The characteristic properties of acids and bases Revision Notes

When acids react with certain metals, they produce a salt and hydrogen gas. This reaction occurs because acids contain hydrogen ions (H⁺) that are displaced by the metal.

Not all metals react with acids; only metals above hydrogen in the reactivity series typically do so, such as magnesium, zinc, and iron.

The general word equation is: Acid + Metal → Salt + Hydrogen gas.

• Hydrogen gas is released, often seen as bubbles or effervescence.
• The salt formed depends on the acid and metal used.
• Example: Hydrochloric acid reacts with zinc to form zinc chloride and hydrogen gas.

Balanced symbol equations show the conservation of atoms and charge during the reaction.

Example: Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g) (zinc + hydrochloric acid → zinc chloride + hydrogen gas)

Another example: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

• State symbols: (s) solid, (aq) aqueous solution, (g) gas.
• Hydrogen gas can be tested by collecting it and applying a lit splint, which produces a 'pop' sound due to combustion.

Acids react with bases to form a salt and water in a neutralisation reaction.

Bases are substances that can accept hydrogen ions or contain hydroxide ions (OH⁻).

The general word equation is: Acid + Base → Salt + Water.

• Example: Hydrochloric acid reacts with sodium hydroxide to form sodium chloride and water.
• This reaction reduces the acidity of the solution, producing a neutral or less acidic solution.
• Neutralisation is exothermic, releasing heat.

Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) (hydrochloric acid + sodium hydroxide → sodium chloride + water)

Ionic equation: H⁺(aq) + OH⁻(aq) → H₂O(l) shows the key reaction between hydrogen and hydroxide ions.

• Water is formed by the combination of H⁺ and OH⁻ ions.
• Salts formed depend on the acid and base used.
• The solution's pH moves towards neutral (pH 7) after neutralisation.

Acids react with carbonates to produce a salt, carbon dioxide gas, and water.

Carbonates contain the carbonate ion (CO₃²⁻). The reaction releases carbon dioxide gas, which can be tested by bubbling through limewater, turning it milky.

The general word equation is: Acid + Carbonate → Salt + Carbon dioxide + Water.

• Example: Hydrochloric acid reacts with calcium carbonate to form calcium chloride, carbon dioxide, and water.
• Effervescence (bubbling) is observed due to carbon dioxide gas.
• The reaction is commonly used to identify carbonates.

Example: 2HCl(aq) + CaCO₃(s) → CaCl₂(aq) + CO₂(g) + H₂O(l) (hydrochloric acid + calcium carbonate → calcium chloride + carbon dioxide + water)

Another example: 2HNO₃(aq) + Na₂CO₃(s) → 2NaNO₃(aq) + CO₂(g) + H₂O(l)

• State symbols indicate the physical states of reactants and products.
• Carbon dioxide gas can be confirmed by its reaction with limewater.

Indicators are substances that change colour depending on the pH of the solution they are in. They are used to identify whether a solution is acidic, neutral, or alkaline.

Acids are substances that release hydrogen ions (H⁺) when dissolved in water, making the solution acidic. The presence of these hydrogen ions causes specific colour changes in indicators.

Common indicators used to test for acidity include litmus, thymolphthalein, and methyl orange. Each has a distinct colour change in acidic solutions.

• Indicators help detect acidity by changing colour.
• Acids release H⁺ ions in aqueous solution.
• Litmus, thymolphthalein, and methyl orange are common acid-base indicators.

Litmus is a natural dye extracted from lichens and is commonly used as a simple acid-base indicator.

In acidic solutions, litmus paper turns red. This is because the hydrogen ions in the acid react with the dye molecules, causing a change in their electronic structure and thus their colour.

In neutral solutions, litmus remains purple, and in alkaline solutions, it turns blue. Therefore, a red colour on litmus paper indicates an acidic environment.

• Litmus paper turns red in acids.
• Litmus is purple in neutral solutions.
• Litmus turns blue in alkaline solutions.

Thymolphthalein is a synthetic indicator used to detect alkaline and acidic conditions.

In acidic solutions, thymolphthalein remains colourless because the indicator molecules are in their protonated form, which does not absorb visible light strongly.

