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Electrolytes and the Dissolving Process

What happens when a substance dissolves, and why do some solutions conduct electricity?

BeginnerSolutionsLast reviewed 6 October 2026
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What is it?

A solution is a uniform mixture. The substance present in the larger amount is the solvent (usually water); the dissolved substance is the solute.

When a substance dissolves, its particles separate and become surrounded by solvent molecules. Water is a polar molecule, with a slightly negative oxygen and slightly positive hydrogens, so:

  • Ionic compounds split into separate ions: NaCl(s)→NaX+(aq)+ClX−(aq)\ce{NaCl(s) -> Na+(aq) + Cl-(aq)}. Each ion is surrounded by water molecules (it is hydrated): the oxygen side faces positive ions and the hydrogen side faces negative ions.
  • Polar molecular compounds such as sugar and ethanol dissolve as whole molecules, held by hydrogen bonds and dipole attractions.

A solution that conducts electricity contains ions that can move. Solutes are classified by how many ions they give:

TypeIn waterConducts?Examples
Strong electrolyteCompletely split into ionsStronglySoluble ionic compounds (NaCl, CaCl₂), strong acids (HCl, HNO₃), strong bases (NaOH)
Weak electrolyteOnly a small fraction forms ionsWeaklyWeak acids (ethanoic acid), weak bases (ammonia)
Non-electrolyteStays as moleculesNoSugar, ethanol, urea

Key idea

Electrical conductivity in solution comes from moving ions. The more ions a solution contains, the better it conducts. That is why salt water conducts, sugar water doesn’t, and vinegar conducts only weakly.

Why does it matter?

  • Your body runs on electrolytes. Nerve signals and heartbeats depend on NaX+\ce{Na+}, KX+\ce{K+} and CaX2+\ce{Ca^2+} ions; sports drinks replace the ions lost in sweat.
  • Counting particles. One mole of CaClX2\ce{CaCl2} gives three moles of ions. This matters for reactions, for colligative properties and for net ionic equations.
  • Hot and cold packs. Some substances absorb heat as they dissolve, others release it.

How does it work?

1. Writing dissociation equations

For a strong electrolyte, show every ion with its coefficient:

CaClX2(s)→CaX2+(aq)+2 ClX−(aq)AlX2(SOX4)X3(s)→2 AlX3+(aq)+3 SOX4X2−(aq)\small\begin{aligned} &\ce{CaCl2(s) -> Ca^2+(aq) + 2Cl-(aq)} \\[8pt] &\ce{Al2(SO4)3(s)} \\[4pt] &\quad \ce{-> 2Al^3+(aq) + 3SO4^2-(aq)} \end{aligned}

A weak electrolyte is shown with an equilibrium arrow, because most of it stays as molecules:

CHX3COOH(aq)⇌HX+(aq)+CHX3COOX−(aq)\small\begin{aligned} &\ce{CH3COOH(aq)} \\[4pt] &\quad \ce{<=> H+(aq) + CH3COO-(aq)} \end{aligned}

2. Ion concentrations

Multiply the concentration of the compound by the number of each ion in its formula. In 0.150 mol/L CaClX2\ce{CaCl2}: [CaX2+]=0.150[\ce{Ca^2+}] = 0.150 mol/L and [ClX−]=2×0.150=0.300[\ce{Cl-}] = 2 \times 0.150 = 0.300 mol/L.

3. Energy changes on dissolving

Dissolving an ionic solid takes two steps:

  1. Breaking up the lattice (separating the ions): needs energy, endothermic.
  2. Hydrating the ions (attracting water molecules): releases energy, exothermic.

If hydration releases more energy than the lattice needs, dissolving is exothermic and the solution warms up (CaClX2\ce{CaCl2}, NaOH\ce{NaOH}). If it releases less, dissolving is endothermic and the solution cools (NHX4NOX3\ce{NH4NO3}, used in instant cold packs).

Think of it like this

Dissolving salt is like a crowd breaking up at a party. Each guest (ion) leaves their own group and is soon surrounded by a circle of new friends (water molecules). If the new friendships are stronger than the old ones, energy is released; if they are weaker, energy has to be put in.

More precisely

Even strong electrolytes are not perfectly “free” ions in concentrated solutions: oppositely charged ions attract each other and form ion pairs, so solutions behave as if they contain slightly fewer particles. The “strong” and “weak” labels describe the extent of ionization, not the concentration: a dilute strong acid can be less acidic than a concentrated weak one. Whether a substance dissolves also depends on entropy (the spreading out of particles), which favours dissolving even when it is slightly endothermic.

