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Trends in Properties of Oxides

Key Concepts

  • Elements react with oxygen to produce oxides of the element.
    element + oxygen → element oxide

  • Group 18 elements (Noble gases) do not form oxides.

  • Trends in the oxides across period 3 of the Periodic Table from left to right:
    (i) ionic to covalent bonding

    (ii) ionic lattice to covalent network to covalent molecules

    (iii) basic to amphoteric to acidic

  • Oxides of metals are generally bases.

  • Oxides of non-metals are generally acids.

  • Oxides of some elements are amphoteric:
    zinc oxide, ZnO, is amphoteric
    aluminium oxide, Al2O3, is amphoteric

  • Some oxides of some elements are neutral:
    water, H2O, is neutral
    carbon monoxide, CO, is neutral
    nitric oxide, NO, is neutral
    nitrous oxide, N2O, is neutral

Producing Oxides

Period 3 elements are the ones most often used to describe the trends in the properties of oxides of elements in the Periodic Table. The relevant section of the Periodic Table is shown below:

  Group
1
Group
2
Groups
3 - 12
Group
13
Group
14
Group
15
Group
16
Group
17
Group
18
Period
3
11
Na
sodium
22.99
12
Mg
magnesium
24.31
  13
Al
aluminium
26.98
14
Si
silicon
28.09
15
P
phosphorus
30.97
16
S
sulfur
32.07
17
Cl
chlorine
35.45
18
Ar
argon
39.95

In order to produce an oxide, we need to react our Period 3 element with oxygen. The most common way to do this is to burn the element in air in a combustion reaction. The element can then react with the oxygen in the air to produce the oxide of the element:

element + oxygen → element oxide

The oxides of all Period 3 elements can be made this way, except:

  • oxides of argon: argon is a Noble Gas (Group 18) so it does not readily form compounds.

  • oxides of chlorine: oxides of chlorine produced in this way are highly unstable.

The chemical equation for the reaction of each Period 3 element with oxygen gas, O2(g), is given below:

Classification Group Period 3
Element
Reaction with Oxygen
metals 1 sodium (Na) (1) 4Na + O2 → 2Na2O
2 magnesium (Mg) 2Mg + O2 → 2MgO
13 aluminium (Al) 2Al + 3O2 → 2Al2O3
semi-metal
(metalloid)
14 silicon (Si) Si + O2 → SiO2
non-metals 15 phosphorus (P) 4P + 3O2 → 2P2O3

4P + 5O2 → 2P2O5

(2) 4P + 3O2 → P4O6

(3) 4P + 5O2 → P4O10

16 sulfur (S) (4) S + O2 → SO2
17 chlorine (Cl) (5) Stable oxides are not formed.
18 argon (Ar) no reaction

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Bonding in the Oxides

Bonding in the oxides of elements may be either ionic or covalent.
The type of bond can be determined by examining the difference in electronegativities between oxygen and the element.
Electronegativity refers to the tendency of a bonded atom to attract electrons to itself.
The electronegativities of each Period 3 element (using the Pauling method) is shown below under the symbol for the element:

  Group
1
Group
2
Groups
3 - 12
Group
13
Group
14
Group
15
Group
16
Group
17
Group
18
Element: Na Mg   Al Si P S Cl Ar
Electronegativity: 0.9 1.2   1.5 1.8 2.1 2.5 3.0 -
Trend low high  

Note that argon as a Noble gas (Group 18) element does not form compounds therefore can not be given a value for electronegativity.
Metals tend to have low electronegativities while non-metals tend to have higher electronegativity.

