Acids and Their Reactions Explained

Properties of Acids

• Acids have a pH below 7 and a sour taste, and are corrosive.
• These neutralize bases to produce salt and water.
• Acids in water form positively charged hydrogen ions (H⁺).
• The presence of H⁺ ions makes a solution acidic.
• An example of this is hydrochloric acid: HCl (aq) → H⁺ (aq) + Cl⁻ (aq).

Acids and Metals

• Only metals above hydrogen in the reactivity series react with dilute acids.
• Metal and acid produce salt and hydrogen gas.
• The salt name depends on the anion within the acid used in in the reaction.
• Hydrochloric acid reacts with magnesium to produce magnesium chloride and hydrogen: Mg + 2HCl → MgCl₂ + H₂
• Sulfuric acid and magnesium react to form magnesium sulfate and hydrogen: Mg + H₂SO₄ → MgSO₄ + H₂
• Nitric acid and magnesium produce magnesium nitrate and hydrogen: Mg + 2HNO₃ → Mg(NO₃)₂ + H₂

Acids and Bases

• Metal oxides and metal hydroxides (alkalis) can act as bases.
• Neutralization occurs when they react with acids.
• Salt and water are produced in acid-base neutralization reactions.
• In general, acid + base → salt + water.
• Hydrochloric acid with magnesium chloride produces magnesium chloride and water: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O
• Magnesium sulfate and water are produced from sulfuric acid and magnesium sulfate: MgO + H₂SO₄ → MgSO₄ + H₂O.
• Magnesium nitrate and water are produced from nitric acid with magnesium nitrate: Mg(OH)₂ + 2HNO₃ → Mg(NO₃)₂ + 2H₂O.

Acids with Metal Carbonates

• The corresponding metal salt, carbon dioxide, and water are produced when acids react with metal carbonates.
• acid + metal carbonate → salt + carbon dioxide + water
• Magnesium chloride, carbon dioxide, and water forms when hydrochloric acid reacts with a magnesium carbonate: MgCO₃ + 2HCl → MgCl₂ + CO₂ + H₂O
• Sulfuric acid reacts with magnesium carbonate to form magnesium sulfate, carbon dioxide, and water: MgCO₃ + H₂SO₄ → MgSO₄ + CO₂ + H₂O
• Magnesium nitrate, carbon dioxide, and water produced from reaction of nitric acid with magnesium carbonate: MgCO₃ + 2HNO₃ → Mg(NO₃)₂ + CO₂ + H₂O

Indicators

• Used to distinguish between acids and alkalis with two-colour indicators.
• Litmus is the most common indicator extracted from lichens in many plants.
• Appear in different colors in acids and alkalis, synthetic indicators are organic compounds.
• Thymolphthalein and methyl orange are often used as synthetic indicators in acid-alkali titrations.
• Litmus turns red in acid and blue in alkali.
• Methyl orange shows red in acid and yellow in alkali.
• Thymolphthalein is colorless in acid but turns blue in alkali.

Synthetic Indicators

• Synthetic indicators are used to identify the endpoint in titrations because they have distinct color changes when an acid is neutralized by alkali and vice versa.
• Because its color transition is not sharp and travels through a purple transition color in neutral solutions, litmus is unsuitable for titrations because endpoint detection is challenging.
• In the form of red and blue papers that can be dipped into solutions or used to test gases, litmus is helpful as an indicator paper.

Properties of Bases and Alkalis

• Bases’ pH values are above 7.
• Alkali refers to a base that dissolves in water.
• In basic conditions:
  • Red litmus paper turns blue.
  • Methyl orange indicator turns yellow.
  • Thymolphthalein indicator turns blue.
• Acids can be neutralized by bases, which create salt and water.
• Metals' oxides or hydroxides are frequently the basis.
• Water is added to alkalis to create negative hydroxide ions (OH-).
• The presence of OH- ions makes the aqueous solution an alkali.
• For instance, sodium hydroxide: NaOH (s) → Na+ (aq) + OH- (aq).

Acids and Bases Reactions

• When bases react with an acid, neutralisation occurs.
• Acids and bases combine and produce salt and water.
• In general: Acid + base → salt + water
• Hydrochloric acid reacts with magnesium chloride to form magnesium chloride and water: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O
• Magnesium sulfate and water are produced from sulfuric acid and magnesium sulfate: MgO + H₂SO₄ → MgSO₄ + H₂O.
• Magnesium nitrate and water are produced from nitric acid and magnesium nitrate: Mg(OH)₂ + 2HNO₃ → Mg(NO₃)₂ + H₂O

Alkalis and Ammonium Salts

• When warmed with an alkali, ammonium salts decompose.
• Even though ammonia is a weak base, it can easily be displaced from the salt by another alkali because it is volatile.
• Where a salt, water, and ammonia are produced, for example: NH₄Cl + NaOH → NaCl + H₂O + NH₃
• To check for the presence of the ammonium ion (NH₄⁺), use this reaction as a chemical test.
• Test for ammonia gas using wet red litmus paper after alkali has been added to the substance and gently warmed.
• The damp litmus paper turns from red to blue if ammonia is present.

