Chapter 9 — Acids, bases, salts and solutions
Classify by behaviour, not by a name alone
An acid in the introductory water-solution account increases the concentration of hydrogen ions (more precisely, hydronium species in water); a base can increase hydroxide-ion concentration or accept a proton in the broader chemical account. For police-exam questions, common aqueous examples are hydrochloric acid, sulfuric acid, sodium hydroxide and calcium hydroxide. Their names do not tell you that every bottle has the same concentration. A highly diluted strong acid and a concentrated weak acid require two different descriptions: strength concerns the degree of ionisation under specified conditions, whereas concentration concerns how much solute is present per quantity of solution.
Acids can turn blue litmus red and bases can turn red litmus blue in the standard aqueous test. Litmus gives a broad classification, not an exact numerical pH. Universal indicator gives a range of colours that can be matched approximately to a scale. If a stem asks for a precise measurement, a calibrated pH meter is the more suitable instrument. Indicator colours can also be obscured by a strongly coloured sample, so the result must be interpreted in context. The chapter's diagrams teach reading results, not a procedure for handling reagents.
At about 25 °C, a neutral aqueous solution has pH 7 in the familiar school convention; lower pH indicates acidity and higher pH indicates basicity. The pH scale is logarithmic, so a difference of one pH unit represents a tenfold change in hydrogen-ion activity under the simplified comparison. The often printed 0–14 range is a standard teaching range for many dilute aqueous solutions, not a universal physical boundary for every solution. Do not answer that pH can never fall below 0 or exceed 14. Likewise, do not classify every substance with hydrogen in its formula as an acid: glucose contains hydrogen, but that fact alone does not make its water solution an acid in the school definition.
Neutralisation and salt names
In a common neutralisation reaction, an acid reacts with a base to form a salt and water. A simple balanced example is HCl + NaOH → NaCl + H₂O. The hydrogen and hydroxide components form water, while the remaining ions form the salt in this example. The word salt in chemistry is broader than table salt. Copper sulfate and sodium carbonate are salts, though neither is ordinary edible sodium chloride. Some acid-base reactions can be described in more general ways and do not fit the naive “every salt is neutral in water” claim. The pH of a salt solution can depend on the ions and their reaction with water.
When naming a salt in a basic exam problem, keep both parent contributions in view. Sulfuric acid with sodium hydroxide can form sodium sulfate under full neutralisation: H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O. One molecule of the diprotic acid in this balanced representation requires two formula units of NaOH for complete neutralisation. If the question gives volumes of two solutions, do not assume they neutralise in a one-to-one volume ratio. Their concentrations and the balanced reaction ratio both matter. The equation gives particle ratios; it does not reveal a volume ratio without concentration information.
Carbonates and hydrogen carbonates can react with suitable acids to release carbon dioxide. For the basic reaction CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂, count each atom to check balance. Carbon dioxide can make limewater appear milky under the standard school test, but a single observed bubble is not a unique chemical identification. In a well-posed MCQ, the question should specify the reagent and relevant observation. Avoid memorising “fizzing means acid” because gas can appear in other reactions as well.
Water of crystallisation is a fixed number of water molecules incorporated in a hydrated crystal's formula unit. The notation CuSO₄·5H₂O indicates five water molecules per copper sulfate formula unit in that specified hydrated material. The dot is a composition notation, not a multiplication sign meaning the formula contains five extra sulfate groups. Heating can change hydration state in some salts; an observed colour change should be linked to the exact stated salt. A general question should not claim every white salt contains water of crystallisation.
Solutions, suspensions and colloids
A true solution is homogeneous at the scale of ordinary observation. The solute is dispersed in the solvent; the amount of each may vary within the range in which a solution forms. Dissolving a small amount of sugar in water does not create a new chemical compound with a fixed formula. Sugar remains sugar in a mixture, although its particles are dispersed. When no more solute dissolves at a stated temperature under the given conditions, the solution is saturated. Solubility can depend on temperature, pressure and the solute–solvent pair; do not apply one substance's temperature trend to all others.
A suspension has dispersed particles large enough to settle under suitable conditions and can often be separated by ordinary filtration. A colloid has dispersed particles between the simple solution and coarse-suspension picture and can scatter light. The Tyndall effect is the scattering of a light beam by dispersed particles, a useful clue for certain colloids. It is not a proof that any visibly cloudy sample is the same colloid. Milk is a familiar colloidal system; a saltwater solution is a true solution for the school comparison. The categories depend on particle behaviour, not solely the sample's colour.
Concentration statements need both numerator and denominator. A 10% mass-by-mass solution means 10 g solute per 100 g solution, not per 100 g solvent. If 10 g solute is added to 90 g solvent and the final solution mass is 100 g, the mass percentage is 10%. If instead 10 g solute is added to 100 g solvent, the total is 110 g and mass percentage is 10/110 × 100 ≈ 9.09%. These two cases are frequently confused because both mention “10 g” and “100 g.” Write the complete fraction before calculating.
