1

Acids An acid is a substance that produces hydrogen ions (H+) when dissolved in water. eg. HCl(aq) → H+(aq) + Cl−(aq) eg. H2SO4 (aq) → 2H+(aq) + SO4

2−aq) Pure acids are usually in the form of simple covalent molecules. The process of acid molecules forming ions in solution is known as ionization. Organic acids: usually weaker & less corrosive, found in fruits/living tissues. Inorganic acids: mineral acids, usually stronger & more corrosive, most are man-made from minerals. Common Acids

Common Name

Chemical Name Formula Description Ions produced in aqueous solutions

hydrochloric acid

aqueous hydrogen chloride

HCl strong acid H+ Cl−

nitric acid aqueous hydrogen nitrate

HNO3 strong acid H+ NO3−

sulfuric acid aqueous hydrogen sulfate

H2SO4 strong acid H+ SO42−

carbonic acid aqueous hydrogen carbonate

H2CO3 weak acid H+ CO32−

acetic acid aqueous ethanoic acid

CH3COOH weak acid H+ CH3COO−

• ”aqueous” means “solution form of”

Properties of Acids

1) Acids have a sour taste. 2) Acid solutions are good electrical conductors. 3) Acids turn blue litmus red. 4) Acids react with bases to form a salt and water ONLY. (neutralization) 5) Acids react with metals to form a salt and hydrogen gas. 6) Acids react with carbonates and hydrogen carbonates to form a salt, water

and carbon dioxide gas.

2

Reactions of Acids

1) acid + base salt + water e.g. nitric acid + sodium oxide sodium nitrate + water 2HNO3(aq) + Na2O(s) 2NaNO3(aq) + H2O(l) Observation: The white solid dissolved to form a colourless solution. e.g. nitric acid + sodium hydroxide sodium nitrate + water HNO3(aq) + NaOH(aq) NaNO3(aq) + H2O(l) Observation: No visible changes. Neutralization is a reaction between hydrogen ions and hydroxide ions to produce water. H+(aq) + OH−(aq) H2O(l)

2) acid + carbonate salt + water + carbon dioxide

e.g. hydrochloric + calcium calcium + water + carbon acid carbonate chloride dioxide 2HCl(aq) + CaCO3(s) CaCl2(aq) + H2O(l) + CO2(g) Observation: The white solid dissolved to form a colourless solution. Effervescence occurred. A colourless and odourless gas is evolved. The gas formed a white precipitate in limewater (aqueous calcium hydroxide).

• ”effervescence” means “bubbling”

• ”precipitate” means “insoluble solid formed in a liquid via a chemical

reaction”

e.g. hydrochloric + potassium potassium + water + carbon acid hydrogen carbonate chloride dioxide HCl(aq) + KHCO3(s) KCl(aq) + H2O(l) + CO2(g)

Observation: Same as previous example.

3) acid + metal salt + hydrogen e.g. sulfuric acid + magnesium magnesium sulfate + hydrogen H2SO4(aq) + Mg(s) MgSO4(aq) + H2(g)

3

Observation: The metal dissolved to form a colourless solution. Effervescence occurred. A colourless and odourless gas is evolved. The gas extinguished a lighted splint with a ‘pop’ sound. Unreactive metals e.g. sulfuric acid + lead lead(II) sulfate + hydrogen H2SO4(aq) + Pb(s) PbSO4(s) + H2(g) Observation: A layer of white deposit is formed on the metal. Little or no effervescence is observed. Lead does not seem to react with sulfuric acid because the initial reaction produces a layer of insoluble lead(II) sulfate around the metal which protects the metal from further reaction. Some metals such as copper, silver, gold and platinum do not react with acids.

Role of Water in Acids When hydrogen chloride gas or citric acid solid is dissolved in non-polar solvents such as toluene, they do not exhibit acidic properties (e.g. turning litmus red or reacting with carbonates to produce carbon dioxide), as their molecules do not ionize. Hence, it is the presence of H+ ions in aqueous solutions that give these acids the acidic properties. Strength of Acids Strong acids are completely ionized in aqueous solution to form H+ ions. e.g. HCl → H+ + Cl−. Weak acids are only partially ionized to form H+ ions with a high proportion of the acid remaining in the undissociated form (ie. remaining as whole molecules, not dissociated into ions). e.g. CH3COOH CH3COO− + H+ Note: Strength refers to the extent of ionization of an acid while concentration refers to the quantity of acid dissolved in the solution. The strength of an acid is not affected by concentration. A dilute acid solution is not the same as a weak acid solution. Likewise, a concentrated acid solution is not the same as a strong acid solution. A strong acid will be completely ionized no matter whether it is concentrated or dilute.

