Chapter 7: Reactions in Aqueous Solutions

Chapter 7: Reactions in Aqueous Solutions

Predicting Whether a Reaction Will Occur

• Four Driving Forces Favor Chemical Change

  • Formation of a solid (Precipitation reactions)

  • Formation of water (Acid-base neutralization reactions)

  • Transfer of electrons (Oxidation-reduction or Redox reactions)

  • Formation of a gas (Gas evolution reactions)

Reactions in Which a Solid Forms (Precipitation)

• Precipitation: The process of formation of an insoluble solid during a chemical reaction when two solutions are mixed.

  • The solid formed is called a precipitate.

  • The reaction is known as a precipitation reaction.

• What Happens When an Ionic Compound Dissolves in Water?

  • When an ionic compound dissolves, its component ions separate from the solid lattice and move around independently in the water. This process is called dissociation.

  • Strong electrolyte: A substance that completely dissociates or ionizes into separated ions when dissolved in water, making the solution an excellent conductor of electricity. All soluble ionic compounds are strong electrolytes.

  • Example: In the reaction $K<em>2CrO</em>4(aq) + Ba(NO<em>3)</em>2(aq) <br>ightarrow Products$, both $K<em>2CrO</em>4$ and $Ba(NO<em>3)</em>2$ dissociate completely into their respective ions in aqueous solution.

• How to Decide What Products Form

  • In the example $K<em>2CrO</em>4(aq) + Ba(NO<em>3)</em>2(aq) <br>ightarrow Products$, the mixed solution initially contains four types of free ions: $K^+, CrO<em>4^{2-}, Ba^{2+},$ and $NO</em>3^-$ .

  • To determine the possible products, one must consider all possible new cation-anion combinations. This involves exchanging the anions of the two reactants (an ion interchange or double displacement reaction).

    • The possible new ion combinations are $KNO<em>3$ (potassium nitrate) and $BaCrO</em>4$ (barium chromate).

    • $KNO_3$ is typically aqueous (soluble in water) based on solubility rules.

    • $BaCrO_4$ is typically a yellow solid (insoluble, thus forms a precipitate).

• Using Solubility Rules

  • Predicting precipitates requires knowledge of solubility rules, which are empirical generalizations based on extensive experimental observations.

  • Soluble solid: A solid that dissolves significantly in water, generally meaning its solubility is greater than approximately $0.1 ext{ M}$ (moles per liter) at room temperature.

  • Insoluble solid: A solid that does not dissolve significantly in water, generally meaning its solubility is less than approximately $0.01 ext{ M}$. Such a solid forms a precipitate when the ions are mixed.

  • Slightly soluble solid: Has limited solubility, falling between soluble and insoluble, but is often treated as insoluble for practical purposes in predicting precipitate formation.

• How to Predict Precipitates When Solutions of Two Ionic Compounds Are Mixed

1. Write the Reactants as Ions: Accurately write the formulas for the reactants as they exist in solution, remembering that soluble ionic compounds (salts) exist as separated ions. This helps visualize all the ions present.

2. Identify Possible Product Combinations: Consider all possible solids that could form by exchanging the anions between the two original cations. This is the essence of a double displacement reaction, producing two new ionic compounds.

3. Apply Solubility Rules: Use established solubility rules (e.g., typically provided in a table like Table 7.1) to decide whether either of the newly formed ionic compounds is insoluble. If one or both are insoluble, that compound(s) will form a precipitate. If both are soluble, no precipitate forms, and therefore no reaction occurs.

Describing Reactions in Aqueous Solutions

• Types of Equations for Reactions in Aqueous Solutions

  • Molecular Equation: Shows the complete formulas of all reactants and products as if they were intact molecules or compounds, even if they exist as dissociated ions in solution. It provides the overall stoichiometry but doesn't accurately represent the species present in solution.

    • Example: $K<em>2CrO</em>4(aq) + Ba(NO<em>3)</em>2(aq) <br>ightarrow BaCrO<em>4(s) + 2KNO</em>3(aq)$

  • Complete Ionic Equation: Represents all substances that are strong electrolytes (soluble ionic compounds, strong acids, strong bases) as separated, dissociated ions. Weak electrolytes, non-electrolytes (like water), and insoluble solids are written in their molecular or undissociated forms.

