Period 3 elements and oxides

​​In a nutshell

Period $$3$$​ metals form metal oxides, these metal oxides can differ in structure and their properties. They can have ionic, giant and simple covalent structures and can behave as either acids, bases or exhibit amphoteric behaviour. 

Where are Period 3 elements? 

Period $$3$$​ elements are the third row of elements in the periodic table.

Sodium and magnesium reactivity

Sodium and magnesium are the first and second Period $$3$$ elements. Sodium is in Group $$1$$ and magnesium is in Group $$2$$. Sodium loses one electron $$(Na^+)$$ and magnesium loses two electrons $$(Mg^{2+})$$​. Sodium is more reactive than magnesium because it's easier to lose one electron than two. It takes more energy for magnesium to react (typically in the form of heat). 

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Below shows the reaction of sodium and magnesium with water:

​$$2Na(s)\ +\ 2H_2O(l)\ \rightarrow\ 2NaOH(aq)\ +\ H_2(g)$$​​

Sodium reacts violently with cold water. A molten ball is formed, fizzing and hydrogen gas is given off. Sodium hydroxide is formed and this forms a strongly alkaline solution ($$pH\ 12-14$$​).

​$$Mg(s)\ +\ 2H_2O(l)\ \rightarrow\ Mg(OH)_2(aq)\ +\ H_2(g)$$​​

Magnesium reacts very slowly with cold water. A thin layer of magnesium hydroxide, not very soluble in water especially cold, prevents the magnesium on the inside of the metal from reacting. This produces few hydroxide ions which in turn forms a weakly alkaline solution $$(pH\ 9-10)$$​.

Magnesium reacts much more vigorously with steam, this is hot water which has more energy. Magnesium oxide is formed and not magnesium hydroxide. 

​$$Mg(s)\ +\ H_2O(g)\ \rightarrow\ MgO(s)\ +\ H_2(g)$$​​

Period 3 elements and oxygen 

Period $$3$$ elements react readily with oxygen to form metal oxides. Typically Period $$3$$ elements are oxidised to their highest oxidation state (the same as their group number). Sulfur is an exception as its highest oxidation state is $$+4$$, a high temperature and catalyst are required to make $$SO_3$$, where sulfur has an oxidation number of $$+6$$.

The equation for the formation of metal oxides is generally:

​$$metal\ +\ oxygen \ \rightarrow\ metal\ oxide$$​​

Examples 

Oxides and melting points 

Metal oxides have differing bonding and structures, which leads to greatly different melting points. 

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​$$Na_2O$$, $$MgO$$ and $$Al_2O_3$$  have very high melting points because they form giant ionic lattices. These are held together by very strong forces of electrostatic attraction, a lot of energy is required to overcome this strong bonding. $$MgO$$ has a higher melting point than $$Na_2O$$ because the $$Mg$$ forms $$2+$$ ions which bonds more strongly to oxygen than the $$1+$$ ions of sodium. $$Al_2O_3$$ has a lower melting point of than $$MgO$$ since the difference in electronegativity between $$Al$$ and $$O$$ is smaller than $$Mg$$ and $$O$$. This means that the aluminium ions don't attract electrons from the oxygen as well as magnesium, this means that the bonding between aluminium and oxygen is partially covalent.

​$$SiO_2$$ has a lower melting point than $$MgO$$ and $$Al_2O_3$$, however a higher melting point than the remaining metal oxides. $$SiO_2$$​ forms a giant covalent structure, this requires a lot of every to overcome this bonding.

​$$P_4O_{10}$$ and $$SO_2$$ have relatively low melting points and they form simple covalent structures. These molecules are held together by weak intermolecular forces (van der Waals), these do not take a lot of energy to overcome.​

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Alkalinity and acidity of oxides

The ionic oxides of $$Na$$ and $$Mg$$ contains oxide ions $$(O^{2-})$$​. When oxide ions dissolve in water they accept protons to form hydroxide ions.

​$$O^{2-}(aq)\ +\ H^+(aq)\ \rightarrow\ \ OH^{-}(aq)$$ 

Both sodium and magnesium hydroxide are alkaline. However, sodium hydroxide is much more soluble in water and therefore can form much more alkaline solutions than magnesium hydroxide.

​$$Na_2O(s)\ +\ H_2O(l)\ \rightarrow\ 2NaOH(aq)\ :\ pH\ 12\ -\ 14$$ 

$$MgO(s)\ +\ H_2O(l)\ \rightarrow\ Mg(OH)_2(aq)\ :\ pH\ 9\ -\ 10$$​​

The simple covalent compounds oxides with the non-metal compounds, phosphorus and sulfur form acidic solutions. All the acids are strong, they split up to form a hydrogen ion and a conjugate base.

The giant covalent structure of $$SiO_2$$ means that it is not soluble in water. However, it'll react with bases to form salts so it is classed as acidic.

$$SiO_2(s)\ +\ 2NaOH(aq)\ \rightarrow\ Na_2SiO_3(aq)\ +\ H_2O(l)$$​​

Aluminium is partially ionic and partially covalent, it is also insoluble in water. It reacts with both acids and bases to form salts. It is therefore classed as amphoteric. 

Neutralisation reactions

The basic and acidic metal oxides can react with acids and bases and be a part of a neutralisation reaction.

Basic metal oxides neutralise acids:

​$$Na_2O(s)\ +\ 2HCl(aq)\ \rightarrow \ 2NaCl(aq)\ +\ H_2O(l)$$ 

$$MgO(s)\ +\ 2HCl(aq)\ \rightarrow \ MgCl_2(aq)\ +\ H_2O(l)$$ 

Acidic metal oxides neutralise bases:

​$$SiO_2(s)\ +\ 2NaOH(aq)\ \rightarrow\ Na_2SiO_3(aq)\ +\ H_2O(l)$$ 

​$$P_4O_{10}(s)\ +\ 12NaOH(aq)\ \rightarrow\ 4Na_3PO_4(aq)\ +\ 6H_2O(l)$$ 

$$SO_2(g)\ +\ 2NaOH(aq)\ \rightarrow\ Na_2SO_3(aq)\ +\ H_2O(l)$$ 

$$SO_3(g)\ +\ 2NaOH(aq)\ \rightarrow\ Na_2SO_4(aq)\ +\ H_2O(l)$$ 

Amphoteric metal oxides can neutralise both acids and bases: 

​$$Al_2O_3(s)\ +\ 3H_2SO_4(aq)\ \rightarrow\ Al_2(SO_4)_3(aq)\ +\ 3H_2O(l)$$ 

$$Al_2O_3(s)\ +\ 2NaOH(aq)\ +\ 3H_2O(l)\ \rightarrow\ 2NaAl(OH)_4(aq)$$​​