For a dilution calculation in which solute amount is conserved, a compact relation is C₁V₁=C₂V₂ when concentration is measured in the same solute-per-volume units and volumes are treated appropriately. If 20 mL of a 10 g/L solution is diluted to 100 mL, the nominal final concentration is 2 g/L because the final volume is five times as large. This is a classroom arithmetic model; with some real liquid mixtures, volumes are not exactly additive, so the final marked volume is the relevant value. Never infer that dilution changed the quantity of solute unless the stem says solute was removed or reacted.
Separation follows a property difference
Filtration separates suitable insoluble solid particles from a liquid when the particles are retained by the filter. It will not remove dissolved sodium chloride from saltwater by ordinary paper filtration. Evaporation can recover a nonvolatile dissolved solid, though it does not recover the evaporated solvent. Distillation uses vaporisation and condensation to collect a component, making it relevant when solvent recovery or separation by boiling behaviour is required. A separating funnel can separate immiscible liquids of different densities after they form layers; it is not the method for two fully miscible liquids.
Chromatography separates components according to their different distributions between a moving phase and a stationary phase. In a simple paper setup, differently moving pigment spots illustrate the concept. A magnet can remove iron pieces from a suitable mixture because iron responds strongly to the magnet, while sand does not in the school example. Sieving uses particle-size difference. Centrifugation accelerates separation in appropriate mixtures by rotation, but it does not create a new substance. An exam item is easiest when you first identify the property used: solubility, boiling behaviour, immiscibility, magnetic response or particle size.
Crystallisation can help obtain a solid with a regular crystal structure from a suitable solution as conditions change. It may be preferred over simply heating to dryness if the solute is affected by strong heating or if purer crystals are the goal. One cannot assume all dissolved substances will crystallise by the same cooling step. In a question about mixed dissolved salts, more information about solubility and temperature dependence is needed. “Cooling always separates all salts” is too broad.
Compare acidity using ratios
In the usual dilute-solution approximation, pH=−log₁₀[H⁺], where the concentration is expressed relative to the standard concentration convention. A pH 3 solution has approximately ten times the hydrogen-ion concentration of a pH 4 solution and one hundred times that of pH 5. It is incorrect to describe pH 3 as “two times more acidic” than pH 5 by subtracting or dividing the pH numbers. The difference of two units represents a factor of 10². This comparison concerns the hydrogen-ion measure, not the total amount of acid in two containers of possibly different volumes.
For a simple strong-acid model in which the acid is fully ionised and water's own contribution is negligible, a tenfold dilution raises pH by approximately one unit. Thus a pH 2 solution becomes approximately pH 3 under that model. This rule cannot be extended indefinitely: as solutions become very dilute, water's contribution matters, and weak acids require their ionisation equilibrium to be considered. An exam stem that explicitly states the model allows the shortcut. Without those conditions, dilution still generally changes acidity but the exact pH change cannot be obtained from the volume ratio alone.
Neutrality means equal hydrogen-ion and hydroxide-ion activities in water, rather than a temperature-independent requirement that pH always equal 7. The familiar pH 7 reference applies near 25 °C. For the intended introductory questions, use the stated temperature and school approximation. A temperature-dependent numerical item needs the relevant water-ionisation information. This is a useful example of how a familiar classroom rule is valid within its conditions and can become misleading when stated without them.
Strength and amount answer different questions
Acid strength describes the tendency to donate protons under specified conditions; concentration describes how much acid is present per stated quantity of solution. Two hydrochloric-acid solutions can have very different concentrations while both contain the same strong acid. A weak acid can be present at a high concentration. The labels “strong” and “concentrated” therefore cannot be used interchangeably. A question about the extent of ionisation needs strength; a question about grams per litre needs concentration. The total acid amount also depends on how much solution is present.
Consider two hypothetical solutions: A contains 2 g solute per litre and B contains 5 g per litre. B is more concentrated on the stated basis. Yet 3 L of A contains 6 g solute, while 1 L of B contains 5 g. The more dilute solution can contain more total solute if enough of it is present. The same distinction applies when comparing acid solutions, salt solutions or sugar solutions. Write amount=concentration×volume and identify whether the question asks about amount or concentration before comparing the numbers.
Neutralisation arithmetic from a balanced equation
For HCl+NaOH→NaCl+H₂O, one mole of HCl reacts with one mole of NaOH in the simplified complete-neutralisation account. If 0.02 mol HCl meets 0.015 mol NaOH, the base is insufficient to neutralise all the acid. The reacted amount is 0.015 mol of each, leaving 0.005 mol HCl in excess. Equal volumes would not guarantee neutralisation if their concentrations differed. Nor can the remaining pH be calculated without the final volume and appropriate solution assumptions. The balanced equation tells the reaction ratio; it does not supply every missing measurement.
For H₂SO₄+2NaOH→Na₂SO₄+2H₂O, complete neutralisation uses two moles of NaOH per mole of sulfuric acid. If the problem provides 0.01 mol H₂SO₄, it needs 0.02 mol NaOH for this stated endpoint. Partial neutralisation can give a different salt composition, so the phrase “complete neutralisation” matters. A one-to-one mole shortcut borrowed from HCl would be wrong here. These are paper calculations for learning ratios and do not direct the handling or mixing of chemicals.
