Grade 10 chemisty – Acids and bases Quiz

Bronsted-Lowry acids are defined as proton donors. In water they increase the concentration of H+ ions. This is the concept used in school chemistry and matches examples like hydrochloric acid (HCl).

Bronsted-Lowry bases accept protons. For example ammonia (NH3) accepts H+ to form NH4+ in water, so it acts as a base.

Pure water at 25 °C has equal concentrations of H+ and OH-, each 1 × 10^-7 mol L^-1, giving pH = -log(1 × 10^-7) = 7, which is classically defined as neutral.

Acidic solutions cause blue litmus paper to turn red because the indicator molecules change structure in presence of excess H+. This is a simple test used in school labs.

Phenolphthalein is colourless in acidic and neutral solutions and turns pink to fuchsia in basic solutions (usually above pH ≈ 8.2). This makes it useful in many titrations.

HCl is a strong acid because it ionises completely in water to give H+ and Cl-. Weak acids like acetic acid only partially ionise.

Acetic acid (vinegar) is a weak acid because it does not fully ionise in water. This is why vinegar has a mild acidic taste compared with strong acids like HCl.

Water can donate a proton to become OH- or accept a proton to become H3O+, so it is amphoteric. This behaviour is central to acid–base chemistry in aqueous solutions.

When HCl donates a proton (H+), the remaining species is Cl-, which is the conjugate base of HCl.

pH = -log[H+]. For [H+] = 1.0 × 10^-3, pH = -log(1.0 × 10^-3) = 3.

At 25 °C, pH + pOH = 14. If pH = 7, then pOH = 14 - 7 = 7.

Kw = [H+][OH-] for pure water at 25 °C equals 1.0 × 10^-14 (mol^2 L^-2). This constant is used to relate pH and pOH.

Neutralisation between an acid and a base usually produces a salt and water (for example HCl + NaOH → NaCl + H2O). Energy (heat) may be released but the chemical products are salt and water.

NaOH is a common laboratory base that produces OH- ions in solution and is used in titrations and tests for acids.

The equivalence point is when stoichiometric amounts have reacted: moles H+ from the acid equal moles OH- from the base (accounting for reaction stoichiometry), not necessarily equal volumes or concentrations.

HCO3- can donate a proton to become CO3^2- or accept a proton to become H2CO3, so it behaves as an amphiprotic species (important in buffering blood and natural waters).

Ammonia acts as a Bronsted-Lowry base because it can accept a proton to form NH4+. HCl and ethanoic acid donate protons, so they are acids.

BF3 has an incomplete octet on boron and can accept an electron pair, making it a Lewis acid. NH3 and Cl- are electron-pair donors (Lewis bases).

NaOH is a strong base and fully dissociates to give [OH-] = 0.01 M. pOH = -log[OH-] = 2, so pH = 14 - 2 = 12.

Red litmus paper turns blue in basic (alkaline) solutions because the indicator responds to lower H+ and higher OH- concentrations.

pH is a logarithmic scale: each unit change is a tenfold change in [H+]. A change from pH 7 to pH 4 is 3 units, so [H+] increases by 10^3 = 1000 times.

H2SO4 is diprotic: it can donate two protons in separate steps (first to HSO4-, then to SO4^2-). HCl and acetic acid are monoprotic (donate one proton).

HCl + NaOH → NaCl + H2O. The salt formed from HCl and NaOH is sodium chloride, common table salt.

Metals above hydrogen in the reactivity series react with acids to produce a salt (e.g., ZnCl2) and hydrogen gas: Zn + 2HCl → ZnCl2 + H2.

The acid dissociation constant Ka measures acid strength; larger Ka means stronger acid (more ionisation). 1.0 × 10^-2 is greater than 1.0 × 10^-5, so A is stronger.

Brønsted-Lowry defines an acid as a proton (H+) donor. The other choices describe bases or incorrect properties.

A Brønsted-Lowry base accepts protons (H+). The other options describe different concepts or are incorrect.

At 25°C pure water and a neutral solution have pH 7. Lower values are acidic and higher values are basic.

Neutralisation of an acid with a base produces a salt and water (for example HCl + NaOH -> NaCl + H2O).

Blue litmus turns red in acidic solutions; in a base it stays blue.

A strong acid fully ionises in water, producing H+ ions; the other statements are false.

HCl is a strong acid so [H+] = 0.01 M. pH = -log[H+] = -log(0.01) = 2.

At 25°C the ionic product of water gives pH + pOH = 14. This relation is used to switch between pH and pOH.

Water can donate or accept a proton, so it is amphoteric. NaOH is a base, HCl is an acid, NaCl is a neutral salt.

When ammonia accepts a proton it becomes the ammonium ion NH4+, which is its conjugate acid.

When HCl donates a proton it forms the chloride ion Cl-, its conjugate base.

Phenolphthalein is colourless in acidic solutions and becomes pink in alkaline solutions.

Universal indicator is green around pH 7, red in acids and blue/purple in strong bases.

