Part 1: The Characteristic Properties of Acids & Bases

Properties of Acids

• Acids have pH values of below 7, have a sour taste and are corrosive.
• Acids form acidic solutions in water.
• In acidic conditions blue litmus paper turns red and methyl orange indicator turns red
• Acids are substances that can neutralize a base, forming salt and water.
• When acids react, they will lose electrons to form positively charged hydrogen ions (H+).
• The presence of H+ ions is what makes a solution acidic.

Example: Hydrochloric Acid

• HCl (aq)   →    H+ (aq)    +    Cl- (aq)

Typical reactions of acids

1. Acids and metals

• Only metals above hydrogen in the reactivity series will react with dilute acids.
• When acids react with metals they form a salt and hydrogen gas: 

• Acid    +    Metal   →    Salt    +    Hydrogen

• Examples of reaction between acids and metals:

2. Acids with bases (alkalis)

• Metal oxides and metal hydroxides can act as bases.
• When they react with acid, a neutralisation reaction occurs.
• In all acid-base neutralisation reactions, salt and water are produced.

• Acid    +    Base   →    Salt    +    Water

• Examples of reaction between acids and bases:

3. Acids with metal carbonates

• Acids will react with metal carbonates to form the corresponding metal salt, carbon dioxide and water:

• Acid  +  Metal Carbonate → Salt  +  Carbon Dioxide  +  Water

• Examples of reaction between acids and carbonates:

Properties of Bases

• Bases have pH values of above 7.
• A base which is water-soluble is referred to as an alkali.

• In basic (alkaline) conditions red litmus paper turns blue and methyl orange indicator turns yellow.
• Bases are substances which can neutralize an acid, forming a salt and water.
• Bases are usually oxides or hydroxides of metals.
• When alkalis react, they gain electrons to form negative hydroxide ions (OH–).
• The presence of the OH– ions is what makes the aqueous solution an alkali.

Example: Sodium Hydroxide

• NaOH (s)   →    Na+ (aq)    +    OH– (aq)

Typical reactions of bases

1. Bases and acids

• When they react with an acid, a neutralisation reaction occurs.
• Acids and bases react together in a neutralisation reaction and produce a salt and water:

• Acid    +    Base   →    Salt    +    Water

• Examples of reaction between bases and acids:

2. Alkalis and ammonium salts

• Ammonium salts undergo decomposition when warmed with an alkali.
• Even though ammonia is itself a weak base, it is very volatile and can easily be displaced from the salt by another alkali.
• Salt, water and ammonia are produced.
• Example:

• NH4Cl + NaOH →NaCl + H2O + NH3

• This reaction is used as a chemical test to confirm the presence of the ammonium ion (NH4+).
• Alkali is added to the substance with gentle warming followed by the test for ammonia gas using damp red litmus paper.
• The litmus paper will turn from red to blue if ammonia is present.

Neutrality, Relative Acidity & Alkalinity

The pH scale

• The pH scale is a numerical scale which is used to show how acidic or alkaline a solution is.
• It goes from 1 – 14 (extremely acidic substances can have values of below 1).
• All acids have pH values of below 7, all alkalis have pH values of above 7.
• The lower the pH then the more acidic the solution is.
• The higher the pH then the more alkaline the solution is.
• A solution of pH 7 is describe as being neutral e.g. water. A Neutral solution is neither acidic nor alkaline.

Indicators

• Acid-Alkali indicators can show whether a solution is acidic , neutral or alkaline.
• The table below show the colors for the common indicators

Universal indicator

• Universal indicator is a mixture of different indicators which is used to measure the pH.
• A drop is added to the solution and the colour is matched with a colour chart which indicates the pH which matches specific colors.