Concentration after dilution and mixing
Suppose a solution contains 8 g solute in a final volume of 200 mL. Its concentration is 8/0.2=40 g/L. If water is added until the final solution volume is 500 mL and no solute is lost or reacted, concentration becomes 8/0.5=16 g/L. The solute amount stays 8 g. Using “500 mL added” instead of “500 mL final volume” would describe a different problem. This small wording difference determines the denominator, so underline “to” or “by” when reading a dilution stem.
When mixing solutions of the same solute without reaction, total solute amount is the sum of the two amounts. For 100 mL at 10 g/L and 200 mL at 25 g/L, the amounts are 1 g and 5 g. Assuming a final volume of 300 mL, concentration is 6/0.3=20 g/L. The simple arithmetic mean of 10 and 25, or 17.5 g/L, is wrong because the volumes differ. The result lies between the starting concentrations and closer to that of the larger portion, providing a plausibility check. If final volume is specified separately, use that measured volume instead of assuming additivity.
Select a method that recovers the required component
The aim of a separation question matters as much as the mixture. Evaporating saltwater can recover salt while allowing water to escape; distillation can collect condensed water as well. Filtration retains suspended sand but allows dissolved salt to pass with the water. A separating funnel is suitable for immiscible liquid layers, whereas fractional distillation relies on differing volatility in suitable miscible-liquid mixtures. A question asking for the solvent and one asking for the solid can therefore need different methods even when they begin with the same mixture.
Chromatography can reveal that an apparently single coloured ink contains several components. Different spots arise because components interact differently with the stationary and mobile phases. Their separation does not require them to have visibly different particle sizes. A single spot is not absolute proof of purity because components can travel together under one set of conditions. For introductory questions, the observation supports a limited conclusion: several separated spots show several detectable components. Avoid extending the result beyond what the method demonstrates.
Molarity and molality: the denominator decides
Molarity is amount of solute in moles divided by final solution volume in litres; molality is amount of solute in moles divided by solvent mass in kilograms. Their similar names conceal two changes: volume versus mass, and solution versus solvent. Suppose 12 g of a solute with molar mass 60 g/mol is present in 400 mL of final solution. There are 12/60=0.20 mol, giving molarity 0.20/0.400=0.50 mol/L. If a separate preparation contains the same solute amount dissolved in 250 g of solvent, its molality is 0.20/0.250=0.80 mol/kg. The first problem does not give solvent mass; the second does not give final solution volume. One concentration cannot be converted into the other from those incomplete data alone.
Warming a closed solution can change its volume while leaving the amounts of solute and solvent unchanged. Molarity can consequently change through thermal expansion. Molality remains unchanged in this comparison because its denominator is solvent mass. This conclusion assumes no evaporation, reaction or loss of material. Mole fraction is another distinct ratio: amount of a component divided by total amount of all components. A mixture containing 1 mol of A and 4 mol of B has mole fractions 0.20 and 0.80, whose sum is one. Mole fraction has no unit. It is not the same as a mass fraction unless the necessary masses happen to produce that result.
Conjugate pairs, buffers and the water-ion relation
A Brønsted acid donates a proton; its conjugate base is what remains after that proton is removed. H₂CO₃ and HCO₃⁻ are one conjugate pair, while HCO₃⁻ and CO₃²⁻ are another. In each pair the formulas differ by one H⁺, including a one-unit change in charge. Bicarbonate can accept a proton to become carbonic acid or donate one to become carbonate. Being negatively charged does not prevent a species from acting as an acid in a suitable reaction. Check the actual proton transfer instead of assigning acid/base behaviour from the charge sign alone.
A buffer resists a large pH change when a limited amount of acid or base is added. A typical acidic buffer contains a weak acid and its conjugate base, such as ethanoic acid and ethanoate supplied by sodium ethanoate. Added H⁺ can be accepted by the conjugate base; added OH⁻ can be neutralised by the weak acid. An ammonia/ammonium-salt mixture provides a familiar basic-buffer example. A buffer does not keep pH exactly fixed under every addition, and it has a finite capacity. Large additions can exhaust a component. Dilution, concentration and component ratio also matter; the label “buffer” is not permission to ignore quantities.
At 25 °C, the introductory dilute-solution relation is K_w=[H⁺][OH⁻]≈1.0×10⁻¹⁴. With pH=−log₁₀[H⁺] and pOH=−log₁₀[OH⁻] under the same approximation, pH+pOH≈14. A solution with pOH 4 therefore has pH 10 at this temperature. Conversely, a fully dissociated strong acid supplying [H⁺]=10⁻² mol/L has pH about 2 when water's contribution is negligible. Do not use pH+pOH=14 at an unspecified different temperature without the appropriate K_w. The more precise treatment uses activities; the concentration formulas here are the school-level dilute approximation.
Gas pressure, vapour pressure and dissolved particles
Henry's law describes a proportional relation between a gas's equilibrium dissolved amount and its partial pressure above the liquid at fixed temperature within the applicable dilute range. If the pressure doubles under otherwise unchanged conditions, the dissolved concentration approximately doubles. Opening a carbonated drink lowers the carbon dioxide pressure above it, allowing excess dissolved gas to leave. The relevant pressure is the partial pressure of that gas, not an unrelated total pressure. Different textbook forms place the Henry constant on different sides of the relation; use the definition and units supplied in a numerical question.