Kw at 25°C equals [H+][OH-] = 1.0 × 10^-14 mol^2 dm^-6; the other numbers are incorrect for Kw.

NH4+ is the conjugate acid of a weak base (NH3), so NH4Cl solution is slightly acidic. The others are salts of strong acid/strong base and are neutral.

CO3^2- is the conjugate base of the weak acid HCO3-, so sodium carbonate solution is basic.

Metals such as zinc react with acids like HCl to produce a salt and hydrogen gas (Zn + 2HCl → ZnCl2 + H2).

Ka is defined as the equilibrium concentration of products over reactants for the acid dissociation: Ka = [H+][A-]/[HA].

A buffer contains a weak acid and its conjugate base (or a weak base and its conjugate acid) and helps maintain a nearly constant pH.

A solution of acetic acid and its salt sodium acetate forms a buffer that resists pH change; the other pairs are not buffers.

A Lewis acid accepts an electron pair. This is a broader definition than Brønsted-Lowry which focuses on protons.

Adding acid to water controls heat released and prevents splashing; adding water to concentrated acid can cause violent splatter. Always use eye protection.

For a strong acid titrated with a strong base the equivalence point is at pH ~7 because the salt formed is neutral.

Phenolphthalein changes colour in the pH range around 8–10 and is commonly used for strong acid vs strong base titrations; methyl red has a lower transition range.

Strength refers to degree of ionisation. At the same concentration a weak acid yields fewer H+ ions in solution than a strong acid.

Sodium hydroxide (caustic soda) is a common strong base found in drain cleaners and used in school labs; the others are not bases or are acids.

Arrhenius acids are defined as substances that, when dissolved in water, increase [H+]. This is the classic school definition used for aqueous solutions.

Brønsted-Lowry defines an acid as a proton donor and a base as a proton acceptor; this applies in many solvents, not only water.

Neutralisation between an acid (HCl) and a base (NaOH) produces a salt (NaCl) and water; this is the typical neutralisation equation.

HCl is a strong acid and fully ionises: [H+] = 0.01 M. pH = −log[H+] = −log(0.01) = 2.

Acid strength refers to how completely it donates protons (ionises); concentration refers to the molarity (amount per litre). A dilute strong acid may have less H+ than a concentrated weak acid.

Phenolphthalein is colourless in acidic and neutral solutions and turns pink in alkaline solutions (roughly above pH 8.2).

At 25°C, Kw = [H+][OH-] = 1.0 × 10^-14. This constant is used to relate [H+] and [OH-] in pure water and aqueous solutions.

pH = −log[H+] = 5. Since pH + pOH = 14 at 25°C, pOH = 14 − 5 = 9.

NaCl comes from a strong acid (HCl) and a strong base (NaOH); neither ion hydrolyses significantly, so the solution is neutral (pH ≈ 7).

Aluminium hydroxide reacts with acids (acting as a base) and with strong bases (acting as an acid), so it is amphoteric.

When H2SO4 donates one proton (H+), it becomes HSO4^-; this remaining species is the conjugate base of the acid.

Lemon juice contains citric acid and is acidic. Baking soda and washing soda are basic; bleach is also basic/oxidising.

At the equivalence point the weak acid has been converted to its conjugate base, which hydrolyses to produce OH-, making the solution basic (pH > 7).

Buffers are typically a weak acid and its salt (conjugate base) or a weak base and its salt; they resist changes in pH when small amounts of acid or base are added.

Red litmus turns blue in alkaline solutions; litmus is a simple indicator commonly used in school practicals.

A strong base such as NaOH fully dissociates in water to give OH- and the corresponding cation, increasing [OH-] significantly.

Acetate ion (CH3COO-) hydrolyses with water to produce OH-, making the solution slightly basic.

A Lewis acid accepts an electron pair and a Lewis base donates an electron pair. BF3 accepts a lone pair from NH3 to form a coordinate covalent bond.

Ionisation of water is endothermic, so increasing temperature increases Kw (more H+ and OH- are produced). Kw = 1.0×10^-14 only at 25°C.

NH4+ is the conjugate acid of the weak base NH3; NH4+ hydrolyses in water producing H+ and making the solution acidic.

HCl is a strong acid and fully ionises: [H+] = 0.001 M = 10^-3, so pH = −log(10^-3) = 3.

Weak acids only partially ionise in water, so at the same concentration they produce fewer H+ ions than strong acids, which ionise completely.

A strong acid–strong base titration has an equivalence point at pH ≈ 7, so an indicator with a colour change near pH 7 gives the most accurate end point.

Diluting a strong acid reduces [H+], so pH (which is −log[H+]) increases; the solution becomes less acidic.

NH3 can accept a proton to become NH4+; they differ by one proton and are therefore a conjugate base–conjugate acid pair.

CO2 from the atmosphere dissolves in rainwater to form carbonic acid (H2CO3), making natural rain slightly acidic (pH about 5.6). Pollutants make it lower.