The importance of pH and soil acidity

• Soil pH is analyzed to indicate the acidity or alkalinity of soil.
• Most plants favor a pH value of between 5 and 8.
• Changes in soil which cause a pH to be outside this range adversely affect plant processes resulting in reduced growth and crop yield.
• Soils may become acid from acid rain, overuse of fertilizers which contain ammonium salts or by the excessive breakdown of organic matter by bacteria.
• Crushed or powdered limestone (calcium carbonate) or lime (calcium oxide) or slaked lime (calcium hydroxide) is added to neutralize the excess acidity in the soil.
• The addition process must be carefully monitored though, as if added in excess, further damage could be done if the pH goes too high.

Proton Transfer, Weak & Strong Acids & Bases

Proton transfer

• The earlier definition of an acid and a base can be extended.
• In terms of proton transfer, we can further define each substance in how they interact with protons.

Acids

• Acids are proton donors as they ionize in solution producing protons, H+ ions.
• These H+ ions make the aqueous solution acidic.

Bases (Alkalis)

• Bases (alkalis) are proton acceptors as they ionize in solution producing OH– ions which can accept protons.
• These OH– ions make the aqueous solution alkaline.

Strong acids and bases

• Acids and alkalis can be either strong or weak, depending on how many ions they produce when dissolved in water.
• Strong acids and bases ionize completely in water, producing solutions of very low pH for an acid or very high pH for a base.
• Strong acids include HCl and H2SO4 and strong bases include the Group I hydroxides.

Weak acids and bases

• Weak acids and bases partially ionize in water and produce pH values which are closer to the middle of the pH scale.
• Weak acids include organic acids such as ethanoic acid, CH3COOH and weak bases include aqueous ammonia.
• For both weak acids and bases, there is usually an equilibrium set-up between the molecules and their ions once they have been added to water.
• Example of a weak acid: propanoic acid.

• CH3CH2COOH ⇌ H+ + CH3CH2COO-

• Example for a weak base: aqueous ammonia.

• NH3 + H2O ⇌ NH4+ + OH-

• In both cases the equilibrium lies to the left, indicating a high concentration of intact acid / base molecules, with a low concentration of ions in solution.

Effect of concentration on strong and weak acids and alkalis

• A concentrated solution of either an acid or a base is one that contains a high number of acid or base molecules per dm3 of solution.
• It does not necessarily mean that the acid or base is strong though, as it may be made from a weak acid or base which does not dissociate completely.
• For example, a dilute solution of HCl will be more acidic than a concentrated solution of ethanoic acid, since most of the HCl molecules dissociate but very few of the CH3COOH do.

Types of Oxides

Acid and basic oxides

• Acidic and basic oxides have different properties and values of pH.
• The difference in their pH stems from whether they are bonded to a metal or a nonmetal element.
• The metallic character of the element influences the acidic or alkaline behaviour of the molecule.

Acid oxides

• Acidic oxides are formed when a nonmetal element combines with oxygen.
• They react with bases to form a salt and water.
• When dissolved in water they produce an acidic solution with a low pH.
• Common examples include SO2 and SiO2.

Basic oxides

• Basic oxides are formed when a metal element combines with oxygen.
• They react with acids to form a salt and water.
• When dissolved in water they produce a basic solution with a high pH.
• Common examples include NaOH, KOH and Ca(OH)2.

Neutral & Amphoteric Oxides

Neutral oxides

• Some oxides do not react with either acids or bases and thus are said to be neutral.
• Examples include N2O, NO and CO.

Neutral oxides

• Amphoteric oxides are a curious group of oxides that can behave. as both acidic and basic, depending on whether the other reactant is an acid or a base
• In both cases a salt and water is formed.
• Two most common amphoteric oxides are zinc oxide and aluminum oxide.
• The hydroxides of both of these elements also behave amphoterically.
• Example of aluminium oxide behaving as a base:

• Al2O3 + 6HCl → 2AlCl3 + 3H2O

• Example for an aluminium oxide behaving as an acid:

• Al2O3 + 2NaOH → 2NaAlO2 + H2O