Raoult's law concerns a component's vapour pressure above an ideal solution. In the simple non-volatile-solute case, solvent vapour pressure is p=x_solvent p°, where p° is the pure solvent's vapour pressure at the same temperature. A solvent mole fraction of 0.96 gives p=0.96p° and a relative lowering of 0.04, or 4%. Henry's law in a gas-solubility question and Raoult's law in a solvent-vapour-pressure question therefore address different comparisons. Real solutions can deviate from the ideal relation, so the model conditions belong with the answer.
Colligative properties depend principally on the number of dissolved particles relative to solvent in the dilute model. The standard group consists of relative lowering of vapour pressure, elevation of boiling point, depression of freezing point and osmotic pressure. Viscosity is not in this group. Dissociation matters: one dissolved formula unit of an electrolyte may give several ions, whereas a molecular solute may remain as single molecules. Association can reduce the effective particle count. Equal masses of different solutes therefore need not produce equal effects; compare particle amounts under the stated assumptions. Antifreeze illustrates freezing-point depression, without implying that any solute is suitable for an actual cooling system.
Osmosis is net solvent movement through a membrane that allows solvent through while restricting relevant solutes. For an introductory equal-pressure comparison, water tends to move toward the side with the higher effective concentration of non-penetrating solute. A red blood cell in a hypertonic surrounding solution loses water and shrinks; in a hypotonic environment it gains water and may rupture. “Hypertonic” compares the surrounding solution with the cell, so the reference matters. These are membrane-transport principles, not instructions for preparing fluids for medical use.
Hardness and colloid classifications
Hardness of water is associated mainly with dissolved calcium and magnesium ions, which interfere with soap lathering. Temporary hardness is commonly linked to hydrogen carbonates and can be reduced by boiling through reactions that remove the relevant ions from solution. Permanent hardness from common calcium or magnesium chlorides and sulfates is not removed simply by boiling. Sodium carbonate can remove hardness ions by forming poorly soluble carbonates; ion-exchange methods use a different mechanism. Clear appearance and ordinary filtration do not establish that dissolved hardness ions have been removed. Boiling for hardness and making water safe to drink are different questions.
For a colloid, identify both the dispersed phase and the surrounding dispersion medium. Fog is liquid droplets in gas; smoke commonly includes solid particles in gas; milk is a liquid-in-liquid emulsion; whipped cream is gas dispersed in liquid; and a gel such as jelly is commonly classified as liquid in a solid network. Fog and smoke are aerosols because their medium is gas. The familiar approximate colloidal size range is 1–1,000 nm; actual materials have distributions and more complex structures. Alum can help suspended particles aggregate so they settle more readily. This coagulation idea is different from claiming that ordinary filter paper directly removes every dissolved salt or every colloidal particle.
Retrieval exercise with worked checks
Calculate the mass percentage when 15 g solute dissolves in 135 g solvent with no loss. The solution mass is 150 g, giving 15/150×100=10%. Compare a pH 4 solution with pH 6 under the usual dilute approximation: the hydrogen-ion measure differs by a factor of 100. Decide whether paper filtration removes dissolved sugar from water: ordinarily it does not. Finally, explain why equal volumes of two unspecified acid and base solutions do not establish complete neutralisation: concentrations and the balanced chemical ratio are also needed. These questions test the denominator, logarithmic comparison, particle behaviour and reaction ratio separately.
Worked choices
Suppose a student must separate iron pins, sand and dissolved salt from water. A sequence of property-based choices could first remove the iron pins magnetically, then filter the sand from the salt solution, then evaporate water to recover the salt. If the task also requires collecting water, use distillation instead of simply evaporating it away. The order matters because filtration is useful for sand but not for the dissolved salt. This is an explanation of separation principles, not a laboratory protocol.
Now suppose a 5 g sample of salt is dissolved in 95 g of water, with no loss. The solution mass is 100 g, so mass-by-mass concentration is 5%. If a second question says “5 g per 100 mL of solution,” that is a mass-by-volume expression instead. The percentage signs look similar, but the denominator units differ. Convert only when enough density or volume information is supplied. A correct MCQ answer will retain the specified unit basis.
Recall and review
Explain the difference between acid strength and concentration. Read pH as a logarithmic aqueous measure without treating 0–14 as an absolute boundary. Balance HCl with NaOH and identify the salt without claiming every salt solution is neutral. Compare true solution, suspension and colloid using particle behaviour. Compute mass percentage with solution in the denominator. Finally, select a separation technique by stating which property differs between the components. Attempt the chapter practice set, then use the explanations to revisit any uncertain concepts.
Figure 9. Original chapter diagram.
Practice 43 questions
Choose one answer for each item. The answer key and explanations follow this chapter set.
