Chemistry · Solutions and electrolytes · Grade 9

Electrolytic dissociation

Why does a lamp glow brightly in salt water, not at all in sugar water, and dimly in vinegar?

The situation

In a school chemistry room in Tashkent, the teacher put eight beakers on the table: distilled water, sugar water, table salt solution, hydrochloric acid, alkali, vinegar, ammonia solution and dry salt. The two graphite electrodes of the lamp apparatus are dipped into the beakers in turn: sometimes the lamp burns brightly, sometimes it barely glows, and sometimes it does not light at all — even in dry salt. What carries current in a liquid, where do the carriers come from, and why do some substances react with one another while others do not?

Topic

Electrolytic dissociation and ion-exchange reactions. Conductivity apparatus: the lamp lights in salt water, not in sugar water, and glows dimly in vinegar — where do the ions come from

Key ideas

Electrolytes and non-electrolytes
Degree of dissociation α
Ionic exchange reactions

Formulas

NaCl \overset{}{\to} Na^{+} + Cl^{-} CH_{3} COOH \overset{}{⇌} CH_{3} COO^{-} + H^{+}

A strong electrolyte dissociates completely; a weak electrolyte dissociates partially and reversibly.

α = \frac{n_{ionised}}{n_{total}} \cdot 100 %

Degree of dissociation: the fraction of molecules that have split into ions.

K = \frac{α^{2} c}{1 - α} \Rightarrow α \approx \sqrt{K / c}

Ostwald's dilution law (weak electrolyte): α increases on dilution. For CH₃COOH, $K = 1.75 \cdot 10^{- 5}$ .

κ = c α (λ_{+} + λ_{-}), R = \frac{K_{cell}}{κ}

The conductivity of a solution depends on the concentration and mobility of the ions; $R$ is the resistance between the electrodes.

Ag^{+} + Cl^{-} \overset{}{\to} AgCl ↓

Net ionic equation: only the ions that form a precipitate, a gas or water take part in the reaction.

Theory

Electrolytes and non-electrolytes

Substances whose solution or melt conducts electric current are electrolytes: salts, acids, alkalis. Those that do not are non-electrolytes: sugar, glucose, alcohol, most organic substances, oxygen, nitrogen. In the lamp apparatus the circuit is completed through the solution: current is carried not by electrons, as in a metal, but by ions in the solution — positive ions (cations) move towards one electrode, negative ions (anions) towards the other. A dry salt crystal also contains ions, but they are fixed in the lattice, so the lamp does not light; when salt is dissolved in water or melted, the ions become free. Distilled water conducts very poorly: it contains very few ions. Lab work (current through a concentrated salt solution): in the “Degree α” mode choose table salt, raise the concentration up to a saturated solution (≈ 5.3 mol/L), and record the current in the log each time — the current rises as the ions increase, but in a concentrated solution the rise slows down.

Arrhenius theory: where the ions come from

Svante Arrhenius (1887) put forward the theory of electrolytic dissociation: when electrolytes dissolve in water they split into ions by themselves, even without any current passing. The water molecule is polar: its oxygen side is negative and its hydrogen sides are positive. When an ionic crystal (NaCl) goes into water, the water molecules attach to Na⁺ with their oxygen side and to Cl⁻ with their hydrogen side and “pull” the ion out of the lattice; the ion moves surrounded by a hydration shell of water molecules (hydration). Polar molecules (HCl) also ionise in water: the proton passes to a water molecule and forms the hydroxonium ion $H_{3} O^{+}$ : $HCl + H_{2} O \overset{}{\to} H_{3} O^{+} + Cl^{-}$ (written more simply, $HCl \overset{}{\to} H^{+} + Cl^{-}$ ). In a weak electrolyte (ethanoic acid) dissociation is reversible: molecules ionise and the ions combine into molecules again — a dynamic equilibrium. The properties of ions are completely different from those of atoms: the Na⁺ ion does not react with water like sodium metal, and the Cl⁻ ion, unlike chlorine gas ( $Cl_{2}$ ), is not poisonous and is colourless.