4

Uses of Acids

1) Hydrochloric acid is used to remove impurities such as rust and scale from metals and aluminium alloys. It is also used in leather processing.

2) Ethanoic acid is the main component in vinegar which is used as a food preservative and flavour enhancer. It is also used in making adhesives such as glues.

3) Phosphoric acid is added to food and beverages to give them a sour taste. 4) Sulfuric acid is used in the production of active ingredients in fertilizers such

as ammonium sulfate and superphosphate. 5) Concentrated sulfuric acid is used in the manufacture of detergents to convert

hydrocarbons into organic acids. The organic acids are then reacted with sodium hydroxide to produce the detergent.

6) Dilute sulfuric acid is used in car batteries. It reacts with lead plates and lead(IV) oxide plates in the battery to generate electrical energy which is used to get the car engine running.

Bases Base: any metal oxide or hydroxide that reacts with an acid to produce a salt and water only. Metal oxides contain the O2− ion while metal hydroxides contain the OH−

ion. Most of the bases are insoluble in water. Alkali (soluble base): a substance that produces hydroxide ions (OH−) in aqueous solution. Alkalis are usually group I metal oxides and hydroxides. Common Alkalis Common Name Chemical Name Formula Description caustic soda sodium hydroxide NaOH strong alkali caustic potash potassium hydroxide KOH strong alkali slaked lime calcium hydroxide Ca(OH)2 strong alkali (but only slightly

soluble in water) ammonia solution

aqueous ammonia NH3 (aq) weak alkali

Note: 1. Aqueous ammonia is formed by dissolving NH3 into water. It is to be

written as NH3 (aq). It is alkaline because some ammonia molecules accept protons from water to produce NH4

+ and OH– ions. NH3(g) + H2O(l) NH4

+(aq) + OH−(aq)

5

Properties of Alkalis

1) Alkalis have a bitter taste and feel soapy. 2) Alkalis turn red litmus blue. 3) Alkalis react with acids to form a salt and water only. 4) Alkalis react with ammonium salts upon warming to produce a salt, water and

ammonia gas. 5) Alkalis react with some salts to produce insoluble hydroxides.

Reactions of Bases

1) acid + base salt + water (refer to Reactions of Acids)

2) alkali + ammonium salt salt + water + ammonia

e.g. sodium + ammonium sodium + water + ammonia hydroxide sulfate sulfate 2NaOH(aq) + (NH4)2SO4 (aq) Na2SO4(aq) + 2H2O(l) + 2NH3(g) Observation: A pungent gas was evolved upon warming the mixture. The gas turned moist red litmus blue.

3) alkali + salt salt + metal hydroxide

(containing metal A) (of metal B) (of metal A) (of metal B) e.g. sodium hydroxide + zinc bromide sodium bromide + zinc hydroxide 2NaOH(aq) + ZnBr2 (aq) 2NaBr(aq) + Zn(OH)2(s) Observation: A white precipitate was formed in a colourless solution.

Controlling the pH of Soil Most plants will not grow well if the pH of the soil is below 5 or above 9. To reduce the acidity of the soil, farmers add bases such as calcium oxide (quicklime) or calcium hydroxide (slaked lime). This is known as ‘liming’ the soil. As these bases are not very soluble in water, they are not easily washed away by rain and thus remaining effective for long periods. Strength of Alkalis Strong alkalis are completely ionized in aqueous solution to form OH− ions. e.g. NaOH → Na+ + OH−

.

6

Weak alkalis are only partially ionized to form OH− ions with a high proportion of the alkali molecules remaining in the undissociated form (ie. remaining as whole molecules, not dissociated into ions). e.g. NH3(g) + H2O(l) NH4

+(aq) + OH−(aq) Uses of Bases

1) Magnesium oxide is used as antacid for relieving gastric pain and for making refractory bricks.

2) Magnesium hydroxide is used in toothpaste to neutralize acid on teeth. It is also used as antacid for relieving indigestion.

3) Sodium hydroxide and potassium hydroxide are used in the manufacture of soap and detergents.

4) Calcium oxide (quicklime) and calcium hydroxide (slaked lime) are used to reduce acidity in soil. Calcium oxide is also used to make iron, concrete and cement.

5) Ammonia solution is used to make fertilizers. It is also used in window cleaning solutions.