1. Milk of magnesia, commonly taken as an antacid, is a suspension of:
A. Magnesium hydroxide, Mg(OH)₂ B. Magnesium sulphate, MgSO₄ C. Magnesium carbonate, MgCO₃ D. Magnesium chloride, MgCl₂
2. According to the Brønsted–Lowry concept, a base is a substance that:
A. Donates a proton (H⁺) B. Accepts an electron pair C. Produces H⁺ ions in water D. Accepts a proton (H⁺)
3. The concentration unit defined as the number of moles of solute per kilogram of solvent is called:
A. Molarity B. Normality C. Molality D. Mole fraction
4. Which of the following is NOT an aerosol, i.e. a colloid in which the dispersion medium is a gas?
A. Smoke B. Fog C. Paint D. Mist
5. In HCO₃⁻ + H⁺ → H₂CO₃, the bicarbonate ion accepts a proton. In HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O, it donates a proton. How does HCO₃⁻ behave in these two reactions?
A. As an acid in both reactions B. As a base first and an acid second C. As an acid first and a base second D. As a base in both reactions
6. An aqueous solution of which of the following salts is acidic?
A. Ammonium chloride (NH₄Cl) B. Sodium chloride (NaCl) C. Sodium carbonate (Na₂CO₃) D. Sodium acetate (CH₃COONa)
7. If the pH of a solution falls from 5 to 3, the hydrogen-ion concentration becomes:
A. 20 times greater B. 100 times smaller C. 100 times greater D. 2 times greater
8. Methyl orange indicator shows which colour in a basic solution?
A. Yellow B. Red C. Pink D. Colourless
9. Phenolphthalein indicator shows which colour in a basic solution?
A. Colourless B. Pink C. Yellow D. Blue
10. A solution that contains more dissolved solute than it can normally hold at that temperature, and which crystallises readily when disturbed, is called:
A. Saturated solution B. Supersaturated solution C. Unsaturated solution D. Colloidal solution
11. The usual reference range for arterial blood pH in a healthy adult is closest to which interval?
A. 6.35 – 6.45 B. 7.35 – 7.45 C. 7.85 – 7.95 D. 5.5 – 6.5
12. Assertion (A): Concentrated acetic acid is a strong acid.
Reason (R): The strength of an acid refers to the extent to which it ionises in water, not to its concentration.
A. Both A and R are true, and R is the correct explanation of A B. Both A and R are true, but R is not the correct explanation of A C. A is true, but R is false D. A is false, but R is true
13. What volume of 0.10 M NaOH is needed to neutralise 25 mL of 0.10 M HCl completely? Assume a 1:1 reaction.
A. 12.5 mL B. 50 mL C. 25 mL D. 100 mL
14. Which compound, commonly called washing soda, is used to remove permanent hardness of water?
A. Sodium carbonate (Na₂CO₃·10H₂O) B. Sodium bicarbonate (NaHCO₃) C. Sodium hydroxide (NaOH) D. Calcium hydroxide (Ca(OH)₂)
15. Which of the following mixtures acts as an acidic buffer solution?
A. Nitric acid + potassium nitrate B. Hydrochloric acid + sodium chloride C. Sodium hydroxide + sodium chloride D. Acetic acid + sodium acetate
16. Farmers commonly treat excessively acidic soil by adding:
A. Elemental sulphur B. Gypsum C. Ammonium sulphate D. Lime (quicklime or slaked lime)
17. Which dissolved salt can cause permanent hardness of water because it supplies magnesium ions with sulphate?
A. Sodium chloride (NaCl) B. Potassium nitrate (KNO₃) C. Magnesium sulphate (MgSO₄) D. Sodium bicarbonate (NaHCO₃)
18. Match these illustrative, approximate pH values for ordinary samples. List I: (a) lemon juice; (b) milk; (c) sea water; (d) milk of magnesia. List II: (i) about 10; (ii) about 6.5; (iii) about 2; (iv) about 8.
A. a-iii, b-iv, c-ii, d-i B. a-iii, b-ii, c-i, d-iv C. a-iii, b-ii, c-iv, d-i D. a-ii, b-iii, c-iv, d-i
19. What is the value of the ionic product of water, K_w = [H⁺][OH⁻], at 25 °C?
A. 1 × 10⁻⁷ B. 1 × 10⁻¹ C. 1 × 10⁻¹⁴ D. 1 × 10¹⁴
20. What is the pH of a 0.001 M solution of hydrochloric acid, assuming complete ionisation?
A. 1 B. 3 C. 11 D. 0.001
21. At 25 °C, what is the pH of 0.001 M NaOH if it dissociates completely and water autoionisation is neglected?
A. 3 B. 7 C. 10 D. 11
22. Temporary hardness of water can be removed simply by:
A. Adding common salt B. Boiling C. Filtration through sand D. Adding chlorine
23. The main buffer system that keeps the pH of human blood within its narrow normal range is:
A. Acetic acid–acetate (CH₃COOH/CH₃COO⁻) B. Carbonic acid–bicarbonate (H₂CO₃/HCO₃⁻) C. Ammonia–ammonium chloride (NH₃/NH₄Cl) D. Hydrochloric acid–sodium chloride (HCl/NaCl)
24. Which pair is a buffer made from a weak base and its conjugate acid?
A. HCl and NaCl B. NaOH and NaCl C. NH₃ and NH₄Cl D. HNO₃ and KNO₃
25. Consider the following statements:
Statement I: An aqueous solution of sodium acetate (CH₃COONa) is basic.
