Periodic Table: Trends in the Properties of Oxides of Elements Chemistry Tutorial

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:
(i) zinc oxide, ZnO, is amphoteric
(ii) aluminium oxide, Al₂O₃, is amphoteric

Some oxides of some elements are neutral:
(i) water, H₂O, is neutral
(ii) carbon monoxide, CO, is neutral
(iii) nitric oxide, NO, is neutral
(iv) nitrous oxide, N₂O, is neutral

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Producing Oxides of Elements

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, O_2(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 of Period 3 Elements

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 in electronegativity	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 element 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 in difference in electronegativities	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, oxygen anion, (O^2-, oxide ion) formed when the oxygen atom pulled 2 electrons of the metal atom(s)

a positively charged metal ion, metal cation, (Na⁺, Mg²⁺, Al³⁺) formed when an electron(s) were removed by the oxygen atom

Electrostatic attraction then acts to keep the negatively charged ions (O^2- 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, SO₃), and chlorine (Cl₂O, Cl₂O₇) are all small, discrete covalent molecules.
The physical properties of these covalent molecules depend on the strength of the intermolecular forces acting between these discrete 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	SO₂	Cl₂O	no oxide
Trend in bond character	ionic		→	→	→	→	covalent
Trend in structure: (25^oC, 100 kPa)	3-dimensional ionic lattice →				3-dimensional covalent network	→ discrete covalent molecules

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Reaction of the Period 3 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 aqueous oxide ions:

General word equation for solvation (hydration)	solid metal oxide	H₂O →	metal cations	+	oxide ions
sodium oxide	Na₂O	H₂O →	2Na⁺_(aq)	+	O^2-_(aq)
magnesium oxide	2MgO	H₂O →	2Mg²⁺_(aq)	+	O^2-_(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, O^2-, does:

oxide ion	+	water	→	hydroxide ions
O^2-	+	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	basic
	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 the experiment shown on the right in the lab.

Sulfur is burnt in air in the 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 Period 3 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 (MgCl_2(aq)) and water.
AND, aluminium oxide (alumina, Al₂O₃) can also react with hot dilute hydrochloric acid to produce aluminium chloride (AlCl_3(aq)() and water.

general word equation	metal oxide	+	hydrochloric acid	→	salt (metal chloride)	+	water
sodium oxide	Na₂O	+	2HCl_(aq)	→	2NaCl_(aq)	+	H₂O_(l)
magnesium oxide	MgO	+	2HCl_(aq)	→	MgCl_2(aq)	+	H₂O_(l)
aluminium oxide	Al₂O₃	+	6HCl_(aq)	heat →	2AlCl_3(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 Period 3 Oxides with Base

We expect acids to react with bases such as sodium hydroxide to produce a salt and water (in a neutralisation reaction).

But, there is a bit of a surprise here.
Aluminium oxide, which we decided was a base in the section above, also reacts with bases!
If hot, concentrated sodium hydroxide solution (NaOH_(aq)) is added to aluminium oxide (Al₂O_3(s)), then complex ions are formed with the sodium such as the sodium tetrahydroxoaluminate ion:

Al₂O_3(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:

SiO_2(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) → Na₃PO_3(aq) + 3H₂O_(l)
H₃PO₄ + 3NaOH_(aq) → Na₃PO_4(aq) + 3H₂O_(l)

Sulfur dioxide, SO_2(g) can be bubbled through aqueous solutions of sodium hydroxide to produce a salt (sodium sulfite) and water:

SO_2(g) + 2NaOH_(aq) → Na₂SO_3(aq) + H₂O_(l)

Sulfur trioxide, SO_3(l), reacts violently with water to produce sulfuric acid, H₂SO_4(aq), so it is the sulfuric acid that will react with the sodium hydroxide base:

H₂SO_4(aq) + 2NaOH_(aq) → Na₂SO_4(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:

HClO_4(aq) + NaOH_(aq) → NaClO_4(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

¹⁰carbon
12.01

7
N

¹¹nitrogen
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.

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Footnotes

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, SO_2(g), is produced.
The oxidation of SO₂ to SO₃ by oxygen is spontaneous, but very slow:
SO_2(g) + ½O_2(g) → SO_3(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.