Kanayati ® Chemistry Acids, Bases & Salts
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
- 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:
Acid | Name Of Products | Equation for reaction |
Hydrochloric acid | Metal Chloride and Hydrogen | M + 2 HCl → MCl2 + H2 |
Sulfuric acid | Metal Sulfate and Hydrogen | M + H2SO4 → MSO4 + H2 |
- 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:
Acid | Name Of Products | Equation for reaction |
Hydrochloric acid | Metal Chloride and Water | MOH + HCl → MCl + H2O |
Sulfuric acid | Metal Sulfate and Water | MO + H2SO4 → MSO4 + H2O |
Nitric acid | Metal Nitrate and Water | M(OH)2 + 2HNO3 → M(NO3)2 + 2H2O |
- 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:
Acid | Name Of Products | Equation for reaction |
Hydrochloric acid | Metal Chloride and Carbon dioxide and Water | MCO3 + 2HCl → MCl2 + CO2 + H2O |
Sulfuric acid | Metal Sulfate and Carbon dioxide and Water | MCO3 + H2SO4 → MSO4 + CO2 + H2O |
Nitric acid | Metal Nitrate and Carbon dioxide and Water | MCO3 + 2HNO3 → M(NO3)2 + CO2 + 2H2O |
Properties of Bases
- Bases have pH values of above 7.
- A base which is water-soluble is referred to as an alkali.
![](https://lh7-rt.googleusercontent.com/docsp/AMfWKRCLFbn0N3roRUe7etxcRLMVnWEZCZ7sDvMKK1354C6sR5BTjyxzorcsYwun02JHYgcjqbuT-5lq6hskSLmMLR4yIXp8DqYy98OcP87cS9C0Vit6u39p5idC_xDtANtsVhNthnuPk4W4ACfIOmh-aYBZS4c2ctgKgkFDDH3rFkCEY3jEKPgCDPCB)
- 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
- 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:
Acid | Name Of Products | Equation for reaction |
Hydrochloric acid | Metal Chloride and Water | MOH + HCl → MCl + H2O |
Sulfuric acid | Metal Sulfate and Water | MO + H2SO4 → MSO4 + H2O |
Nitric acid | Metal Nitrate and Water | M(OH)2 + 2HNO3 → M(NO3)2 + 2H2O |
- 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.
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Indicators
- Acid-Alkali indicators can show whether a solution is acidic , neutral or alkaline.
- The table below show the colors for the common indicators
![](https://lh7-rt.googleusercontent.com/docsp/AMfWKRDkNshp-ydxbz4Oce97hlFWf5HQIZv-jsr0eNB5_cax9Qt0grYErt0XdEFI8_KGEtvNUTnxGt2KI-ch50zSkNkxK4IoRPiNA5SPRzSql_IKSuk1HcQdILr53u5cuQyv0eItXoje6bHnS3vXHdE8qQ3M5oERdVGB0Hk2fyMkMNqBhZuOTVi4ya1p)
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.
![](https://lh7-rt.googleusercontent.com/docsp/AMfWKRDFWVEtRZ7FJmVUKAAEg2ss6GKucmEbrqFg1NW1imKFT3ZdtbJhjK4Qcc03o0Y66oRBbH9swZIEVAl69HelEYY-Ec18TMqlfM4rleWiNgPfhJwgYm4FaS-ZnM_pHulTzKK_A_wV5cTpXMXIrrCY46A8-q0APPJ8l8UDhTP3JHxgFiFpXaLCIyf7)
![](https://lh7-rt.googleusercontent.com/docsp/AMfWKRBWPaR6cQVdVs7Fu1VRUV2r-tVSjCLQBIqr8WflsxFFQKM5laj1DvezjNK1NXW_J5AOqDa0gpvMUAT6-h7XRqC4bc_kTqMIRFVgmW3ESK-722kIYpkWk8OszbhhQahnWypl-hj0xfNtKv6CbRTn4DrrWw9VgkAkpy53x-XTXdSTJIdZhigVvyTJ)
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.
- 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