Visualise it

A sodium ion and a chloride ion in water. Around the positive sodium ion, water molecules point their slightly negative oxygen atoms inward. Around the negative chloride ion, water molecules point their slightly positive hydrogen atoms inward.
Hydrated ions: water molecules turn their oppositely charged ends towards each ion.

Worked example

Worked example: Classifying solutes

Question: Classify each as a strong, weak or non-electrolyte: (a) KNO₃ (b) glucose (c) HCl (d) NH₃ (e) ethanol

  1. (a) Soluble ionic compound: strong.
  2. (b) Molecular, does not ionize: non-electrolyte.
  3. (c) Strong acid, fully ionized: strong.
  4. (d) Weak base, only partly ionized: weak.
  5. (e) Molecular, does not ionize: non-electrolyte.

Worked example: Ion concentrations

Question: What are the ion concentrations in 0.0500 mol/L AlX2(SOX4)X3\ce{Al2(SO4)3}?

  1. AlX2(SOX4)X3→2 AlX3++3 SOX4X2−\ce{Al2(SO4)3 -> 2Al^3+ + 3SO4^2-}
  2. [AlX3+]=2×0.0500 mol/L=0.100 mol/L[\ce{Al^3+}] = 2 \times 0.0500\ \text{mol/L} = 0.100\ \text{mol/L}
  3. [SOX4X2−]=3×0.0500 mol/L=0.150 mol/L[\ce{SO4^2-}] = 3 \times 0.0500\ \text{mol/L} = 0.150\ \text{mol/L}
  4. Total ions: 0.2500.250 mol/L, five times the concentration of the compound.

Worked example: Mixing two solutions

Question: 25.0 mL of 0.200 mol/L NaCl is mixed with 25.0 mL of 0.100 mol/L CaClX2\ce{CaCl2}. What is [ClX−][\ce{Cl-}] in the mixture?

  1. Moles of ClX−\ce{Cl-} from NaCl: 0.200 mol/L×0.0250 L=5.00×10−3 mol0.200\ \text{mol/L} \times 0.0250\ \text{L} = 5.00 \times 10^{-3}\ \text{mol}

  2. From CaClX2\ce{CaCl2} (2 Cl⁻ each): 2×0.100 mol/L×0.0250 L=5.00×10−3 mol2 \times 0.100\ \text{mol/L} \times 0.0250\ \text{L} = 5.00 \times 10^{-3}\ \text{mol}

  3. Total volume 0.0500 L:

    [ClX−]=1.000×10−2 mol0.0500 L=0.200 mol/L\small\begin{aligned} &[\ce{Cl-}] \\[4pt] &= \frac{1.000 \times 10^{-2}\ \text{mol}}{0.0500\ \text{L}} \\[4pt] &= 0.200\ \text{mol/L} \end{aligned}

Worked example: A cold pack

Question: 10.0 g of NHX4NOX3\ce{NH4NO3} dissolves in 100. g of water. ΔH(solution) = +25.7 kJ/mol. By how much does the temperature fall?

  1. Moles: 10.0 g80.05 g/mol=0.1249 mol\dfrac{10.0\ \text{g}}{80.05\ \text{g/mol}} = 0.1249\ \text{mol}

  2. Heat absorbed: 0.1249 mol×25.7 kJ/mol=3.210 kJ=3210 J0.1249\ \text{mol} \times 25.7\ \text{kJ/mol} = 3.210\ \text{kJ} = 3210\ \text{J}

  3. This heat comes from the 110. g of solution:

    ΔT=qmc=3210 J110. g×4.184 Jg °C=6.98 °C\small\begin{aligned} &\Delta T = \frac{q}{mc} \\[4pt] &= \frac{3210\ \text{J}}{110.\ \text{g} \times 4.184\ \tfrac{\text{J}}{\text{g °C}}} \\[4pt] &= 6.98\ \text{°C} \end{aligned}

  4. The solution cools by about 7 °C.

Common mistake

Common mistake: Thinking all dissolved substances conduct

Sugar dissolves completely but forms no ions, so sugar solution does not conduct. Dissolving and ionizing are different things.

Common mistake: Forgetting the subscripts

CaClX2\ce{CaCl2} gives two chloride ions per formula unit, so [ClX−][\ce{Cl-}] is twice the concentration of the compound.

Common mistake: Splitting polyatomic ions

NaX2SOX4\ce{Na2SO4} gives 2 NaX++SOX4X2−\ce{2Na+ + SO4^2-}. The sulfate ion stays together; it does not break into S and O.

Notation note

  • Square brackets mean concentration in mol/L: [ClX−][\ce{Cl-}].
  • (aq) means “dissolved in water” (aqueous).