Using the Pauling method, oxygen has an electronegativity of 3.5 and is the second most electronegative element (the most electronegative is fluorine).
Let us calculate the difference in electronegativities between each element and oxygen and compare the results in order to determine the likely character (ionic or covalent) of the bond between the element and oxygen:

  Group
1
Group
2
Group
3 - 12
Group
13
Group
14
Group
15
Group
16
Group
17
Group
18
Period 3
bonds
Na-O Mg-O   Al-O Si-O P-O S-O Cl-O Ar
-
difference in
electronegativity
3.5 - 0.9
= 2.6
3.5 - 1.2
= 2.3
  3.5 - 1.5
= 2.0
3.5 - 1.8
= 1.7
3.5 - 2.1
= 1.4
3.5 - 2.5
= 1.0
3.5 - 3.0
= 0.5
Ar
-
Trend large difference small difference  

Clearly, oxygen is much better at attracting electrons to itself than the metals (Na, Mg and Al) so the difference in electronegativities between the metals and oxygen is large. So the oxygen atom pulls electrons away from the metal atom resulting in the formation of two ions:

  • a negatively charged oxygen ion (O2-, oxide ion) formed when the oxygen atom pulled 2 electrons of the metal atom(s)

  • a positively charged metal ion (Na+, Mg2+, Al3+) formed when an electron(s) were removed by the oxygen atom
Electrostatic attraction then acts to keep the negatively charged ions (O2- anions) and positively charged ions (metal cations) in an ordered 3-dimensional array, or lattice, of ions in the solid oxide.
The metal oxides will be ionic solids, ions held together by electrostatic attraction known as ionic bonds.
The physical properties of these ionic metal oxides include:

  • Solids that are hard and brittle.

  • High melting point and high boiling point.

  • Solids that do not conduct electricity.

  • Molten metal oxides (liquids) that do conduct electricity.

  • If dissolved in water, metal oxides do conduct electricity.

The other Period 3 elements, Si, P, S and Cl, are more like oxygen in electronegativity, that is, they are also quite good at attracting electrons to themselves so electrons are not pulled off one atom and given to the other, instead, the electrons making up the bond between these atoms and oxygen atoms are shared between the two atoms. This type of bond is known as covalent bond.

Silicon, the semi-metal (metalloid) in period 3, forms an oxide, SiO2, which is a 3-dimensional covalent network called silica as is similar in structure to that of diamond. Silica, SiO2, is relatively hard, has a high melting and boiling point, so it is a solid at room temperature and pressure, and does not conduct electricity.

The oxides of phosphorus (P4O6, P4O10), sulfur (SO2, SO), and chlorine (Cl2O, Cl2O7) are all small, discrete covalent molecules.
The physical properties of these covalent molecules:

  • low melting point and low boiling point

  • Solids do not conduct electricity.

  • Molten (liquid) oxides do not conduct electricity.

Therefore, we see a trend in the type of bond formed between a Period 3 element oxygen, and hence, a trend in the type of oxides formed, their structure and their properties:

  Group
1
Group
2
Group
3 - 12
Group
13
Group
14
Group
15
Group
16
Group
17
Group
18
Period 3
oxides
Na2O MgO   Al2O3 SiO2 P4O6
P4O10
SO2
SO3
Cl2O
Cl2O7
no
oxide
Trend in bond character ionic covalent  
Trend in structure:
(25oC, 100 kPa)
3-dimensional ionic lattice → 3-dimensional
covalent network
→ discrete covalent molecules  

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Reaction of the Oxides with Water

Let us now consider what happens when the oxides of each Period 3 element is placed in water, H2O(l).

First, we expect ionic solids (Period 3 metal oxides) to dissolve in water so the solution will be composed of aqueous metal ions and oxide ions:

General word equation
for solvation (hydration)
solid metal oxide H2O
metal cations + oxide ions
sodium oxide Na2O H2O
2Na+(aq) + O2-(aq)
magnesium oxide 2MgO H2O
2Mg2+(aq) + O2-(aq)
aluminium oxide no reaction, strength of ionic bonds
holding lattice together is too strong
(see heat of solution)

Na+ and Mg2+ do not react with water molecules, but the oxide ion, O2-, does:

oxide ion + water hydroxide ions
O2- + H2O(l) 2OH-(aq)

So the overall equations for the reaction between Na2O(s) and MgO(s) with water are:
Na2O(s) + H2O(l) → 2Na+(aq) + 2OH-(aq)     or     Na2O(s) + H2O(l) → 2NaOH(aq)

2MgO(s) + H2O(l) → 2Mg2+(aq) + 2OH-(aq)     or     2MgO(s) + H2O(l) → 2Mg(OH)2(aq)


These metal hydroxides, sodium hydroxide (NaOH) and magnesium hydroxide (Mg(OH)2) are both bases.

