Precipitation, Acid-Base, and Redox Reactions
Sofia pours two clear liquids together in a beaker, and the mixture instantly turns cloudy yellow. Nothing was heated and nothing was stirred. What just happened, and could she have predicted it before pouring? By the end of this page, you will be able to sort reactions into three big families and predict which ones will make a solid.
Three families of reactions
Thousands of reactions happen in water, but most of them fit into three families. Each family is defined by what moves or what forms.
| Family | What happens | Clue |
|---|---|---|
| Precipitation | Ions combine to form an insoluble solid | A solid, written (s), appears from two aqueous solutions |
| Acid-base (neutralization) | An acid and a base react to form a salt and water | $\text{H}_2\text{O}(l)$ is a product |
| Oxidation-reduction (redox) | Electrons transfer from one species to another | Charges or oxidation numbers change |
Precipitation and acid-base reactions are both double replacement reactions: the partners swap, $\text{AB} + \text{CD} \rightarrow \text{AD} + \text{CB}$. Many synthesis, decomposition, and combustion reactions, and all single replacement reactions, are redox, because electrons move in those.
Classify each reaction as precipitation, acid-base (neutralization), or oxidation-reduction, and give a short reason for each: (a)…
Classify each reaction as precipitation, acid-base (neutralization), or oxidation-reduction, and give a short reason for each:
(a) $\text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$
(b) $\text{Zn}(s) + \text{CuSO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{Cu}(s)$
(c) $\text{AgNO}_3(aq) + \text{NaCl}(aq) \rightarrow \text{AgCl}(s) + \text{NaNO}_3(aq)$
Precipitation and solubility rules
Why do some mixtures make a solid and others do not? When an ionic compound dissolves, its ions separate and move freely. If two of those ions attract each other strongly enough to stay together as a solid, they leave the solution. Whether that happens depends on the compound, and chemists summarize it as solubility rules.
Always check the soluble rules first, because they have priority. For example, $\text{Na}_2\text{CO}_3$ is soluble even though most carbonates are not, since all sodium compounds are soluble.
Predicting a double replacement
If both possible products are soluble, there is no reaction: the ions just stay mixed in the water.
Use the general solubility rules (compounds of $\text{Na}^+$, $\text{K}^+$, $\text{NH}4^+$ and $\text{NO}3^-$ are soluble; most carbonates are insoluble except those of…
Use the general solubility rules (compounds of $\text{Na}^+$, $\text{K}^+$, $\text{NH}_4^+$ and $\text{NO}_3^-$ are soluble; most carbonates are insoluble except those of group 1 ions) to decide whether each compound is soluble or insoluble in water: (a) $\text{KNO}_3$, (b) $\text{CaCO}_3$, (c) $\text{Na}_2\text{CO}_3$.
Sofia mixes aqueous solutions of lead(II) nitrate, $\text{Pb(NO}3)2$, and potassium iodide, KI. Write the balanced molecular equation for the double replacement…
Sofia mixes aqueous solutions of lead(II) nitrate, $\text{Pb(NO}_3)_2$, and potassium iodide, KI. Write the balanced molecular equation for the double replacement reaction. Use the rules that nitrates are soluble and that most iodides are soluble except those of $\text{Pb}^{2+}$, $\text{Ag}^+$ and $\text{Hg}_2^{2+}$ to identify the precipitate, and include state symbols.
Net ionic equations
In the $\text{BaCl}_2$ + $\text{Na}_2\text{SO}_4$ reaction, only the barium and sulfate ions actually join together. The sodium and chloride ions are in the same state before and after. Net ionic equations show only what really changes.
Jayden mixes aqueous barium chloride, $\text{BaCl}2$, with aqueous sodium sulfate, $\text{Na}2\text{SO}4$. (a) Write the balanced net ionic equation, given that…
Jayden mixes aqueous barium chloride, $\text{BaCl}_2$, with aqueous sodium sulfate, $\text{Na}_2\text{SO}_4$. (a) Write the balanced net ionic equation, given that $\text{BaSO}_4$ is insoluble. (b) Identify the spectator ions. (c) Explain why this is not a redox reaction.
Acid-base reactions
You already know that an acid releases $\text{H}^+$ in water and a base releases $\text{OH}^-$ (or accepts $\text{H}^+$). When they meet, the $\text{H}^+$ and $\text{OH}^-$ combine into water, and the remaining ions form a salt.
$$\text{acid} + \text{base} \rightarrow \text{salt} + \text{water}$$ $$\text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$$This is why the beaker warms up: neutralization releases energy. The net ionic equation for a strong acid and a strong base is always the same:
$$\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$$Acid-base reactions are double replacement reactions, but the clue that separates them from precipitation is the water product, not a solid.