In alkaline solutions, thymolphthalein turns blue, indicating the presence of hydroxide ions (OH⁻). Thus, the absence of colour (colourless) indicates acidity.

• Thymolphthalein is colourless in acids.
• It turns blue in alkaline solutions.
• Colour change depends on protonation state of the indicator molecule.

Methyl orange is a synthetic indicator that changes colour over a pH range of approximately 3.1 to 4.4, making it useful for detecting acidic conditions.

In acidic solutions, methyl orange turns red due to the protonation of the indicator molecule, altering its electronic structure and colour.

In neutral and alkaline solutions, methyl orange appears yellow. Therefore, a red colour indicates an acidic solution.

• Methyl orange turns red in acids.
• It is yellow in neutral and alkaline solutions.
• Colour change occurs due to protonation of the indicator molecule.

The following table summarises the colour changes of the three indicators when exposed to acidic solutions:

• Litmus: red in acid
• Thymolphthalein: colourless in acid
• Methyl orange: red in acid

Bases are substances that can neutralise acids to form salts and water. Chemically, bases are often metal oxides or metal hydroxides.

Metal oxides are compounds composed of metal cations combined with oxide ions (O²⁻). For example, copper(II) oxide (CuO) and magnesium oxide (MgO) are metal oxides.

Metal hydroxides contain metal cations combined with hydroxide ions (OH⁻). Examples include sodium hydroxide (NaOH) and calcium hydroxide (Ca(OH)₂).

• Bases react with acids in neutralisation reactions.
• Common bases include metal oxides and metal hydroxides.
• Metal oxides and hydroxides are ionic compounds containing O²⁻ or OH⁻ ions.

Alkalis are a special group of bases that dissolve in water to produce hydroxide ions (OH⁻) in solution, making the solution alkaline.

All alkalis are bases, but not all bases are alkalis. The key difference is solubility in water.

For example, sodium hydroxide (NaOH) is an alkali because it dissolves in water to give OH⁻ ions, whereas copper(II) oxide (CuO) is a base but not an alkali because it is insoluble in water.

• Alkalis are bases soluble in water.
• Alkalis produce hydroxide ions (OH⁻) in aqueous solution.
• Solubility distinguishes alkalis from other bases.

When a base reacts with an acid, a neutralisation reaction occurs, producing a salt and water.

For example, the reaction of sodium hydroxide (an alkali) with hydrochloric acid is:

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l) (balanced equation)

• Alkalis dissociate in water: NaOH(s) → Na⁺(aq) + OH⁻(aq)
• Bases that are insoluble do not dissociate significantly in water.
• Neutralisation involves H⁺ ions from acid reacting with OH⁻ ions from alkali to form water.

Common bases include: copper(II) oxide (CuO), magnesium oxide (MgO), calcium hydroxide (Ca(OH)₂), and aluminium hydroxide (Al(OH)₃).

Common alkalis include: sodium hydroxide (NaOH), potassium hydroxide (KOH), and aqueous ammonia (NH₃ solution, which produces OH⁻ ions by reaction with water).

Solubility is the key factor: Ca(OH)₂ is only sparingly soluble and is sometimes called a weak alkali.

• Metal oxides like CuO and MgO are bases but insoluble.
• Metal hydroxides like NaOH and KOH are alkalis due to their solubility.
• Alkalis release OH⁻ ions in solution, increasing pH.

To identify whether a substance is a base or an alkali, test its solubility in water.

If it dissolves and produces OH⁻ ions (detected by blue litmus paper turning blue or pH > 7), it is an alkali.

If it does not dissolve appreciably, it is a base but not an alkali.

• Test solubility in water to distinguish alkalis from insoluble bases.
• Use indicators or pH measurement to confirm presence of OH⁻ ions.
• Recognise that all alkalis are bases, but not all bases are alkalis.

Bases are substances that can accept protons (H⁺ ions) or donate electron pairs. Chemically, they neutralise acids by reacting with hydrogen ions to form water and a salt. Many bases are metal oxides or hydroxides. When soluble in water, these bases are called alkalis and produce hydroxide ions (OH⁻) in solution, making the solution alkaline.

Bases typically have a bitter taste and a slippery feel (though these properties are not tested in the laboratory for safety reasons). They change the colour of acid-base indicators: for example, bases turn red litmus paper blue and methyl orange indicator yellow.

