NaOH+P4O10→4Na3PO4+6H2O
5.Sulfur Oxides
Two oxides
are also considered: sulfur dioxide, SO2, and sulfur trioxide, SO3.
Sulfur
dioxide: Sulfur dioxide is fairly soluble in water, reacting to give a solution
of sulfurous acid (also known as sulfuric(IV)acid), H2SO3, as you will see in the
reaction below. This species only exists in solution, and any attempt to
isolate it gives off sulfur dioxide.
SO2+H2O→H2SO3
The
protonated acid has the following structure:
Sulfurous
acid is also a relatively weak acid, with a pKa of around 1.8, but slightly
stronger than the two phosphorus-containing acids above. A reasonably
concentrated solution of sulfurous acid has a pH of about 1.
Sulfur
dioxide also reacts directly with bases such as sodium hydroxide solution.
Bubbling sulfur dioxide through sodium hydroxide solution first forms sodium
sulfite solution, followed by sodium hydrogen sulfite solution if the sulfur dioxide is in excess.
SO2+2NaOH→Na2SO3+H2O
Na2SO3+H2O→2NaHSO3
Another
important reaction of sulfur dioxide is with the base calcium oxide to form
calcium sulfite (also known as calcium sulfate(IV)). This is of the important
methods of removing sulfur dioxide from flue gases in power stations.
CaO+SO2→CaSO3
Sulfur
trioxide: Sulfur trioxide reacts violently with water to produce a fog of
concentrated sulfuric acid droplets.
SO3+H2O→H2SO4
Below is the protonated structure:
Sulfuric
acid is a strong acid, and solutions will typically have a pH around 0. The
acid reacts with water to give a hydronium ion (a hydrogen ion in solution) and
a hydrogen sulfate ion. This reaction runs essentially to completion:
H2SO4(aq)+H2O(l)→H3P++HSO−4(aq)
The second
proton is more difficult to remove. In fact, the hydrogen sulfate ion is a
relatively weak acid, similar in strength to the acids discussed above.
This
reaction is more appropriately described as an equilibrium:
HSO−4(aq)+H2O⇌H3O+(aq)+SO2−4(aq)
It is useful
if you understand the reason that sulfuric acid is a stronger acid than
sulfurous acid. You can apply the same reasoning to other acids that you will come across on this article.
Sulfuric
acid is stronger than sulfurous acid because when a hydrogen ion is lost from
one of the -OH groups on sulfuric acid, the negative charge left on the oxygen
is spread out (delocalized) over the ion by interacting with the doubly-bonded
oxygen atoms.
It follows that more double bonded oxygen atoms in the ion make
more delocalization possible; more delocalization leads to greater stability,
making the ion less likely to recombine with a hydrogen ion and revert to the non-ionized
acid.
Sulfurous
acid only has one double bonded oxygen, whereas sulfuric acid has two; the
extra double bond provides much more effective delocalization, a much more
stable ion, and a stronger acid. Sulfuric acid displays all the reactions characteristic
of a strong acid. For example, a reaction with sodium hydroxide forms sodium
sulfate; in this reaction, both of the acidic protons react with hydroxide ions
as shown:
2NaOH+H2SO4→Na2SO4+2H2O
In
principle, sodium hydrogen sulfate can be formed by using half as much sodium
hydroxide; in this case, only one of the acidic hydrogen atoms is removed.
Sulfur
trioxide itself also reacts directly with bases such as calcium oxide, forming
calcium sulfate:
CaO+SO3→CaSO4
This
reaction is similar to the reaction with sulfur dioxide discussed above.
6.Chlorine
Oxides
Chlorine
forms several oxides, but only two (chlorine(VII) oxide, Cl2O7, and
chlorine(I)oxide, Cl2O) would be touched here. Chlorine(VII) oxide is also known
as dichlorine heptoxide, and chlorine(I) oxide as dichlorine monoxide.
Chlorine(VII)
oxide: Chlorine(VII) oxide is the highest oxide of chlorine—the chlorine atom
is in its maximum oxidation state of +7. It continues the trend of the highest
oxides of the Period 3 elements towards being stronger acids. Chlorine(VII)
oxide reacts with water to give the very strong acid, chloric(VII) acid, also
known as perchloric acid.
Cl2O7+H2O→2HClO4
As in
sulfuric acid, the pH of typical solutions of perchloric acid are around 0.
Neutral chloric(VII) acid has the following structure:
When the
chlorate(VII) ion (perchlorate ion) forms by loss of a proton (in a reaction
with water, for example), the charge is delocalized over every oxygen atom in
the ion. That makes the ion very stable, making chloric(VII) acid very strong.
Chloric(VII)
acid reacts with sodium hydroxide solution to form a solution of sodium
chlorate(VII):
NaOH+HClO4→NaClO4+H2O
Chlorine(VII)
oxide itself also reacts directly with sodium hydroxide solution to give the
same product:
2NaOH+Cl2O7→2NaClO4+H2O
Chlorine(I)
oxide: Chlorine(I) oxide is far less acidic than chlorine(VII) oxide. It reacts
with water to some extent to give chloric(I) acid, HOCl−HOCl−
also known as hypochlorous acid.
Cl2O+H2O⇌2HOCl
The
structure of chloric(I) acid is exactly as shown by its formula, HOCl. It has
no doubly-bonded oxygens, and no way of delocalizing the charge over the
negative ion formed by loss of the hydrogen. Therefore, the negative ion formed
not very stable, and readily reclaims its proton to revert to the acid.
Chloric(I) acid is very weak (pKa = 7.43) and reacts with sodium hydroxide
solution to give a solution of sodium chlorate(I) (sodium hypochlorite):
NaOH+HOCl→NaOCl+H2O
Chlorine(I)
oxide also reacts directly with sodium hydroxide to give the same product:
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