Oxidation-reduction (redox) reactions
In redox reactions, electrons move from one species to another. Two words describe the two halves.
Take $2\,\text{Na}(s) + \text{Cl}_2(g) \rightarrow 2\,\text{NaCl}(s)$. Each Na atom becomes $\text{Na}^+$ by losing one electron (oxidized). Each Cl atom becomes $\text{Cl}^-$ by gaining one electron (reduced).
Oxidation numbers
When no ions are visible (as in covalent compounds), track electrons with oxidation numbers. A few rules are enough:
-
An atom in an uncombined element is 0 (Zn, $\text{Cl}_2$, Cu).
-
A monatomic ion has an oxidation number equal to its charge ($\text{Na}^+ = +1$).
-
Hydrogen is usually $+1$ and oxygen usually $-2$.
-
The oxidation numbers in a neutral compound add to 0 (in an ion, to its charge).
A reaction is redox if any oxidation number changes. If the number increases, that atom is oxidized; if it decreases, it is reduced.
Define oxidation and reduction in terms of the transfer of electrons. Then, in the reaction $2\,\text{Na}(s) + \text{Cl}2(g) \rightarrow 2\,\text{NaCl}(s)$, identify…
Define oxidation and reduction in terms of the transfer of electrons. Then, in the reaction $2\,\text{Na}(s) + \text{Cl}_2(g) \rightarrow 2\,\text{NaCl}(s)$, identify which species is oxidized and which is reduced.
For the reaction $\text{Mg}(s) + 2\,\text{HCl}(aq) \rightarrow \text{MgCl}2(aq) + \text{H}2(g)$, assign oxidation numbers to Mg, H, and Cl on both sides of the equation.…
For the reaction $\text{Mg}(s) + 2\,\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2(g)$, assign oxidation numbers to Mg, H, and Cl on both sides of the equation. Identify the species oxidized and reduced, the oxidizing agent, and the reducing agent.
Putting it all together: classifying and identifying
When you meet a new reaction, ask these questions in order:
-
Is any element in its uncombined form (a metal such as Zn, or $\text{O}_2$)? If so, it is probably redox.
-
Do two aqueous solutions form an insoluble solid? That is precipitation.
-
Is water a product, from an acid and a base? That is acid-base.
Notice that a solid forming does not mean precipitation. In $\text{Cu}(s) + 2\,\text{AgNO}_3(aq) \rightarrow \text{Cu(NO}_3)_2(aq) + 2\,\text{Ag}(s)$, the silver solid forms because $\text{Ag}^+$ gained electrons from copper. The net ionic equation is $\text{Cu}(s) + 2\,\text{Ag}^+(aq) \rightarrow \text{Cu}^{2+}(aq) + 2\,\text{Ag}(s)$. Copper goes from 0 to $+2$ (oxidized), silver from $+1$ to 0 (reduced). The blue color is $\text{Cu}^{2+}$ in solution.
A lab tech has four unlabeled beakers, each containing a clear, colorless aqueous solution of one of these: $\text{NaOH}$, $\text{HNO}3$, $\text{Na}2\text{SO}4$, or…
A lab tech has four unlabeled beakers, each containing a clear, colorless aqueous solution of one of these: $\text{NaOH}$, $\text{HNO}_3$, $\text{Na}_2\text{SO}_4$, or $\text{Pb(NO}_3)_2$. She mixes pairs and records the results:
| Mixture | Observation |
|---|---|
| A + B | White solid forms |
| A + C | No visible change; the mixture feels warm |
| B + D | White solid forms |
| C + D | No visible change |
Assume $\text{Pb(OH)}_2$ and $\text{PbSO}_4$ are insoluble white solids. Identify A, B, C, and D, and write a balanced net ionic equation for each of the three mixtures that shows a reaction. Classify each reaction type.
Emma drops a strip of copper metal into a solution of silver nitrate, $\text{AgNO}3(aq)$, and after an hour she sees a gray, feathery solid on the strip while the…
Emma drops a strip of copper metal into a solution of silver nitrate, $\text{AgNO}_3(aq)$, and after an hour she sees a gray, feathery solid on the strip while the solution turns pale blue. (a) Write the balanced net ionic equation. (b) Assign oxidation numbers and identify what is oxidized and what is reduced. (c) Explain how each observation supports your equation. (d) Explain why this reaction is redox and not a precipitation reaction, even though a solid forms.
Back to Sofia's beaker: when she mixed lead(II) nitrate and potassium iodide, the solubility rules predicted insoluble $\text{PbI}_2$, so the cloudy solid was a precipitate (bright yellow, in fact).