• Bases accept protons (H⁺ ions) or donate electron pairs.
• They neutralise acids to form salts and water.
• Soluble bases are called alkalis and produce OH⁻ ions in solution.
• Bases change indicator colours: red litmus to blue, methyl orange to yellow.

When a base reacts with an acid, a neutralisation reaction occurs, producing a salt and water. This reaction involves the base accepting H⁺ ions from the acid, forming water molecules.

For example, sodium hydroxide (a base) reacts with hydrochloric acid as follows:

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l) (balanced equation)

• Neutralisation: acid + base → salt + water.
• Example: NaOH + HCl → NaCl + H₂O.
• This reaction removes H⁺ ions, reducing acidity.

Bases react with ammonium salts to release ammonia gas (NH₃). This occurs because the base removes H⁺ ions from the ammonium ion (NH₄⁺), forming ammonia.

For example, when sodium hydroxide reacts with ammonium chloride:

NaOH(aq) + NH₄Cl(aq) → NaCl(aq) + NH₃(g) + H₂O(l) (balanced equation)

• Bases react with ammonium salts to release NH₃ gas.
• NH₄⁺ ions lose H⁺ ions to the base, forming NH₃.
• Ammonia gas is alkaline and has a characteristic pungent smell.
• This reaction is used to test for ammonium ions.

Soluble bases (alkalis) dissociate in water to produce hydroxide ions (OH⁻). These ions increase the pH of the solution, making it alkaline.

For example, sodium hydroxide dissociates as: NaOH(s) → Na⁺(aq) + OH⁻(aq)

The presence of OH⁻ ions causes indicator colour changes: red litmus paper turns blue, and methyl orange turns yellow.

• Alkalis produce OH⁻ ions in aqueous solution.
• OH⁻ ions cause alkaline pH (>7).
• Indicators change colour in presence of OH⁻ ions (red litmus → blue).

Bases include metal oxides and hydroxides that may be insoluble or soluble in water. Alkalis are a subset of bases that dissolve in water to produce OH⁻ ions.

For example, copper(II) oxide (CuO) is a base but insoluble in water, so it does not form an alkaline solution. Sodium hydroxide (NaOH) is an alkali because it dissolves and produces OH⁻ ions.

This distinction is important when predicting reactions and pH changes.

• Bases may be soluble or insoluble in water.
• Alkalis are soluble bases producing OH⁻ ions in solution.
• Not all bases are alkalis (e.g., metal oxides like CuO).

Alkalis are a subset of bases that dissolve in water to produce hydroxide ions (OH⁻). These hydroxide ions increase the pH of the solution, making it alkaline (pH greater than 7). Alkalis can be metal hydroxides such as sodium hydroxide (NaOH) or potassium hydroxide (KOH). Their aqueous solutions are slippery to the touch and conduct electricity due to the presence of ions.

• Alkalis are soluble bases.
• They release hydroxide ions (OH⁻) in aqueous solution.
• Alkalis have pH values above 7, typically between 8 and 14.

Litmus is a natural dye used as an acid-base indicator. It is commonly available as red or blue litmus paper. When exposed to an alkaline solution, red litmus paper turns blue, indicating the presence of alkali.

This colour change occurs because the alkali causes the litmus dye molecules to accept hydroxide ions, altering their electronic structure and thus their colour.

• Red litmus paper turns blue in alkalis.
• Blue litmus paper remains blue in alkalis.
• Litmus changes colour around pH 7, making it a simple indicator for acidity or alkalinity.

Thymolphthalein is a synthetic indicator that is colourless in acidic and neutral solutions but turns blue in alkaline solutions. It changes colour at a higher pH range than litmus, typically between pH 9.3 and 10.5.

This makes thymolphthalein useful for detecting strongly alkaline solutions. The colour change is due to the deprotonation of the indicator molecule in the presence of hydroxide ions.

• Thymolphthalein is colourless below pH 9.3.
• It turns blue in solutions with pH above 9.3, indicating strong alkalinity.
• Used to detect strong alkalis such as concentrated sodium hydroxide solutions.

Methyl orange is an indicator that changes colour in acidic to neutral pH ranges. In acidic solutions, it appears red, and as the solution becomes less acidic and more alkaline, it changes to yellow.