Neutralization Reactions

• Acids provide hydrogen ions, H⁺.
• Hydroxide ions, OH–, are produced by bases (or alkalis).
• When they react with one another during a neutralization reaction, H⁺ ions combine with OH– ions to produce water.
• As water has a pH of 7, this is the net ionic equation for all acid-base neutralizations that results in neutral solutions: H⁺ (aq) + OH– (aq) ⟶ H₂O (l)
• Certain acid reactions do not result in neutralisation.
• For instance, even though a salt is produced when a metal reacts with an acid, the absence of water produced prevents it from satisfying the definition of neutralization.

pH Scale

• The pH scale ranges from 1 to 14.
• Acids have pH values below 7, alkalis above 7.
• Lower pH corresponds to a more acidic solution.
  • pH 0-2 solutions are strong acid solutions
  • pH 3-6 solutions are weak acid solutions.
• Higher pH = More Alkaline solution:
  • pH 8-11 solutions are weak alkali solutions.
  • pH 12-14 solutions are strong alkali solutions.
• A pH of 7 is considered neutral.

pH Scale and Hydrogen Ions

• Acids contain hydrogen ions in the solution.
• The pH is lower, and the acid is stronger.
• Hydroxide ions in a solution increases the pH
• Therefore, pH is a concentration of H+ ions with an inverse relationship.
• Since the pH scale is logarithmic: each change of 1 indicates a 10 times change in concentration.

Universal Indicator

• An acid with a pH of 3 has ten times the concentration of H⁺ ions than an acid with a pH of 4. An acid with a pH of 2 has 10 x 10 = 100 times the concentration of H⁺ ions than an acid with a pH of 4
• As acid strength correlates with the quantity of hydrogen ions in solution, the more hydrogen ions, the lower the pH and vice versa.
• The mixture of indicators used to measure pH is called the universal indicator.
• To determine the pH, a drop is added to the solution and the drop's color is then compared to a color chart demonstrating the pH correlation.

Acids, Bases, and Proton Transfer

• Acids are proton donors as they ionize in solution producing protons, which are H⁺ ions.
• These H+ ions result in the formation of an acidic aqueous solution.
• These bases accept the protons provided by the acid, making them proton acceptors.
• Strong acids completely dissociate (or ionise) in water, producing low pH solutions.
• Hydrochloric acid (HCl) and sulfuric acid (H₂SO₄) are strong acids.
• Example of a strong acid: hydrochloric acid: HCl (aq) → H⁺ (aq) + Cl⁻ (aq)

Weak Acids

• Lower pH values, but pH values still less than 7, are produced when weak acids only partially dissociate (or ionize) in water.
• After weak acids have been added to water, the molecules and their ions become an equilibrium.
• For example, propanoic acid: CH₃CH₂COOH ⇌ H⁺ + CH₃CH₂COO⁻
• The equilibrium lies to the left indicating a high intact acid concentration and a low H⁺ ion concentration.
• Ethanoic acid is another weak acid example that reacts with alkalis like sodium hydroxide to produce ethanoate salts.
• ethanoic acid + sodium hydroxide → sodium ethanoate + water can be described as CH₃COOH + NaOH → CH₃COONa + H₂O.

Concentration and Acid Strength

• A concentrated acid solution has more acid molecules per dm³.
• The acid is not necessarily strong, because the molecules do not completely dissociate.
• For example, a dilute HCl solution is more acidic than a concentrated ethanoic acid solution, since most of the HCl molecules dissociate, but not the CH₃COOH do.

Classifying Oxides

• Oxides are made up of one or more oxygen atoms and one other element, for example: MgO, ZnO, K₂O, CO₂, SO₂, H₂O
• Acid-base characteristics determine how oxides are classified.
• Acidic oxides are formed when a non-metal element and oxygen combine.
• They react with bases forming salt and water.
• Acidic oxides, when dissolved in water, produce an acidic solution with a low pH.
• Common examples include: CO₂, SO₂, NO₂ and SiO₂
• Formed from a metal element combines with oxygen, basic oxides react with acids forming salt and water.
• Common examples include CuO and CaO which produce a basic solution with a high pH when dissolved in water.

Amphoteric and Neutral Oxides

• Metals form basic oxides whereas non-metals form acidic oxides.
• Neutral oxides do not react with either acids or bases, where examples include N₂O, NO and CO.
• Certain oxides are amphoteric and can act as both acid and base depending on the reactant which forms salt and water.
• Zinc oxide (ZnO) and aluminum oxide (Al₂O₃) are 2 amphoteric oxides and their hydroxides behave amphoterically.
• Aluminum oxide behaving as a base: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
• Aluminum oxide behaving as an acid: Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O