3.1D Properties of Oxides
Definition of Acids and Bases
From last year...
A Bronsted-Lowry acid is a substance that donates protons, H+
e.g. HCl (aq), HNO3 (aq)
HCl (aq) = H+ and Cl-
start with H
A Bronsted-Lowry base is a substance that accepts protons, H+
e.g. NaOH(aq)
NaOH = Na+ and OH- (accepts H+ to form H2O)
end with OH
Definition of Acids
Another way to describe acids and bases is through their ability to accept or donate a pair of electrons.
A Lewis acid can accept an electron pair
e.g. H+, Fe2+, Al3+
A Lewis base can donate an electron pair
Cl-, OH-, NH3
Metal Oxides
Many metal oxides are Lewis bases. They react with water to form hydroxides by donating an electron pair to hydrogen in water.
Reactions of alkali metal oxides with water have the general equation
M = Metal
M2O(s) + H2O(l) > 2 MOH(aq)
Reactions of group 2 metal oxides with water have the general equation
MO(s) + H2O(l) > M(OH)2 (aq)
Non-Metallic Oxides
Non-metallic oxides are Lewis acids. They react with water to form other acids by accepting an electron pair from oxygen in water.
Examples
CO2 (g) + H2O(l) > H2CO3 (aq)
SO3 (l) + H2O(l) > H2SO4 (aq)
P4O10 (s) + 6 H2O > 4 H3PO4 (aq)
Amphoteric
Amphoteric substances are able to behave both as a Lewis acid and a
Lewis base.
Aluminum oxide acts as a Lewis base when reacting with NaOH(aq)
Al2O3 (s) + 2 NaOH(aq) + 3 H2O(l) > 2 Na[Al(OH)4](aq)
Aluminum oxide acts as a Lewis acid when reacting with HCl (aq)
Al2O3 (s) + 6 HCl (aq) > 2 AlCl3 + 3 H2O(l)
Trends
Going across a period, the oxides of the elements become less basic and more acidic.
Oxide | Na2O(s) | MgO(s) | Al2O3 (s) | SiO2 (s) | P4O10 (s) | SO3 (l) |
|---|---|---|---|---|---|---|
Acid or Base? | Basic | Basic | Amphoteric | Acidic | Acidic | Basic |
Acid Rain
Pure water has a pH of 7.0
Rainwater is naturally acidic due to the presence of dissolved CO2 which forms carbonic acid. It has a pH of 5.6
Oxides of sulfur and nitrogen are more acidic than oxides of carbon. When these gases dissolve in rainwater, the rain is more acidic than normal. This is known as , which has a pH of less than 5.6.
Oxides of sulfur and nitrogen are produced naturally by volcanic eruptions and decomposing vegetation. They can also be released as pollutants from industrial processes such as the burning of fossil fuels.
Ocean Acidification
Oceans absorb a large proportion of the CO2 released into the atmosphere. As a result, carbonic acid is formed in the ocean. The increased acidity can affect biodiversity, including coral reefs and shellfish
Definition of Acids and Bases
From last year...
A Bronsted-Lowry acid is a substance that donates protons, H+
e.g. HCl (aq), HNO3 (aq)
HCl (aq) = H+ and Cl-
start with H
A Bronsted-Lowry base is a substance that accepts protons, H+
e.g. NaOH(aq)
NaOH = Na+ and OH- (accepts H+ to form H2O)
end with OH
Definition of Acids
Another way to describe acids and bases is through their ability to accept or donate a pair of electrons.
A Lewis acid can accept an electron pair
e.g. H+, Fe2+, Al3+
A Lewis base can donate an electron pair
Cl-, OH-, NH3
Metal Oxides
Many metal oxides are Lewis bases. They react with water to form hydroxides by donating an electron pair to hydrogen in water.
Reactions of alkali metal oxides with water have the general equation
M = Metal
M2O(s) + H2O(l) > 2 MOH(aq)
Reactions of group 2 metal oxides with water have the general equation
MO(s) + H2O(l) > M(OH)2 (aq)
Non-Metallic Oxides
Non-metallic oxides are Lewis acids. They react with water to form other acids by accepting an electron pair from oxygen in water.
Examples
CO2 (g) + H2O(l) > H2CO3 (aq)
SO3 (l) + H2O(l) > H2SO4 (aq)
P4O10 (s) + 6 H2O > 4 H3PO4 (aq)
Amphoteric
Amphoteric substances are able to behave both as a Lewis acid and a
Lewis base.
Aluminum oxide acts as a Lewis base when reacting with NaOH(aq)
Al2O3 (s) + 2 NaOH(aq) + 3 H2O(l) > 2 Na[Al(OH)4](aq)
Aluminum oxide acts as a Lewis acid when reacting with HCl (aq)
Al2O3 (s) + 6 HCl (aq) > 2 AlCl3 + 3 H2O(l)
Trends
Going across a period, the oxides of the elements become less basic and more acidic.
Oxide | Na2O(s) | MgO(s) | Al2O3 (s) | SiO2 (s) | P4O10 (s) | SO3 (l) |
|---|---|---|---|---|---|---|
Acid or Base? | Basic | Basic | Amphoteric | Acidic | Acidic | Basic |
Acid Rain
Pure water has a pH of 7.0
Rainwater is naturally acidic due to the presence of dissolved CO2 which forms carbonic acid. It has a pH of 5.6
Oxides of sulfur and nitrogen are more acidic than oxides of carbon. When these gases dissolve in rainwater, the rain is more acidic than normal. This is known as , which has a pH of less than 5.6.
Oxides of sulfur and nitrogen are produced naturally by volcanic eruptions and decomposing vegetation. They can also be released as pollutants from industrial processes such as the burning of fossil fuels.
Ocean Acidification
Oceans absorb a large proportion of the CO2 released into the atmosphere. As a result, carbonic acid is formed in the ocean. The increased acidity can affect biodiversity, including coral reefs and shellfish