The transition occurs around pH 3.1 to 4.4. In alkaline solutions (pH above 4.4), methyl orange is yellow. This indicator is therefore useful for distinguishing acidic from neutral or alkaline solutions, but it does not change colour further in strongly alkaline conditions.

• Methyl orange is red in acidic solutions (pH < 3.1).
• It changes to orange and then yellow as pH increases.
• In alkalis (pH > 4.4), methyl orange is yellow.

The three indicators respond differently to alkalis due to their distinct pH transition ranges. Litmus changes colour at neutral pH, thymolphthalein at higher alkaline pH, and methyl orange at acidic to neutral pH.

These differences allow chemists to select appropriate indicators depending on the expected pH range of the solution being tested.

• Litmus: red to blue at pH 7 (alkalis turn litmus blue).
• Thymolphthalein: colourless to blue at pH 9.3–10.5 (alkalis turn it blue).
• Methyl orange: red to yellow at pH 3.1–4.4 (alkalis turn it yellow).

When acids dissolve in water, they release hydrogen ions (H⁺) into the solution. These ions are responsible for the acidic properties of the solution.

Alkalis are soluble bases that release hydroxide ions (OH⁻) when dissolved in water. The presence of these ions gives the solution alkaline properties.

Understanding which ions are present in solution is essential for predicting and explaining the reactions and properties of acids and alkalis.

• Acidic solutions contain H⁺ ions.
• Alkaline solutions contain OH⁻ ions.

Acids ionise in water to produce hydrogen ions. For example, hydrochloric acid (HCl) dissociates completely in water:

HCl (aq) → H⁺ (aq) + Cl⁻ (aq) (complete dissociation)

The hydrogen ion is often represented as H⁺, but in aqueous solution it exists as the hydronium ion (H₃O⁺) formed by H⁺ associating with water molecules. However, for simplicity, H⁺ is used in equations and explanations at this level.

• H⁺ ions cause the acidic properties of solutions.
• Strong acids fully dissociate to release H⁺ ions.
• Weak acids partially dissociate, releasing fewer H⁺ ions.

Alkalis are bases that dissolve in water to produce hydroxide ions (OH⁻). For example, sodium hydroxide (NaOH) dissociates completely:

NaOH (aq) → Na⁺ (aq) + OH⁻ (aq) (complete dissociation)

The OH⁻ ions are responsible for the alkaline properties of the solution, such as a slippery feel and ability to neutralise acids.

• OH⁻ ions cause alkaline properties.
• Alkalis are soluble bases producing OH⁻ ions in solution.
• Examples include NaOH and KOH.

The presence of H⁺ or OH⁻ ions can be identified using indicators such as litmus paper, which turns red in acidic solutions and blue in alkaline solutions.

Ionic equations focus on the ions involved in reactions. For example, neutralisation between an acid and an alkali involves H⁺ and OH⁻ ions:

H⁺ (aq) + OH⁻ (aq) → H₂O (l) (formation of water)

• Indicators detect H⁺ or OH⁻ ions by colour change.
• Ionic equations show the key ions in reactions.
• Neutralisation involves H⁺ and OH⁻ forming water.

A solution is tested with blue litmus paper, which turns red. This indicates the solution contains H⁺ ions and is acidic.

If red litmus paper turns blue, the solution contains OH⁻ ions and is alkaline.

If there is no colour change, the solution is neutral and contains neither excess H⁺ nor OH⁻ ions.

• Blue litmus turning red indicates H⁺ ions (acid).
• Red litmus turning blue indicates OH⁻ ions (alkali).
• No change indicates neutrality.

The acidity or alkalinity of an aqueous solution depends on the concentration of hydrogen ions (H⁺) present. Acids increase the concentration of H⁺ ions in solution, while alkalis increase the concentration of hydroxide ions (OH⁻). The pH scale is a numerical scale from 0 to 14 that quantifies the hydrogen ion concentration:

pH = -log[H⁺] where [H⁺] is the molar concentration of hydrogen ions.

A low pH corresponds to a high hydrogen ion concentration (acidic solution). A high pH corresponds to a low hydrogen ion concentration (alkaline solution). A pH of 7 indicates neutrality, where [H⁺] = [OH⁻].