Remember this

Remember this

  • Solute dissolves in solvent; water is polar and hydrates ions.
  • Strong electrolytes: soluble ionic compounds, strong acids and bases (fully ionized). Weak: weak acids and bases. Non-electrolytes: molecular solutes such as sugar.
  • Ion concentration = compound concentration × number of that ion in the formula.
  • Dissolving: lattice breaking (endothermic) + hydration (exothermic). The balance decides hot or cold.

Test yourself

Check your understanding before moving on.

Flashcards

10 cards

  1. Question
    What are the solute and the solvent in a solution?
    Answer

    Solute: the substance dissolved. Solvent: the substance it dissolves in (usually present in larger amount, often water).

  2. Question
    What is an electrolyte?
    Answer

    A substance that forms ions in water, so its solution conducts electricity.

  3. Question
    Give examples of strong, weak and non-electrolytes.
    Answer

    Strong: NaCl, HCl, NaOH. Weak: ethanoic acid, ammonia. Non-electrolyte: sugar, ethanol.

  4. Question
    What does "hydrated" mean for an ion?
    Answer

    The ion is surrounded by water molecules, with their oppositely charged ends facing it.

  5. Question
    Which end of a water molecule faces a Na⁺ ion?
    Answer

    The oxygen (δ−) end.

  6. Question
    Write the dissociation equation for CaCl₂.
    Answer

    CaClX2(s)→CaX2+(aq)+2 ClX−(aq)\ce{CaCl2(s) -> Ca^2+(aq) + 2Cl-(aq)}

  7. Question
    What is [Cl⁻] in 0.150 mol/L CaCl₂?
    Answer

    2 × 0.150 mol/L = 0.300 mol/L

  8. Question
    Why does sugar solution not conduct electricity?
    Answer

    Sugar dissolves as whole molecules and forms no ions.

  9. Question
    What two energy steps occur when an ionic solid dissolves?
    Answer

    Breaking the lattice (endothermic) and hydrating the ions (exothermic).

  10. Question
    Why does dissolving NH₄NO₃ make water colder?
    Answer

    Hydration releases less energy than breaking the lattice needs, so heat is taken from the water (ΔH positive).

Quiz

7 questions

  1. Question 1EasyWhich solution conducts electricity best?
    Show answer

    Answer: 0.10 mol/L NaCl

    NaCl is a strong electrolyte, fully split into ions. Ethanoic acid is only partly ionized, sugar forms no ions and pure water has very few.

  2. Question 2EasyWhich substance is a non-electrolyte?
    Show answer

    Answer: Ethanol

    Ethanol dissolves as molecules and does not ionize. The others are a strong base, a strong acid and a soluble salt.

  3. Question 3EasyAround a dissolved Cl⁻ ion, water molecules point which atoms towards the ion?
    Show answer

    Answer: Hydrogen atoms

    The hydrogen atoms carry a partial positive charge (δ+), attracted to the negative chloride ion.

  4. Question 4MediumWhat is [Cl⁻] in 0.150 mol/L FeCl₃?
    Show answer

    Answer: 0.450 mol/L

    Each FeCl₃ gives three Cl⁻ ions: 3 × 0.150 mol/L = 0.450 mol/L.

  5. Question 5MediumWhich 0.10 mol/L solution has the highest total concentration of ions?
    Show answer

    Answer: CaCl₂

    CaCl₂ gives 3 ions per formula unit (0.30 mol/L of ions); NaCl gives 2 (0.20 mol/L); ethanoic acid gives very few; glucose gives none.

  6. Question 6MediumDissolving a salt makes the solution warmer. What does this tell you?
    Show answer

    Answer: Hydration releases more energy than breaking the lattice needs

    A warmer solution means dissolving is exothermic: the energy released by hydrating the ions exceeds the energy needed to separate them.

  7. Question 7Hard10.0 g of NH₄NO₃ (80.05 g/mol, ΔH(soln) = +25.7 kJ/mol) dissolves in water. How much heat is absorbed?
    Show answer

    Answer: 3.21 kJ

    n = 10.0 g ÷ 80.05 g/mol = 0.1249 mol; q = 0.1249 mol × 25.7 kJ/mol = 3.21 kJ.

Notes and downloads

  • Worksheet

    Electrolytes and Dissolving Worksheet

    8 questions on electrolytes, dissociation equations, ion concentrations, mixing solutions and energy changes on dissolving. Answer key included.

    BeginnerFree

References

  1. Brown, T. L.; LeMay, H. E., Jr.; Bursten, B. E.; Murphy, C. J.; Woodward, P. M.; Stoltzfus, M. W. Chemistry: The Central Science, 15th ed.; Pearson, 2022.

Practise this topic with flashcards and a quiz at chemistryclarity.com/chemistry/electrolytes-and-dissolving/

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