Lesson 6 · 35 min

Reactions in aqueous solution

Electrolytes and ions; the three driving forces — precipitate, gas, water; molecular, complete ionic and net ionic equations; the activity series and displacement.

In this lesson

Distinguish strong electrolytes, weak electrolytes and non-electrolytes
Use the solubility table and “precipitate, gas or water” to decide whether an exchange reaction occurs
Write molecular, complete ionic and net ionic equations and identify the spectator ions
Predict displacement reactions from the activity series (standard electrode potentials)

Pour colourless lead nitrate solution into colourless potassium iodide and a bright yellow “golden rain” appears at once. Each solution was clear on its own; mixed, they make a visible solid — some of the particles in solution met, locked together and left the solution. The key to reactions in water is a change of viewpoint: look at the ions, not the compounds.

Predict first

Silver nitrate solution is mixed with sodium chloride solution. What do you see?

Pick an answer and lock it in before you look. A wrong guess teaches more than no guess.

Nothing — both salts are soluble A white precipitate of sodium nitrate A white precipitate of silver chloride Bubbles of gas

Electrolytes: what is in the solution

Compounds that release ions in water and make it conduct are electrolytes. Soluble salts, strong acids (HCl, HNO₃, H₂SO₄) and strong bases (NaOH, KOH) ionise almost completely — strong electrolytes: 0.1 mol/L $\ce{Pb(NO3)2}$ is really 0.1 mol/L $\ce{Pb^2+}$ and 0.2 mol/L $\ce{NO3-}$ . Acetic acid and ammonia ionise only slightly (weak electrolytes); sugar and ethanol dissolve without ionising (non-electrolytes).

An exchange (double-displacement) reaction goes when two ions combine and leave the solution. They can leave in three ways:

as a precipitate (an insoluble solid) such as $\ce{AgCl}$ or $\ce{BaSO4}$ ;
as a gas such as $\ce{CO2}$ ;
as water (or another weak electrolyte), as in neutralisation.

Solubility of some acids, bases and salts (room temperature): sol = soluble, sl = slightly soluble, ins = insoluble, — = does not exist or decomposes in water
	$\ce{OH-}$	$\ce{NO3-}$	$\ce{Cl-}$	$\ce{SO4^2-}$	$\ce{CO3^2-}$
$\ce{H+}$	— (water)	sol	sol	sol	sol, gas (CO₂)
$\ce{Na+}$ , $\ce{K+}$	sol	sol	sol	sol	sol
$\ce{Ba^2+}$	sol	sol	sol	ins	ins
$\ce{Ca^2+}$	sl	sol	sol	sl	ins
$\ce{Cu^2+}$	ins	sol	sol	sol	—
$\ce{Fe^3+}$	ins	sol	sol	sol	—
$\ce{Ag+}$	—	sol	ins	sl	ins
$\ce{Pb^2+}$	ins	sol	sl	ins	ins

Source: 人民教育出版社《化学》九年级下册附录 I（部分酸、碱和盐的溶解性表，室温）; Pb²⁺ row after S. S. Zumdahl & S. A. Zumdahl, Chemistry, 10th ed., Table 4.1.

Three equations: write, split, cancel, check

Write the balanced molecular equation: $\ce{BaCl2} + \ce{Na2SO4} \mathrel{\ce{->}} \ce{BaSO4 v} + \ce{2NaCl}$ .
Split every soluble strong electrolyte into its ions to get the complete ionic equation; keep precipitates, gases, water and weak electrolytes as formulas: $\ce{Ba^2+} + \ce{2Cl-} + \ce{2Na+} + \ce{SO4^2-} \mathrel{\ce{->}} \ce{BaSO4(s)} + \ce{2Na+} + \ce{2Cl-}$ .
Cancel ions that appear unchanged on both sides — here $\ce{Na+}$ and $\ce{Cl-}$ , the spectator ions — leaving the net ionic equation: $\ce{Ba^2+} + \ce{SO4^2-} \mathrel{\ce{->}} \ce{BaSO4(s)}$ .
Check that atoms and charge balance and the coefficients are in lowest terms.

\ce{Ba^2+(aq) + SO4^2-(aq) -> BaSO4(s)} \qquad \ce{H+(aq) + OH-(aq) -> H2O(l)}

A net ionic equation shows only the particles that change: any soluble barium salt with any soluble sulfate gives the left-hand equation.

$(aq)$: dissolved (hydrated) particle
$(s)$: solid — here a precipitate
$(l)$: liquid — here the water formed

Worked example

Weighing sulfate as a precipitate

Excess $\ce{Na2SO4}$ solution is added to 25.0 mL of 0.100 mol/L $\ce{BaCl2}$ . Write the net ionic equation and find the mass of $\ce{BaSO4}$ precipitated.

Precipitate · gas · water · redox

Reactions-in-water bench

Every bottle is 0.100 mol/L

In the tube: 10.0 mL of solution A

Add solution B

Volume of B added

The tube after mixing

What you see

pale blue precipitate of Cu(OH)₂ (0.50 mmol)
Insoluble Cu(OH)₂ precipitates.