Silica, SiO2, does not react with water. The covalent bonds holding the silicon and oxygen atoms together in the 3-dimensional lattice are just too strong to be broken apart by the water molecules.

The oxides of the non-metals are all small, discrete covalent molecules, and they all react with water molecules to form oxyacids (an acid in which oxygen is attached to the non-metal).

We can summarise the reactions of water with the oxides of Period 3 elements as:

Bond
Type
Group Period 3
Oxide
Reaction with water Nature of
aqueous
product
ionic
lattice
1 sodium oxide, Na2O Na2O + H2O → 2NaOH basic
2 magnesium oxide, MgO MgO + H2O → Mg(OH)2
13 alumina, Al2O3 (6) no reaction  
covalent
network
14 silica, SiO2 no reaction
covalent
molecules
15 phosphorus(III) oxide, P4O6
diphosphorus trioxide

phosphorus(V) oxide, P4O10
diphosphorus pentoxide

(7) P4O6 + 6H2O → 4H3PO3
 

(8) P4O10 + 6H2O → 4H3PO4
 

acidic
16 sulfur dioxide, SO2

sulfur trioxide, SO3

(9) SO2 + H2O → H2SO3

SO3 + H2O → H2SO4

17 chlorine(I) oxide, Cl2O
dichlorine monoxide

chlorine(VII) oxide, Cl2O7
dichlorine heptoxide

Cl2O + H2O → 2HOCl
 

Cl2O7 + H2O → 2HClO4
 

We expect solutions of the metal oxides in water to display the properties of bases including the ability to turn red litmus blue, while we expect aqueous solutions of the non-metal oxides to display the properties of acids, including the ability to turn blue litmus red.

You have probably performed this experiment in the lab.

Sulfur is burnt in air in a gas jar containing a piece of blue litmus paper. A fine mist of water is then sprayed into the gas jar and the blue litmus paper changes colour from blue to red indicating the presence of an acid.

Performing a similar experiment with magnesium instead of sulfur and blue litmus paper instead of red litmus paper, then, the red litmus paper changes colour from red to blue indicating the presence of a base.

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Reaction of the Oxides with Acid

We expect bases to react with acids such as hydrochloric acid.
We saw above, that the oxides of Group 1 and Group 2 metals (sodium oxide and magnesium oxide) produce basic aqueous solutions.
Sodium oxide (Na2O(s)) reacts with dilute hydrochloric acid (HCl(aq)) to produce a salt (NaCl(aq)) and water (H2O).
Magnesium oxide (MgO(s)) reacts with warm dilute hydrochloric acid to produce magnesium chloride (MgCl2(aq)) and water.
AND, aluminium oxide (alumina, Al2O3) can also react with hot dilute hydrochloric acid to produce aluminium chloride (AlCl3(aq)() and water.

general word equation metal
oxide
+ hydrochloric
acid
salt + water
sodium oxide Na2O + 2HCl(aq) 2NaCl(aq) + H2O(l)
magnesium oxide MgO + 2HCl(aq) MgCl2(aq) + H2O(l)
aluminium oxide Al2O3 + 6HCl(aq) heat
2AlCl3(aq) + 3H2O(l)

On this basis we would conclude that the oxides of sodium, magnesium and aluminium are all basic ...... but .......

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Reaction of the Oxides with Base

We expect acids to react with bases such as sodium hydroxide to produce a salt and water.

But, there is a bit of a surprise here, aluminium oxide, which we decided was a base above, also reacts with bases!
If hot, concentrated sodium hydroxide solution (NaOH(aq)) is added to aluminium oxide (Al2O3(s)), then complex ions are formed with the sodium such as the sodium tetrahydroxoaluminate ion:

Al2O3(s) + 2NaOH(aq) + 3H2O(l) → 2NaAl(OH)4(aq)

In this reaction, Al2O3 is acting as an acid!
Aluminium oxide, Al2O3, is said to be amphoteric, it can act as either an acid or as a base.

