Acids and Bases (6092) - Review Flashcards

A comprehensive set of practice questions covering acids, bases, pH, indicators, neutralisation, titration, and applications in agriculture and industry.

What ions do acids produce in aqueous solution?

Hydrogen ions (H+) are produced when acids dissociate in water.

What ions do alkalis (bases) produce in aqueous solution?

Hydroxide ions (OH−) are produced when alkalis dissolve in water.

How can neutrality, relative acidity, and alkalinity be described in terms of H+, OH− concentrations, Universal Indicator color, and the pH scale?

More H+ means higher acidity and lower pH; more OH− means higher alkalinity and higher pH; neutral solutions have equal H+ and OH− and pH ≈ 7; Universal Indicator changes color across the pH scale to reflect these changes.

How do strong acids differ from weak acids in terms of ionisation in water?

Strong acids ionise completely in water; weak acids ionise only partially.

What are the characteristic properties of acids in reactions with metals, bases, and carbonates?

Acids react with reactive metals to form salts and hydrogen gas; with carbonates to form salts, carbon dioxide, and water; with bases to form salts and water (neutralisation).

What is the neutralisation reaction between hydrogen ions and hydroxide ions?

H+ (aq) + OH− (aq) → H2O (l).

Why is pH control important in soils and how can excess acidity be treated?

To maintain healthy plant growth; excess acidity can be treated with calcium hydroxide (lime) which neutralises acidity.

What are the characteristic properties of bases?

Bitter taste, slippery feel, change indicator colors (usually blue/purple with universal indicators), conduct electricity; react with acids to form salts and water; can react with ammonium salts to release ammonia.

How are oxides classified as acidic, basic, amphoteric, or neutral based on metallic/non-metallic character?

Metal oxides are typically basic; non-metal oxides are typically acidic; some oxides are amphoteric (e.g., ZnO, Al2O3); some oxides are neutral (e.g., CO, NO, H2O).

Give examples of strong and weak inorganic acids and name some organic acids listed in the notes.

Strong inorganic acids: HCl, H2SO4, HNO3. Weak inorganic acids: H2CO3, H3PO4. Organic acids include ethanoic acid (CH3COOH), lactic acid (CH3CH(OH)COOH), citric acid, malic acid, tartaric acid, amino acids (NH2CH2COOH).

What is the basicity of acids and how is it related to ionisation?

Basicity is the total number of H+ formed when the acid ionises in water (e.g., HCl and HNO3 = 1 H+; H2SO4 = 2 H+; H3PO4 = 3 H+).

What is pH and how is it calculated?

pH = −log10[H+]; pH < 7 is acidic, pH = 7 is neutral, pH > 7 is alkaline.

Which common indicators are used, and how do they indicate pH changes?

Litmus (red in acid, blue in base); Methyl orange (red in acid, yellow in base); Phenolphthalein (colorless in acid, pink in base); universal indicator covers the full pH range.

What is a universal indicator and what does it represent?

A mixture that changes color at different pH values; it represents acidity/alkalinity on the 0–14 pH scale.

What is a pH meter and why is it more reliable than indicators?

An electronic device with a pH probe that gives precise pH readings (often to two decimal places), more reliable than color indicators.

Describe the general reaction of acids with metals, with an example.

Acid + reactive metal → salt + H2 gas (e.g., 2HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)). Ionic form: 2H+(aq) + Mg(s) → Mg2+(aq) + H2(g).

Describe the general reaction of acids with carbonates, with an example.

Acid + carbonate → salt + CO2 + H2O (e.g., CuCO3(s) + H2SO4(aq) → CuSO4(aq) + CO2(g) + H2O(l)). Ionic form: CO3^2−(aq) + 2H+(aq) → CO2(g) + H2O(l).

Describe the neutralisation reaction between acids and bases, with ionic form.

Acid + base → salt + water; H+ (aq) + OH− (aq) → H2O (l). Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l).

What happens in an alkali–ammonium salt reaction and provide an example?

Alkali + ammonium salt → salt + water + ammonia (e.g., Ca(OH)2(aq) + 2NH4Cl(aq) → CaCl2(aq) + 2NH3(g) + 2H2O(l)). Ionic: NH4+(aq) + OH−(aq) → NH3(g) + H2O(l).

What is a precipitation reaction involving hydroxides and give an example?

Alkali + salt → insoluble hydroxide precipitate + soluble salt (e.g., 2NaOH(aq) + CuSO4(aq) → Cu(OH)2(s) + Na2SO4(aq)). Ionic: 2OH−(aq) + Cu2+(aq) → Cu(OH)2(s).

How can precipitation reactions be used to identify cations?

The color of the precipitate and its solubility in excess NaOH or NH3 help identify the cation present.

What is a titration and what is it used for?

A laboratory method to determine the concentration of an acid or alkali by precisely adding solution with a burette; involves a pH change and a titration curve.

What is a titration curve and what information does it provide?

Plot of pH against added volume of titrant; shows the end point and helps select an appropriate indicator.

What is the typical end point for a strong acid–strong base titration and why?

End point around pH 7 (neutral) because the reaction goes to completion with equal moles of H+ and OH−.

What can temperature change reveal in neutralisation reactions?

Neutralisation is exothermic; temperature increases as acid and base react.

Which fertilisers are formed by neutralisation of ammonia with acids, and what are their formulas?

Ammonium nitrate NH4NO3 and ammonium sulfate (NH4)2SO4 are formed from NH3 with HNO3 and H2SO4, respectively.

Why is calcium carbonate often preferred over calcium hydroxide for neutralising acidic soils?

Calcium carbonate does not react with ammonium fertilisers to release NH3 (no nitrogen loss) and is less likely to overshoot pH because it is less soluble; it also stays in soil longer.

List some uses of acids and sulfuric acid as given in the notes.

HCl: rust removal and metal cleaning; H2SO4: fertilisers, detergents, electrolytes, paints; HNO3: fertilisers, explosives; organic acids: in foods and industry (citric, acetic, lactic, etc.).