Statement II: An aqueous solution of potassium nitrate (KNO₃) is neutral.
Which of the above is/are correct?
A. Only I B. Only II C. Both I and II D. Neither I nor II
26. At 25 °C, a solution has pH 4. What is its pOH under the usual dilute-solution approximation?
A. 4 B. 7 C. 10 D. 14
27. Which pair differs by exactly one proton and is therefore a Brønsted conjugate acid–base pair?
A. HCl and NaCl B. H₂SO₄ and SO₄²⁻ C. NaOH and H₂O D. HSO₄⁻ and SO₄²⁻
28. Which of the following is NOT a strong acid in water?
A. Hydrochloric acid (HCl) B. Nitric acid (HNO₃) C. Perchloric acid (HClO₄) D. Hydrofluoric acid (HF)
29. Which of the following is NOT a colligative property?
A. Osmotic pressure B. Elevation of boiling point C. Depression of freezing point D. Viscosity
30. Which of the following is NOT used as an antacid to relieve excess stomach acid?
A. Sodium bicarbonate (NaHCO₃) B. Calcium carbonate (CaCO₃) C. Sodium hydroxide (NaOH) D. Magnesium hydroxide (Mg(OH)₂)
31. Arrange the following in increasing order of the size of their dispersed particles:
1. Muddy water (suspension)
2. Common salt solution
3. Milk
A. 3, 2, 1 B. 2, 3, 1 C. 2, 1, 3 D. 1, 3, 2
32. What happens to a red blood cell placed in a hypertonic salt solution?
A. It swells as water moves in B. It shrinks as water moves out C. It bursts immediately D. It remains unchanged
33. Alum is added to muddy water during purification because it causes the suspended colloidal clay particles to:
A. Peptise into a stable sol B. Dissolve completely C. Coagulate and settle down D. Evaporate
34. Which of the following concentration terms changes with a change in temperature?
A. Molarity B. Molality C. Mole fraction D. Mass percentage
35. Which compound is commonly added to the water in car radiators as an antifreeze in cold regions?
A. Ethylene glycol B. Acetone C. Benzene D. Ethyl acetate
36. At roughly constant temperature, soft drinks hold more dissolved CO₂ under higher CO₂ pressure and bubble when opened. Which gas-solubility law explains the pressure effect?
A. Raoult's law B. Boyle's law C. Charles's law D. Henry's law
37. At the same temperature and within the range where Henry’s law applies, what happens to the equilibrium dissolved concentration of a gas when its partial pressure above the liquid doubles?
A. It halves B. It remains unchanged C. It approximately doubles D. It becomes four times as large
38. For an ideal dilute solution of a non-volatile solute, which law relates the relative lowering of solvent vapour pressure to solute mole fraction?
A. Henry's law B. Raoult's law C. Graham's law D. Boyle's law
39. Match List-I (colloid) with List-II (dispersed phase in dispersion medium).
List-I: (a) Milk (b) Fog (c) Whipped cream (d) Jelly
List-II: (i) Gas in liquid (ii) Liquid in liquid (iii) Liquid in solid (iv) Liquid in gas
A. a-ii, b-iv, c-i, d-iii B. a-ii, b-i, c-iv, d-iii C. a-iii, b-iv, c-i, d-ii D. a-ii, b-iv, c-iii, d-i
40. The size of particles in a colloidal solution lies approximately in the range:
A. 0.01 nm to 0.1 nm B. 1 mm to 10 mm C. 1 nm to 1000 nm D. 0.1 mm to 1 mm
41. 20 g of common salt is dissolved in 80 g of water. What is the mass percentage of salt in the solution?
A. 25% B. 16% C. 80% D. 20%
42. 4 g of sodium hydroxide (molar mass 40 g/mol) is dissolved in water to make 500 mL of solution. What is the molarity of the solution?
A. 0.1 M B. 0.4 M C. 0.8 M D. 0.2 M
43. 5.85 g of sodium chloride (molar mass 58.5 g/mol) is dissolved in 100 g of water. What is the molality of the solution?
A. 1 m B. 0.1 m C. 0.5 m D. 10 m
Answers and explanations
1. A. Milk of magnesia is a suspension of magnesium hydroxide, a mild base that neutralises excess stomach acid: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O. Magnesium sulphate is Epsom salt, a laxative, and magnesium carbonate is the mineral magnesite.
2. D. In the Brønsted–Lowry theory an acid is a proton donor and a base is a proton acceptor, e.g. NH₃ + H⁺ → NH₄⁺. Accepting an electron pair defines a Lewis acid, and producing H⁺ in water is the Arrhenius definition of an acid.
3. C. Molality (m) = moles of solute ÷ kg of solvent. Because it is based on mass, it does not change with temperature, which makes it the unit used in colligative-property formulae. Molarity is moles per litre of solution and normality is equivalents per litre.