The pH Scale

The pH is a measure of how acidic or alkaline a solution is in water. The pH scale is numbered between 0 ~ 14. An acidic solution has a pH < 7. In an acidic solution, the concentration of H+ is greater than the concentration of OH−. Having a pH < 7 does not mean that OH− ions are absent in the acidic solution. The smaller the pH, the more acidic is the solution and the greater is the concentration of H+ ions. A stronger acid will have a lower pH compared to a weaker acid with the same concentration. An alkaline solution has a pH > 7. In an alkaline solution, the concentration of H+ is less than the concentration of OH−. Having a pH > 7 does not mean that H+ ions are absent in the alkaline solution. The bigger the pH, the more alkaline is the solution and the greater is the concentration of the OH- ions. A stronger alkali will have a higher pH compared to a weaker alkali with the same concentration. In a neutral solution (pH = 7), the concentration of H+ is equal to the concentration of OH−.

7

Indicators pH can be measured with indicators and pH meters. An indicator is a chemical that is able to change its colour when it is in an acidic solution to another colour when it is in an alkaline solution. The Universal Indicator contains a mixture of dyes and gives different colours in solutions of different pH values. It is available in paper or solution form. The pH values and corresponding colours of the universal indicator are as shown below.

http://www.abundanthealthcenter.com/blog/wp-content/uploads/2011/04/PH-Scale.jpg

A pH meter is an electrical device which is more accurate than indicators. It requires calibration and is dipped into the solution directly to measure pH. Variation of pH during a neutralisation reaction with an alkali can be studied using data logging and a computer. The pH meter measures the pH as alkali is added to acid in the beaker.

8

pH values of common substances

http://staff.jccc.net/pdecell/chemistry/phscale.gif

9

Oxides Non-Metallic Oxides Acidic oxides are oxides of non-metals. Most of them dissolve with water to produce acids.

e.g. sulfur dioxide + water sulfurous acid SO2(g) + H2O(l) H2SO3(aq)

Acidic oxides also react with alkalis to form a salt and water.

e.g. sulfur dioxide + potassium hydroxide potassium sulfite + water SO2(g) + 2KOH(aq) K2SO3(aq) + H2O(l)

Examples of Acidic Oxides

Acidic Oxide Formula Acid Produced with Water

sulphur trioxide SO3 sulphuric acid, H2SO4 sulphur dioxide SO2 sulphurous acid, H2SO3 carbon dioxide CO2 carbonic acid, H2CO3

phosphorous(V) oxide P4O10 phosphoric acid, H3PO4

Silicon(IV) oxide is an example of an acidic oxide that does not dissolve in water. It is however still considered an acidic oxide since it reacts with sodium hydroxide (a base) to form sodium silicate (a salt) and water.

e.g. silicon(IV) oxide + sodium hydroxide sodium silicate + water SiO2(s) + 2NaOH(aq) Na2SiO3(aq) + H2O(l)

Some non-metal oxides form neutral oxides which are insoluble in water. Neutral oxides show neither basic nor acidic properties.

Examples of Neutral Oxides

Neutral Oxide Formula water H2O

carbon monoxide CO nitric oxide NO

10

Metallic Oxides Basic oxides are oxides of metals. They are solids at room temperature. All basic oxides are insoluble except those of calcium, barium, and group I metals.

e.g. calcium oxide + water calcium hydroxide CaO(s) + H2O(l) Ca(OH)2(aq)

Calcium hydroxide is sparingly soluble in water. In water it forms a dilute solution called limewater. Limewater is thus Ca(OH)2 (aq). Basic oxides also react with acids to form a salt and water.

e.g. calcium oxide + nitric acid calcium nitrate + water CaO(s) + 2HNO3(aq) Ca(NO3)2(aq) + H2O(l)

Examples of Basic Oxides

Basic Oxide Formula

magnesium oxide MgO sodium oxide Na2O calcium oxide CaO

copper(II) oxide CuO

Some metal oxides form amphoteric oxides which react with both acids and bases to form a salt and water.

e.g. zinc oxide + sodium hydroxide sodium zincate + water ZnO(s) + 2NaOH(aq) Na2ZnO2(aq) + H2O(l)

e.g. zinc oxide + hydrochloric acid zinc chloride + water

ZnO(s) + 2HCl(aq) ZnCl2(aq) + H2O(l)

Examples of Amphoteric Oxides

Amphoteric Oxide Formula Salt produced in NaOH(aq) aluminium oxide Al2O3 sodium aluminate (NaAlO2)

lead(II) oxide PbO sodium plumbate(II) (Na2PbO2)

zinc oxide ZnO sodium zincate (Na2ZnO2)

11

How to Classify an Unknown Oxide

Uses of Sulfur Dioxide

1) Sulfur dioxide is used as a bleach in the manufacture of wood pulp for paper. The wood pulp used for making paper is coloured because it contains dyes. When sulfur dioxide is added to the pulp, it removes oxygen from the dyes, causing the pulp to turn white.

2) Sulfur dioxide is used as a food preservative. In the food industry, sulfur dioxide is added to food in small amounts to prevent the growth of moulds and bacteria.

Is it a metallic oxide?