    • Example: $2K^+(aq) + CrO<em>4^{2-}(aq) + Ba^{2+}(aq) + 2NO</em>3^-(aq) <br>ightarrow BaCrO<em>4(s) + 2K^+(aq) + 2NO</em>3^-(aq)$

    • Spectator ions: Ions that are present in solution both before and after the reaction and do not undergo any chemical change (i.e., they do not participate directly in the reaction). They appear on both sides of the complete ionic equation in identical forms and are ultimately removed to get the net ionic equation. In the example, $K^+$ and $NO_3^-$ are spectator ions.

  • Net Ionic Equation: An equation that includes only those components (ions or molecules) that are directly involved in the chemical transformation. Spectator ions are removed (cancelled out) from the complete ionic equation. This equation provides the most fundamental and accurate representation of the actual chemical change occurring.

    • Example: $Ba^{2+}(aq) + CrO<em>4^{2-}(aq) ightarrow BaCrO</em>4(s)$

• Concept Check 4: Write the molecular, complete ionic, and net ionic equations for the reaction between cobalt(II) chloride and sodium hydroxide.

  • Molecular Equation: $CoCl<em>2(aq) + 2NaOH(aq) ightarrow Co(OH)</em>2(s) + 2NaCl(aq)$

  • Complete Ionic Equation: $Co^{2+}(aq) + 2Cl^-(aq) + 2Na^+(aq) + 2OH^-(aq) <br>ightarrow Co(OH)_2(s) + 2Na^+(aq) + 2Cl^-(aq)$

  • Net Ionic Equation: $Co^{2+}(aq) + 2OH^-(aq) <br>ightarrow Co(OH)_2(s)$

Reactions That Form Water: Acids and Bases

• Arrhenius Acids and Bases

  • Strong acid: A strong electrolyte that completely ionizes (or dissociates) in water, producing $H^+$ ions (protons). In reality, these $H^+$ ions immediately react with water to form hydronium ions ($H_3O^+$), but for simplicity, $H^+$ is often used.

    • Examples:

      • Hydrochloric acid: $HCl(aq) <br>ightarrow H^+(aq) + Cl^-(aq)$ (or $HCl(aq) + H<em>2O(l) ightarrow H</em>3O^+(aq) + Cl^-(aq))$

      • Nitric acid: $HNO<em>3(aq) ightarrow H^+(aq) + NO</em>3^-(aq)$

      • Sulfuric acid: $H<em>2SO</em>4(aq) <br>ightarrow H^+(aq) + HSO_4^-(aq)$ (only the first proton dissociates completely; the second is weak).

  • Strong base: A substance that completely dissociates or ionizes to produce hydroxide ions (OH^-$) in water. These are typically soluble metal hydroxides, particularly those of Group 1 and the heavier Group 2 metals.

    • Most common examples are sodium hydroxide ($NaOH$) and potassium hydroxide ($KOH$).

    • Examples:

      • NaOH(s)
        ightarrow Na^+(aq) + OH^-(aq)$</p></li><li><p>$KOH(s)
        ightarrow K^+(aq) + OH^-(aq)$</p></li></ul></li></ul></li><li><p><strong>Products of Strong Acid and Strong Base Reaction</strong>: The reaction between a strong acid and a strong base always produces water and an ionic compound known as a <strong>salt</strong>. This type of reaction is also called a <strong>neutralization reaction</strong> because the acidic and basic properties of the reactants are neutralized.</p><ul><li><p>The <strong>net ionic equation</strong> for such reactions is always:$H^+(aq) + OH^-(aq)
        ightarrow H_2O(l)$. This highlights the core chemical change, the formation of water from its constituent ions.</p></li><li><p>$H^+$is the acidic species, and$OH^-$is the basic species directly involved in water formation.</p></li></ul></li></ul></li><li><p><strong>Summary of Strong Acids and Strong Bases</strong></p><ul><li><p>Common strong acids:$HCl$(hydrochloric acid),$HBr$(hydrobromic acid),$HI$(hydroiodic acid),$HNO3$(nitric acid),$HClO4$(perchloric acid), and$H2SO4$(sulfuric acid, for its first proton dissociation). These acids completely dissociate (ionize) in water into an$H^+ ion and a conjugate anion.