4. C. Smoke (solid particles in air), fog and mist (liquid droplets in air) are aerosols because their dispersion medium is a gas. Paint is a sol — solid pigment particles dispersed in a liquid medium.
5. B. A Brønsted base accepts a proton and a Brønsted acid donates one. Bicarbonate accepts H⁺ to form H₂CO₃ in the first reaction, so it acts as a base. It transfers a proton to OH⁻ in the second reaction, so it acts as an acid. A negatively charged species is not necessarily restricted to acting as a base.
6. A. NH₄Cl is the salt of a strong acid (HCl) and a weak base (NH₃); NH₄⁺ hydrolyses (NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺), making the solution acidic. NaCl (strong acid + strong base) is neutral, while Na₂CO₃ and sodium acetate (weak acid + strong base) are basic.
7. C. pH is logarithmic: each unit is a ten-fold change in [H⁺]. [H⁺] goes from 10⁻⁵ M to 10⁻³ M, an increase of 10² = 100 times. A lower pH means a more acidic solution.
8. A. Methyl orange is red in acidic solution and yellow in basic solution, changing over roughly pH 3.1–4.4. Pink in base is the behaviour of phenolphthalein, not methyl orange.
9. B. Phenolphthalein is colourless in acidic and neutral solutions and turns pink/magenta in basic solutions (about pH 8–10), which makes it the common indicator for strong acid–strong base titrations.
10. B. A supersaturated solution holds more solute than its equilibrium solubility and is unstable, so crystals form quickly when it is disturbed or seeded — honey crystallising over time is an example. A saturated solution holds exactly the maximum amount and an unsaturated one less.
11. B. A commonly reported arterial blood gas reference range is pH 7.35–7.45, slightly basic. The carbonic-acid/bicarbonate system contributes to acid–base regulation along with breathing and kidney function. Clinical interpretation depends on the person and laboratory; this item tests a reference interval, not a diagnosis.
12. D. A is false and R is true. Acetic acid is a weak acid at any concentration because only a small fraction of its molecules ionise (CH₃COOH ⇌ H⁺ + CH₃COO⁻). As R states, 'strong' and 'weak' describe degree of ionisation, whereas 'concentrated' and 'dilute' describe amount per volume.
13. C. HCl + NaOH → NaCl + H₂O is a 1:1 reaction. HCl amount = 0.10 mol/L × 0.025 L = 0.0025 mol, so NaOH needs 0.0025 mol. At 0.10 M, volume = 0.0025/0.10 = 0.025 L = 25 mL.
14. A. Washing soda is sodium carbonate decahydrate. Its carbonate ions precipitate Ca²⁺ and Mg²⁺ as insoluble carbonates (Ca²⁺ + CO₃²⁻ → CaCO₃↓), removing permanent hardness caused by chlorides and sulphates. NaHCO₃ is baking soda.
15. D. A buffer needs a weak acid and its conjugate base (or a weak base and its conjugate acid). CH₃COOH/CH₃COONa fits: acetate absorbs added H⁺ and acetic acid neutralises added OH⁻. Mixtures based on strong acids or bases cannot resist pH change.
16. D. Lime (CaO or Ca(OH)₂) or powdered limestone (CaCO₃) neutralises soil acidity. Sulphur and ammonium sulphate make soil more acidic, and gypsum is used to reclaim alkaline (sodic) soils, not acidic ones.
17. C. Magnesium sulphate supplies Mg²⁺ together with sulphate, a usual source of permanent hardness. Boiling does not remove this form as it can remove bicarbonate-derived temporary hardness. Water-softening methods, including carbonate precipitation or ion exchange, can remove the hardness ions.
18. C. A typical lemon-juice sample is strongly acidic, around pH 2; milk is mildly acidic, near 6.5; sea water is mildly basic, around 8; and a milk-of-magnesia suspension is basic, around 10. Thus a-iii, b-ii, c-iv, d-i. Real pH values vary with sample composition, temperature and measurement method, so the numbers are illustrative.
19. C. In pure water at 25 °C, [H⁺] = [OH⁻] = 10⁻⁷ M, so K_w = 10⁻⁷ × 10⁻⁷ = 1 × 10⁻¹⁴. This is why pH + pOH = 14 at 25 °C. 10⁻⁷ M is the concentration of each ion, not their product.
20. B. HCl is a strong acid, so [H⁺] = 0.001 M = 10⁻³ M. pH = −log[H⁺] = −log(10⁻³) = 3. The value 11 is the pOH (14 − 3), and 0.001 is the concentration itself, not the pH.
21. D. NaOH gives [OH⁻] = 10⁻³ M, so pOH = 3. At 25 °C, pH + pOH = 14, hence pH = 11. The pH is basic; 3 is the pOH, not the pH.
22. B. Temporary hardness is due to dissolved calcium and magnesium bicarbonates; boiling converts them into insoluble carbonates that precipitate, e.g. Ca(HCO₃)₂ → CaCO₃↓ + H₂O + CO₂. Permanent hardness is not removed by boiling.