Dissociation of acids, alkalis and salts

According to the theory of dissociation, an acid is an electrolyte that gives only hydrogen ions H⁺ as cations in water: $HCl \overset{}{\to} H^{+} + Cl^{-}$ . Polybasic acids dissociate in steps: $H_{2} SO_{4} \overset{}{\to} H^{+} + HSO_{4}^{-}$ , $HSO_{4}^{-} \overset{}{⇌} H^{+} + SO_{4}^{2 -}$ ; in phosphoric acid the first step $H_{3} PO_{4} \overset{}{⇌} H^{+} + H_{2} PO_{4}^{-}$ is the strongest. The common properties of acids (sour taste, indicator colour, reaction with metals) depend on H⁺ ions. An alkali is an electrolyte that gives only hydroxide ions OH⁻ as anions: $NaOH \overset{}{\to} Na^{+} + OH^{-}$ , $Ba (OH)_{2} \overset{}{\to} Ba^{2 +} + 2 OH^{-}$ . Salts split into metal cations (or $NH_{4}^{+}$ ) and anions of the acid residue: $Na_{2} SO_{4} \overset{}{\to} 2 Na^{+} + SO_{4}^{2 -}$ , $Al_{2} (SO_{4})_{3} \overset{}{\to} 2 Al^{3 +} + 3 SO_{4}^{2 -}$ . In a solution the sums of positive and negative charges are equal — the solution is electrically neutral.

Strong and weak electrolytes, degree α

The degree of dissociation $α$ is the ratio of the number of molecules that have split into ions to the total number of molecules. Strong electrolytes dissociate practically completely in dilute solution ( $α \to 1$ ): HCl, HNO₃, H₂SO₄, HBr, HI, alkalis (NaOH, KOH, Ba(OH)₂) and almost all salts. In weak electrolytes $α$ is small: CH₃COOH, H₂CO₃, H₂S, HF, NH₃·H₂O, water itself. Textbooks often use a conventional boundary for a 0.1 mol/L solution: $α$ > 30 % — strong, < 3 % — weak. In a 0.1 mol/L solution of ethanoic acid $α$ ≈ 1.3 %; it increases on dilution (Ostwald's law, $α \approx \sqrt{K / c}$ ): at 0.001 mol/L it is ≈ 12 %. But the concentration of ions ( $c α$ ) falls — which is why the current still decreases on dilution. When the temperature rises, α usually increases as well. In strong electrolytes the “apparent” degree found from conductivity, $Λ / Λ^{\circ}$ , is less than one (≈ 0.85 in 0.1 mol/L NaCl): the ions are fully separated, but oppositely charged ions slow each other down.

Ionic exchange reactions

Reactions between solutions of electrolytes actually take place between ions. A reaction goes to completion if the ions form a precipitate ( $Ag^{+} + Cl^{-} \overset{}{\to} AgCl ↓$ ), a gas ( $CO_{3}^{2 -} + 2 H^{+} \overset{}{\to} H_{2} O + CO_{2} ↑$ ) or a weakly dissociating substance, for example water ( $H^{+} + OH^{-} \overset{}{\to} H_{2} O$ ). The order for writing the equation: 1) the molecular equation; 2) the complete ionic equation — strong electrolytes are written as ions, while precipitates, gases, water and weak electrolytes are written as molecules; 3) identical ions on both sides (spectator ions) are cancelled — leaving the net ionic equation. It shows the essence of the reaction: Ag⁺ and Cl⁻ give an AgCl precipitate no matter which salts they come from. When NaCl and KNO₃ are mixed, no pair of ions gives a precipitate, a gas or water — all the ions stay in solution and no reaction takes place.

Safety and the model

Safety: use the conductivity apparatus only with a safe low-voltage (12 V) source, do not touch the electrodes or the solution, and before changing the solution lift the electrodes and rinse them with distilled water. Wear protective goggles when working with acids and alkalis; pour acid into water, never the other way round. AgNO₃ stains skin black and barium salts are poisonous — wear gloves and keep them away from drink and food. Model: 25 °C; the conductivity of strong 1:1 electrolytes comes from the Debye–Hückel–Onsager equation, and that of concentrated NaCl solutions from the CRC Handbook table; the $α$ of weak electrolytes is calculated exactly from $K$ , and in concentrated ethanoic acid the change of medium is accounted for by the empirical factor $(1 - 0.12 c)$ (it gives the measured maximum — ≈ 1.85 mS/cm, ≈ 17 %). The source is alternating voltage — almost no electrolysis occurs at the electrodes. Lamp: 12 V / 0.1 A, incandescent filament (current ∝ U^0.55, light ∝ U^3.4). The “drop” model is magnified about $10^{8}$ times and the motion is greatly slowed down; only some of the water molecules are shown.

Did you know?

Svante Arrhenius published the theory of electrolytic dissociation in 1887 and received the 1903 Nobel Prize in Chemistry for it.
The conductivity of very pure water at 25 °C is only ≈ 0.055 µS/cm, while that of seawater is ≈ 50 mS/cm — about a million times greater: the difference is made by the ions of salts dissolved in the water.

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

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