• pH scale ranges from 0 to 14
• pH < 7 indicates acidic solution (high [H⁺])
• pH = 7 indicates neutral solution ([H⁺] = [OH⁻])
• pH > 7 indicates alkaline solution (low [H⁺])

Universal indicator paper is a strip of paper impregnated with a mixture of indicators that change colour over a wide pH range. When dipped into a solution, the paper changes colour according to the solution's pH. This allows a quick, visual comparison of acidity or alkalinity.

To use universal indicator paper: dip a clean strip into the solution, remove it, and immediately compare the resulting colour to a standard pH colour chart provided with the indicator. This gives an approximate pH value and indicates whether the solution is acidic, neutral, or alkaline.

• Universal indicator paper shows a range of colours from red to violet
• Colour corresponds to pH value
• Used for quick, approximate pH measurement

The universal indicator colour changes correspond to specific pH values and solution types:

pH 1–3: Red to orange — strongly acidic (high [H⁺])

pH 4–6: Yellow to green — weakly acidic to neutral transition (moderate [H⁺])

• pH 1–3: Red to orange (strong acid)
• pH 4–6: Yellow to green (weak acid to near neutral)
• pH 7: Green (neutral)
• pH 8–11: Blue (weak alkali)
• pH 12–14: Violet (strong alkali)

By observing the colour of universal indicator paper after dipping into different solutions, you can compare their hydrogen ion concentrations:

A solution producing a red or orange colour has a higher hydrogen ion concentration than one producing yellow or green.

A green colour indicates neutrality, where hydrogen ion concentration equals hydroxide ion concentration (approximately 1 × 10⁻⁷ mol/dm³).

• More acidic solutions produce colours towards red
• More alkaline solutions produce colours towards violet
• Neutral solutions produce green colour

A solution turns universal indicator paper orange. Using the colour chart, orange corresponds to pH 3. Calculate the hydrogen ion concentration:

Formula: pH = -log[H⁺] → [H⁺] = 10^(-pH) Substitution: [H⁺] = 10^(-3) mol/dm³ Answer: [H⁺] = 1 × 10⁻³ mol/dm³ This indicates a strongly acidic solution with a high concentration of hydrogen ions.

• Use pH = -log[H⁺] to calculate hydrogen ion concentration
• Example: pH 3 → [H⁺] = 1 × 10⁻³ mol/dm³

Neutralization is a chemical reaction in which an acid reacts with an alkali to produce water and a salt. This reaction involves the combination of hydrogen ions (H⁺) from the acid and hydroxide ions (OH⁻) from the alkali to form water (H₂O). The general word equation is: Acid + Alkali → Salt + Water.

An acid is a substance that releases hydrogen ions (H⁺) in aqueous solution, while an alkali is a soluble base that releases hydroxide ions (OH⁻) in aqueous solution. The reaction between these ions is the basis of neutralization.

• Neutralization involves proton transfer from acid to alkali.
• Water is formed by the combination of H⁺ and OH⁻ ions.
• A salt is also produced, depending on the acid and alkali used.

In aqueous solution, acids dissociate to release H⁺ ions, and alkalis dissociate to release OH⁻ ions. For example, hydrochloric acid (HCl) dissociates as: HCl (aq) → H⁺ (aq) + Cl⁻ (aq). Sodium hydroxide (NaOH) dissociates as: NaOH (aq) → Na⁺ (aq) + OH⁻ (aq).

The neutralization reaction occurs between H⁺ and OH⁻ ions: H⁺ (aq) + OH⁻ (aq) → H₂O (l). This ionic equation shows the essential chemical change: formation of water from hydrogen and hydroxide ions.

• Acids provide H⁺ ions in solution.
• Alkalis provide OH⁻ ions in solution.
• H⁺ and OH⁻ ions combine to form water molecules.

The general balanced chemical equation for neutralization is: Acid + Alkali → Salt + Water.

For example, when hydrochloric acid reacts with sodium hydroxide: HCl (aq) + NaOH (aq) → NaCl (aq) + H₂O (l).

Another example is the reaction between sulfuric acid and potassium hydroxide: H₂SO₄ (aq) + 2KOH (aq) → K₂SO₄ (aq) + 2H₂O (l). This shows that the number of hydroxide ions must balance the number of hydrogen ions from the acid for complete neutralization.