23. B. Blood pH is held near 7.35–7.45 chiefly by the carbonic acid–bicarbonate buffer: HCO₃⁻ mops up extra H⁺, and H₂CO₃ releases H⁺ when pH rises. HCl with NaCl is not a buffer at all, since a strong acid has no useful conjugate base.
24. C. NH₃ is a weak base and NH₄⁺, supplied by NH₄Cl, is its conjugate acid. Together they can resist modest pH changes. HCl and HNO₃ are strong acids, while NaOH is a strong base; the listed combinations with their salts do not form the corresponding weak-base/conjugate-acid buffer.
25. C. Both are correct. Sodium acetate comes from a weak acid and a strong base; acetate ions hydrolyse (CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻), so the solution is basic. KNO₃ comes from a strong acid (HNO₃) and a strong base (KOH); neither ion hydrolyses, so pH ≈ 7.
26. C. At 25 °C, pH + pOH = 14 for the standard dilute aqueous approximation. Thus pOH = 14 − 4 = 10. The pH and pOH refer to different ion concentrations and are not generally equal.
27. D. HSO₄⁻ can donate one H⁺ to become SO₄²⁻, so the two species are a conjugate acid–base pair. H₂SO₄ to SO₄²⁻ differs by two protons; NaOH is a compound rather than the one-proton conjugate partner of water in the listed form.
28. D. HF is a weak acid in water because its ionisation is incomplete; HCl, HNO₃ and HClO₄ are commonly treated as strong aqueous acids. Do not infer acid strength from fluorine electronegativity alone. The H–F bond and solvent interactions matter. The question asks about aqueous ionisation, not how hazardous or corrosive a substance is.
29. D. Colligative properties depend only on the number of solute particles: relative lowering of vapour pressure, elevation of boiling point, depression of freezing point and osmotic pressure. Viscosity depends on the nature of the substances, so it is not colligative.
30. C. Antacids must be mild bases. Milk of magnesia (Mg(OH)₂), baking soda (NaHCO₃) and calcium carbonate neutralise stomach HCl. Sodium hydroxide is a strong, corrosive alkali (caustic soda) and would damage tissue.
31. B. True solution (salt, < 1 nm) < colloid (milk, 1–1000 nm) < suspension (muddy water, > 1000 nm). So the order is 2, 3, 1; suspension particles are large enough to settle on standing.
32. B. A hypertonic solution has a higher solute concentration than the cell, so water leaves the cell by osmosis and it shrinks (crenation). In a hypotonic solution water enters and the cell swells or bursts; in an isotonic one there is no net change.
33. C. Alum is a coagulant. Dissolved aluminium species hydrolyse, helping neutralise particle surface charges and forming aluminium-hydroxide flocs that capture fine clay. The enlarged aggregates can then settle or be filtered. The process is more than free Al³⁺ simply neutralising every clay particle.
34. A. Molarity uses the volume of solution, and volume expands or contracts with temperature, so molarity changes. Molality, mole fraction and mass percentage are based only on masses/moles, which do not change with temperature.
35. A. Ethylene glycol dissolved in water depresses its freezing point (a colligative effect), so the coolant does not freeze and crack the engine. Acetone, benzene and ethyl acetate are volatile solvents unsuitable as coolants.
36. D. Henry’s law relates the equilibrium concentration of a dissolved gas to its partial pressure above a sufficiently dilute solution at a fixed temperature. Opening a carbonated drink lowers the CO₂ partial pressure, so excess dissolved gas escapes as bubbles. Raoult’s law concerns an ideal-solution vapour-pressure relationship.
37. C. Henry’s law gives dissolved equilibrium concentration proportional to the gas’s partial pressure at fixed temperature, within its range of validity. Doubling partial pressure therefore approximately doubles the dissolved concentration. This is an application, rather than another definition, of the law.
38. B. For the stated ideal solution and non-volatile solute, Raoult’s law gives p = x_solvent p°; therefore (p° − p)/p° = x_solute. These equalities are model relations with conditions, not an exact rule for every real solution. Henry’s law instead describes dissolved-gas concentration versus partial pressure.
39. A. Milk is an emulsion (liquid fat in liquid water), fog is an aerosol (liquid droplets in gas), whipped cream is a foam (gas in liquid) and jelly is a gel (liquid in solid).
40. C. Colloidal particles are about 1–1000 nm; true-solution particles are below 1 nm and suspension particles are larger than about 1000 nm, which is why suspensions settle while colloids do not.
41. D. Mass of solution = 20 + 80 = 100 g. Mass % = (mass of solute ÷ mass of solution) × 100 = (20 ÷ 100) × 100 = 20%. Dividing by the mass of water alone (20/80) wrongly gives 25%.
42. D. Moles of NaOH = 4/40 = 0.1 mol. The final solution volume is 500 mL = 0.5 L, so molarity = 0.1/0.5 = 0.2 mol/L. Use the final solution volume, rather than the starting water volume; the distractors represent other arithmetic errors.
43. A. Moles of NaCl = 5.85/58.5 = 0.1 mol. Mass of solvent = 100 g = 0.1 kg. Molality = 0.1 mol ÷ 0.1 kg = 1 mol/kg (1 m). The answer 0.1 m forgets to convert grams of water to kilograms.