Cu(OH)₂ · pale blue

Precipitate colours are observed colours documented in reference texts

Molecular equation

\ce{CuSO4(aq) + 2NaOH(aq) -> Cu(OH)2(s) + Na2SO4(aq)}

Complete ionic equation

\ce{Cu^2+(aq) + SO4^2-(aq) + 2Na+(aq) + 2OH-(aq) -> Cu(OH)2(s) + 2Na+(aq) + SO4^2-(aq)}

Net ionic equation

\ce{Cu^2+(aq) + 2OH-(aq) -> Cu(OH)2(s)}

Spectator ions: SO₄²⁻, Na⁺

Particles: ions in the mixture (schematic)

+ −

One

\ce{Cu^2+}

gets only one

\ce{OH-}

: half the copper stays dissolved and the liquid stays pale blue — raise B to 20.0 mL and look again. Then try lead nitrate with sodium chloride (slightly soluble

\ce{PbCl2}

stays dissolved at this concentration), hydrochloric acid with sodium carbonate (no fizz until the acid is in excess), and copper sulfate with potassium iodide (a redox reaction, not a precipitation).

Redox: displacing metals

Dip a zinc strip into blue copper sulfate solution: red-brown copper soon coats the zinc and the blue fades, $\ce{Zn(s)} + \ce{Cu^2+(aq)} \mathrel{\ce{->}} \ce{Zn^2+(aq)} + \ce{Cu(s)}$ . Zinc loses electrons (it is oxidised) and copper ions gain them (they are reduced). Whether such a reaction goes is decided by the activity series: K, Ca, Na, Mg, Al, Zn, Fe, Sn, Pb, (H), Cu, Hg, Ag, Pt, Au. A metal displaces any metal after it from solutions of its salts, and metals before hydrogen release hydrogen from dilute hydrochloric acid.

The quantitative basis is the standard electrode potential $E^\circ$ : the lower a metal’s $E^\circ$ , the more readily it gives up electrons. Displacement is spontaneous when the dissolved ion’s $E^\circ$ is higher than that of the metal’s own ion.

Precipitate · gas · water · redox

Reactions-in-water bench

Metal strip

Solution

The tube after mixing

Displacement happens: E°(Ag⁺) − E°(Cu²⁺) = 0.46 V > 0

Deposit of silver on the strip: grey-white crystals (“silver tree”)
Solution colour change (computed): before → after

Molecular equation

\ce{Cu(s) + 2AgNO3(aq) -> Cu(NO3)2(aq) + 2Ag(s)}

Net ionic equation

\ce{Cu(s) + 2Ag+(aq) -> Cu^2+(aq) + 2Ag(s)}

Activity series (standard reduction potential E° / V)

+0.800 Ag⁺ / Ag Solution
+0.342 Cu²⁺ / Cu Metal strip
+0.000 H⁺ / H₂
−0.126 Pb²⁺ / Pb
−0.447 Fe²⁺ / Fe
−0.762 Zn²⁺ / Zn
−2.372 Mg²⁺ / Mg

Copper in silver nitrate grows a grey-white “silver tree”, and the colourless solution turns blue (computed from the

\ce{Cu^2+}

formed). Then put copper in hydrochloric acid, or silver in copper sulfate, and use the potential ladder to explain why nothing happens.

Common pitfall Misconception: “insoluble” means none dissolves at all

“Insoluble” means a very small solubility (under roughly 0.01 g per 100 g of water at 20 °C), not zero; whether a “slightly soluble” salt appears depends on the ion concentrations. That is why 0.1 mol/L lead nitrate and 0.1 mol/L sodium chloride never give a precipitate on the bench, whatever the volumes: $Q = [\ce{Pb^2+}][\ce{Cl-}]^2$ reaches at most $1.5\times10^{-4}$ (at a 1 : 2 volume ratio), below $K^{\prime} \approx 1.7\times10^{-4}$ — while 1 mol/L solutions of each throw down white $\ce{PbCl2}$ at once. The other misconception is that spectator ions “disappear” — they stay in solution throughout; they just do not change.

Concept check

What is the net ionic equation for hydrochloric acid reacting with sodium hydroxide solution?

\ce{HCl} + \ce{NaOH} \mathrel{\ce{->}} \ce{NaCl} + \ce{H2O}

\ce{H+} + \ce{OH-} \mathrel{\ce{->}} \ce{H2O}

\ce{Na+} + \ce{Cl-} \mathrel{\ce{->}} \ce{NaCl}

No reaction — no precipitate forms

Try the calculation

Excess NaCl solution is added to 20.0 mL of 0.100 mol/L $\ce{AgNO3}$ . What mass of AgCl precipitates?

Key takeaways

Strong electrolytes ionise completely; exchange reactions go when a precipitate, a gas or water forms.
Ionic equations: write the molecular equation, split strong electrolytes into ions, cancel spectators, check atoms and charge.
A net ionic equation keeps only the particles that change, and covers every reaction of that kind.
A more active metal (lower E°) displaces a less active one from its salt; metals above hydrogen release H₂ from dilute hydrochloric acid.

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Data and references

人民教育出版社《化学》九年级下册（2012），附录 I 部分酸、碱和盐的溶解性表（室温）；《普通高中教科书化学必修第一册》（2019），第一章第二节离子反应。
S. S. Zumdahl, S. A. Zumdahl, Chemistry, 10th ed. (Cengage, 2018), Ch. 4 (types of chemical reactions and solution stoichiometry), Table 4.1.
CRC Handbook of Chemistry and Physics, 97th ed.: “Aqueous solubility of inorganic compounds” and “Electrochemical series” (P. Vanýsek).
N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed. — colours of precipitates and deposits.