Silicon dioxide, the oxide of a semi-metal or metalloid, is only very weakly acidic, it will react with hot concentrated hydrochloric acid to produce a sodium silicate and water:

SiO2(s) + 2NaOH(aq) → Na2SiO3 + H2O(l)

The Period 3 non-metal oxides react with bases, so they are all acidic.
Using aqueous sodium hydroxide, NaOH(aq), as the base, we can see that the oxides of phosphorus react with water to produce acids, and these acids can then react with the sodium hydroxide in a neutralisation reaction:

H3PO3 + 3NaOH(aq) → Na3PO3(aq) + 3H2O(l)

H3PO4 + 3NaOH(aq) → Na3PO4(aq) + 3H2O(l)


Sulfur dioxide, SO2(g) can be bubbled through aqueous solutions of sodium hydroxide to produce a salt (sodium sulfite) and water:
SO2(g) + 2NaOH(aq) → Na2SO3(aq) + H2O(l)

Sulfur trioxide, SO3(l), reacts violently with water to produce sulfuric acid, H2SO4(aq), so it is the sulfuric acid that will react with the sodium hydroxide base:
H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l)

Similarly, the oxides of chlorine readily form oxyacids with water, so it is the reaction between these acids and the sodium hydroxide base that are represented below:
HClO4(aq) + NaOH(aq) → NaClO4(aq) + H2O(l)
HOCl(aq) + NaOH(aq) → NaOCl(aq) + H2O(l)

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Trends in Acidity of Oxides of Elements in the Periodic Table

From the oxides of the Period 3 elements, we note the following general trend in acidity as we go across the period from left (Group 1) to right (Group 17):
basic oxides (Groups 1, 2) → amphoteric oxide (Al2O3) → acidic oxides (oxyacids)

The same trend, basic → amphoteric → acidic, can be seen in each period of the Periodic Table as shown below using colours to represent the acidic character of the oxides:

  •       no oxides

  •       neutral oxides

  •       basic oxides

  •       amphoteric acids

  •       acidic oxides

Period 1 1
H
hydrogen
1.008
  2
He
helium
4.003
  Group 1 Group 2 Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12 Group 13 Group 14 Group 15 Group 16 Group 17 Group 18
  basic → → → → → → → → → → → → → → amphoteric → → → → → acidic → none
Period 2 3
Li
lithium
6.941
4
Be
beryllium
9.012
  5
B
boron
10.81
6
C
(10)carbon
12.01
7
N
(11)nitrogen
14.01
8
O
oxygen
16.00
9
F
fluorine
19.00
10
Ne
neon
20.18
Period 3 11
Na
sodium
22.99
12
Mg
magnesium
24.31
  13
Al
aluminium
26.98
14
Si
silicon
28.09
15
P
phosphorus
30.97
16
S
sulfur
32.07
17
Cl
chlorine
35.45
18
Ar
argon
39.95
Period 4 19
K
potassium
39.10
20
Ca
calcium
40.08
21
Sc
scandium
44.96
22
Ti
titanium
47.87
23
V
vanadium
50.94
24
Cr
chromium
52.00
25
Mn
manganese
54.94
26
Fe
iron
55.85
27
Co
cobalt
58.93
28
Ni
nickel
58.69
29
Cu
copper
63.55
30
Zn
zinc
65.41
31
Ga
gallium
69.72
32
Ge
germanium
72.64
33
As
arsenic
74.92
34
Se
selenium
78.96
35
Br
bromine
79.90
36
Kr
krypton
83.80
Period 5 37
Rb
rubidium
85.47
38
Sr
strontium
87.62
39
Y
yttrium
88.91
40
Zr
zirconium
91.22
41
Nb
niobium
92.91
42
Mo
molybdenum
95.94
43
Tc
technetium
[97.91]
44
Ru
ruthenium
101.1
45
Rh
rhodium
102.9
46
Pd
palladium
106.4
47
Ag
silver
107.9
48
Cd
cadmium
112.4
49
In
indium
114.8
50
Sn
tin
118.7
51
Sb
antimony
121.8
52
Te
tellurium
127.6
53
I
iodine
126.9
54
Xe
xenon
131.3
Period 6 55
Cs
caesium
132.9
56
Ba
barium
137.3
57
La
lanthanum
138.9
72
Hf
hafnium
178.5
73
Ta
tantalum
180.9
74
W
tungsten
183.8
75
Re
rhenium
186.2
76
Os
osmium
190.2
77
Ir
iridium
192.2
78
Pt
platinum
195.1
79
Au
gold
197.0
80
Hg
mercury
200.6
81
Tl
thallium
204.4
82
Pb
lead
207.2
83
Bi
bismuth
209.0
84
Po
polonium
[209.0]
85
At
astatine
[210.0]
86
Rn
radon
[222.0]