• Equations must be balanced for atoms and charge.
• Water is always a product formed by H⁺ and OH⁻ combining.
• The salt formed depends on the acid’s anion and the alkali’s cation.

During neutralization, hydrogen ions (H⁺) from the acid react with hydroxide ions (OH⁻) from the alkali. These ions combine through an acid-base reaction to form water molecules (H₂O).

This process can be viewed as proton transfer: the acid donates a proton (H⁺) which is accepted by the hydroxide ion (OH⁻), resulting in water. This is why neutralization is often described as an acid-base reaction.

• Proton transfer is the key step in neutralization.
• Water molecules form from H⁺ and OH⁻ ions.
• This reaction removes the acidic and alkaline properties, producing a neutral solution if balanced.

Neutralization reactions typically occur in aqueous solution at room temperature. They are usually exothermic, releasing heat, which can be observed as a temperature rise.

Indicators such as litmus or methyl orange change colour during neutralization, signalling the reaction’s progress.

The reaction requires the acid and alkali to be in solution so ions can move freely and react.

• Neutralization is exothermic; temperature increases.
• Colour changes in indicators confirm neutralization.
• Reaction occurs in aqueous solution where ions are free to move.

In chemistry, acids and bases are classified by their ability to transfer protons (H⁺ ions) during chemical reactions. This proton transfer concept is central to the Brønsted-Lowry theory, which defines acids as substances that donate protons and bases as substances that accept protons.

Understanding this definition helps explain many chemical reactions involving acids and bases, including neutralisation and reactions in aqueous solutions.

• Acids release H⁺ ions (protons) to other substances.
• Bases accept H⁺ ions from acids or other proton donors.

An acid is any substance that can donate a proton (H⁺ ion) to another species. When an acid dissolves in water, it releases H⁺ ions into the solution, increasing the concentration of hydrogen ions.

For example, hydrochloric acid (HCl) dissociates completely in water:

HCl (aq) → H⁺ (aq) + Cl⁻ (aq) (complete dissociation)

• Acids increase the concentration of H⁺ ions in aqueous solution.
• The proton donated is a hydrogen ion (H⁺), which is essentially a bare proton.
• Strong acids fully dissociate, weak acids partially dissociate in water.

A base is any substance that can accept a proton (H⁺ ion) from another species. Bases often contain lone pairs of electrons that can bond with the proton.

For example, ammonia (NH₃) acts as a base by accepting a proton from water:

NH₃ (aq) + H₂O (l) ⇌ NH₄⁺ (aq) + OH⁻ (aq) (proton transfer)

• Bases accept H⁺ ions, forming conjugate acids.
• Many bases are metal oxides or hydroxides, but the key is proton acceptance.
• Alkalis are bases that dissolve in water to produce OH⁻ ions.

In acid-base reactions, the acid donates a proton to the base. After donating a proton, the acid becomes its conjugate base, and after accepting a proton, the base becomes its conjugate acid.

For example, in the reaction: HCl + NH₃ → NH₄⁺ + Cl⁻

HCl is the acid (proton donor) and Cl⁻ is its conjugate base (after losing H⁺). NH₃ is the base (proton acceptor) and NH₄⁺ is its conjugate acid (after gaining H⁺).

• Acid-base reactions involve proton transfer between conjugate pairs.
• The conjugate base is the species formed after an acid loses a proton.
• The conjugate acid is the species formed after a base gains a proton.

Write a balanced equation showing the proton transfer when ethanoic acid reacts with ammonia.

Ethanoic acid (CH₃COOH) donates a proton to ammonia (NH₃), which accepts it:

CH₃COOH (aq) + NH₃ (aq) ⇌ CH₃COO⁻ (aq) + NH₄⁺ (aq) (proton transfer)

• Ethanoic acid is the proton donor (acid).
• Ammonia is the proton acceptor (base).
• The products are the conjugate base (CH₃COO⁻) and conjugate acid (NH₄⁺).

A strong acid is a substance that completely dissociates into its ions when dissolved in water. This means that every molecule of the acid breaks apart to release hydrogen ions (H⁺) and the corresponding negative ions. Because dissociation is complete, no undissociated acid molecules remain in solution.