Note that the following:

  • Group 18 (Noble gas) elements do not generally form oxides at room temperature and pressure.

  • Only metal oxides can be basic.

    If the metal oxide is soluble in water it produces an alkaline solution in the form of a metal hydroxide.

  • Soluble oxides of non-metals are acidic.

    The higher oxides of transition metals can also be acidic.

  • Water (H2O), and the relatively insoluble gases carbon monoxide (CO), nitric oxide (NO) and nitrous oxide (N2O) are neutral.

    The more soluble gases carbon dioxide (CO2) and nitrogen dioxide (NO2) are acidic.

  • Amphoteric oxides generally occur between the basic oxides on the left hand side of the Periodic Table and the acidic oxides on the right.

    Both aluminium oxide (Al2O3) and zinc oxide (ZnO) are amphoteric oxides.


What would you like to do now?

1 We will consider only the "normal" oxides. Apart from the normal oxide, Na2O, sodium also forms an ionic peroxide, Na2O2.
However, it should be noted that when Na combusts, Na2O will also react with O2 to produce Na2O2, so the major product of the combustion of sodium is Na2O2.

2 Limiting the supply of oxygen during combustion produces the lower oxide, P4O6 instead of P5O10

3 When phosphorus burns in excess oxygen, phosphoric oxide, P4O10, is produced.

4 When sulfur combusts, sulfur dioxide, SO2(g), is produced.
The oxidation of SO2 to SO3 by oxygen is spontaneous, but very slow:
SO2(g) + ½O2(g) → SO3(l)

5Several of the oxides of chlorine are prone to explode: ClO2, Cl2O, Cl2O3 and Cl2O7.
These appear to be shock sensitive rather than thermally sensitive.
Even so, ClO2 and Cl2O are both used commercially as bleaching agents, particular for bleaching paper and flour.
Commercially, ClO2 is prepared by the exothermic reaction between sodium chlorate in about 4 mol L-1 H2SO4 containing 0.05-0.25 mol L-1 chloride ion with sulfur dioxide:
2NaClO3 + SO2 + H2SO4 → 2ClO2 + 2NaHSO4

Cl2O can be prepared by treating freshly prepared yellow mercuric oxide with chlorine gas or with a solution of chlorine in carbon tetrachloride:
2Cl2 + 2HgO → HgCl2.HgO + Cl2O

6 There is only one formula for the oxide of aluminium, Al2O3, known as alumina, however, a number of polymorphs and hydrated species exist.
There are 2 forms of anhydrous Al2O3 known as α-Al2O3 and γ-Al2O3.
α-Al2O3 is very hard and resistant to hydration and attack by acids.
γ-Al2O3 is softer, readily takes up water and dissolves in acids.
There are several hydrated forms of alumina including AlO.OH and Al(OH)3, but these are prepared in alkaline solutions, not by reacting alumina with water.

7 This reaction occurs in cold water. If hot water is used, a range of products are formed such as PH3, phosphoric acid and element P.

8 This reaction occurs readily, making P4O10 a good drying agent, but it does produce a mixture of acids, depending on the quantity of water and other conditions.

9 There is no doubt that gaseous SO2 dissolves in water, but the acid, H2SO3 hasn't been isolated. Nevertheless, this equation is commonly accepted as a representation of the reaction.

10 Carbon dioxide, CO2, is an acidic oxide, but carbon monoxide, CO, is a neutral oxide.

11 Nitrogen dioxide, NO2, is an acidic oxide, but nitric oxide, NO, and nitrous oxide, N2O, are neutral oxides.

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