For example, hydrochloric acid (HCl) is a strong acid. When dissolved in water, it dissociates fully:

HCl (aq) → H⁺ (aq) + Cl⁻ (aq) (complete dissociation)

• Strong acids release all their hydrogen ions into solution.
• No equilibrium exists because dissociation goes to completion.
• Examples include hydrochloric acid, nitric acid, and sulfuric acid (first dissociation).

A weak acid only partially dissociates in water. This means that in solution, an equilibrium exists between undissociated acid molecules and ions. Only a fraction of the acid molecules release hydrogen ions, while the rest remain intact.

For example, ethanoic acid (CH₃COOH) is a weak acid. Its dissociation in water is represented by an equilibrium:

CH₃COOH (aq) ⇌ H⁺ (aq) + CH₃COO⁻ (aq) (partial dissociation)

• Weak acids do not release all their hydrogen ions.
• An equilibrium is established between dissociated ions and undissociated molecules.
• The position of equilibrium depends on the acid's strength and concentration.
• Examples include ethanoic acid, citric acid, and carbonic acid.

The partial dissociation of weak acids can be described by an equilibrium constant, Ka, which quantifies the extent of ionisation. The larger the Ka, the stronger the acid (more dissociation).

For ethanoic acid: Ka = [H⁺][CH₃COO⁻] / [CH₃COOH] (concentrations in mol dm⁻³)

Because the dissociation is partial, the solution contains a mixture of ions and undissociated acid molecules. This affects the pH and reactivity of the acid solution.

• Weak acid dissociation is reversible and establishes equilibrium.
• Ka values indicate acid strength quantitatively.
• Lower Ka means weaker acid and less dissociation.

Since strong acids fully dissociate, their solutions have a higher concentration of hydrogen ions (H⁺) compared to weak acids at the same concentration. This results in lower pH values for strong acid solutions.

For example, a 0.1 mol dm⁻³ solution of hydrochloric acid has a pH close to 1, while a 0.1 mol dm⁻³ solution of ethanoic acid has a pH around 2.9 due to partial dissociation.

This difference in ion concentration affects the acid's reactivity and strength in chemical reactions.

• Strong acids produce higher [H⁺] and lower pH at equal concentrations.
• Weak acids produce lower [H⁺] and higher pH due to partial dissociation.
• pH measurement can help distinguish strong and weak acids experimentally.

Write the dissociation equations for hydrochloric acid and ethanoic acid in water, indicating whether dissociation is complete or partial.

Answer: Hydrochloric acid dissociates completely: HCl (aq) → H⁺ (aq) + Cl⁻ (aq)

Ethanoic acid dissociates partially: CH₃COOH (aq) ⇌ H⁺ (aq) + CH₃COO⁻ (aq)

• Use → for complete dissociation (strong acids).
• Use ⇌ for partial dissociation (weak acids).

A strong acid is an acid that completely dissociates into its ions when dissolved in water. This means that every molecule of the acid breaks apart to release hydrogen ions (H⁺) and the corresponding negative ions. Because of this, strong acids produce a high concentration of H⁺ ions in solution, which is responsible for their characteristic acidic properties such as low pH and high reactivity.

• Complete dissociation means 100% of acid molecules ionise.
• High concentration of H⁺ ions results from complete dissociation.
• Strong acids have a low pH, typically close to 0–1 in concentrated form.

Hydrochloric acid (HCl) is a common example of a strong acid. When HCl gas dissolves in water, it ionises completely into hydrogen ions and chloride ions. The process can be represented by the balanced ionic equation:

HCl (aq) → H⁺ (aq) + Cl⁻ (aq) This equation shows that each HCl molecule produces one H⁺ ion and one Cl⁻ ion in solution. Because the dissociation is complete, no undissociated HCl molecules remain in the solution.

• HCl is a strong acid because it fully ionises in water.
• The ionic equation is HCl (aq) → H⁺ (aq) + Cl⁻ (aq).
• No equilibrium exists because dissociation goes to completion.

Complete dissociation of hydrochloric acid means that the concentration of hydrogen ions in solution equals the initial concentration of the acid. For example, if 0.1 mol/dm³ of HCl is dissolved in water, the concentration of H⁺ ions will also be 0.1 mol/dm³. This high concentration of H⁺ ions causes the solution to be strongly acidic, which can be measured by a low pH value (around 1 for 0.1 mol/dm³ HCl).

• [H⁺] = initial concentration of HCl due to complete dissociation.
• High [H⁺] results in low pH (acidic solution).
• pH can be calculated using pH = -log[H⁺].

Unlike strong acids, weak acids only partially dissociate in water. This means that only a fraction of the acid molecules release hydrogen ions, resulting in an equilibrium between undissociated acid molecules and ions. For example, ethanoic acid (CH₃COOH) is a weak acid and dissociates partially:

CH₃COOH (aq) ⇌ H⁺ (aq) + CH₃COO⁻ (aq) Because hydrochloric acid dissociates completely, it produces a higher concentration of H⁺ ions than a weak acid of the same concentration, making it more acidic and reactive.

• Weak acids partially dissociate; strong acids fully dissociate.
• Equilibrium exists in weak acid dissociation, not in strong acids.
• Strong acids have higher [H⁺] and lower pH than weak acids at the same concentration.

Calculate the pH of a 0.01 mol/dm³ hydrochloric acid solution, assuming complete dissociation.

Step 1: Write the dissociation equation: HCl (aq) → H⁺ (aq) + Cl⁻ (aq)

Step 2: Determine [H⁺]: Since HCl fully dissociates, [H⁺] = 0.01 mol/dm³ = 1.0 × 10⁻² mol/dm³. Step 3: Calculate pH using the formula: pH = -log[H⁺] pH = -log(1.0 × 10⁻²) = 2.00 Step 4: Interpret the result: The pH of 2.00 confirms the solution is acidic due to the high concentration of hydrogen ions from complete dissociation of HCl.

• Use pH = -log[H⁺] to calculate acidity.
• Complete dissociation means [H⁺] equals acid concentration.
• Example: 0.01 mol/dm³ HCl has pH 2.00.

Acids are substances that release hydrogen ions (H⁺) when dissolved in water. The extent to which an acid releases H⁺ ions is called dissociation or ionisation.

A strong acid completely dissociates in aqueous solution, meaning all acid molecules release H⁺ ions.

A weak acid only partially dissociates in solution, so only some molecules release H⁺ ions while others remain undissociated. This partial dissociation establishes an equilibrium between the acid molecules and ions in solution.

• Strong acid: complete dissociation in water.
• Weak acid: partial dissociation in water, equilibrium established.

Ethanoic acid, commonly known as acetic acid, has the molecular formula CH3COOH.

Its structural formula is: H3C–C(=O)–OH, showing a methyl group (CH3) attached to a carboxyl group (COOH).

The carboxyl group is responsible for the acidic properties of ethanoic acid.

• Molecular formula: CH3COOH
• Structural formula: H3C–C(=O)–OH

When ethanoic acid dissolves in water, it partially ionises to produce hydrogen ions (H⁺) and ethanoate ions (CH3COO⁻).

The dissociation is reversible and establishes an equilibrium:

CH3COOH (aq) ⇌ H⁺ (aq) + CH3COO⁻ (aq) (equilibrium arrow)

• Partial ionisation leads to equilibrium between undissociated acid and ions.
• Only a fraction of molecules release H⁺ ions.

Because dissociation is partial, the concentration of H⁺ ions in ethanoic acid solution is lower than in a strong acid of the same concentration.

The position of equilibrium lies mostly to the left (undissociated acid).

This explains why ethanoic acid solutions have higher pH values (less acidic) compared to strong acids at the same molarity.

• Equilibrium lies towards undissociated acid.
• Lower H⁺ ion concentration than strong acids.
• Higher pH compared to strong acids of equal concentration.

Write the ionic equation for the dissociation of ethanoic acid in water.

Step 1: Write the formula of ethanoic acid: CH3COOH

Step 2: Show dissociation into ions: H⁺ and CH3COO⁻ ions form in solution, but only partially, so use equilibrium arrows: ⇌ Step 3: Write the ionic equation: CH3COOH (aq) ⇌ H⁺ (aq) + CH3COO⁻ (aq)

• Use ⇌ to indicate partial dissociation and equilibrium.
• Include state symbols (aq) for aqueous species.