      • Strong bases: Metal hydroxide compounds that are very soluble in water. Key examples include hydroxides of Group 1 metals (LiOH, NaOH, KOH, RbOH, CsOH) and heavier Group 2 metals (Ca(OH)2, Sr(OH)2, Ba(OH)2). They completely dissociate into separated metal cations and hydroxide ions ($OH^-$) when dissolved.

      • The net ionic equation for the reaction of any strong acid with any strong base always expresses the production of water: H^+(aq) + OH^-(aq)
        ightarrow H_2O(l)$.</p></li><li><p>The other product, a salt, remains dissolved in water as separated ions and can be obtained as a solid by evaporating the water.</p></li><li><p>The reaction of$H^+$and$OH^-$is fundamentally an acid-base neutralization reaction.</p></li></ul></li><li><p><strong>Concept Check 5</strong>: The net ionic equation for the reaction of$HNO3$and$LiOH$is$H^+ + OH^- ightarrow H2O$. (This correctly shows the neutralization of a strong acid by a strong base).</p></li></ul><h5 id="15310153-7ffa-4dbb-97df-492a7d4d62c6" data-toc-id="15310153-7ffa-4dbb-97df-492a7d4d62c6" collapsed="false" seolevelmigrated="true">Reactions of Metals with Nonmetals (Oxidation-Reduction)</h5><ul><li><p><strong>Oxidation-Reduction Reaction (Redox)</strong>: A chemical reaction that fundamentally involves a <strong>transfer of electrons</strong> between atoms or ions. When an atom loses electrons, its oxidation state increases, and it is said to be <strong>oxidized</strong>; when an atom gains electrons, its oxidation state decreases, and it is said to be <strong>reduced</strong>.</p><ul><li><p>Example:$2Mg(s) + O_2(g) ightarrow 2MgO(s)$</p><ul><li><p>In this reaction,$Mg$metal undergoes oxidation (loses$2e^-$per atom) to form$Mg^{2+}$ions, while$O_2$gas undergoes reduction (each$O$atom gains$2e^-$) to form$O^{2-}$ions within the ionic compound$MgO$. The electrons are transferred from magnesium to oxygen.</p></li></ul></li><li><p>Reactions between a metal and a nonmetal invariably involve a transfer of electrons from the metal (which typically has a lower electronegativity and tends to lose electrons) to the nonmetal (which typically has a higher electronegativity and tends to gain electrons).</p></li></ul></li><li><p><strong>Concept Check 6</strong>: In the reaction of Al with$Fe2O3$, metal Al loses$3e^-$(is oxidized) and$Fe^{3+}$in$Fe2O3$gains these$3e^-$(is reduced) to become$Fe$metal.</p></li><li><p><strong>Characteristics of Oxidation-Reduction Reactions</strong></p><ul><li><p>Any reaction between a metal and a nonmetal can be confidently classified as an oxidation-reduction reaction, as it intrinsically involves electron transfer and changes in oxidation states.</p></li><li><p>Reactions between two nonmetals can also be oxidation-reduction reactions. These are often recognized by the presence of$O_2$as a reactant or product (e.g., combustion or synthesis reactions involving oxygen) or by tracking changes in the oxidation states of the elements involved. When two nonmetals react, the compound formed is typically covalent, not ionic, but electron sharing is unequal, leading to changes in oxidation states.</p></li></ul></li></ul><h5 id="4a4bdf2e-5772-4606-8aac-1f06322c9f13" data-toc-id="4a4bdf2e-5772-4606-8aac-1f06322c9f13" collapsed="false" seolevelmigrated="true">Ways to Classify Reactions</h5><ul><li><p><strong>Driving Forces for a Reaction</strong> (reiterated, with common reaction type):</p><ul><li><p>Formation of a solid (Precipitation Reaction)</p></li><li><p>Formation of water (Acid-Base Reaction/Neutralization)</p></li><li><p>Transfer of electrons (Oxidation-Reduction Reaction)</p></li><li><p>Formation of a gas (Gas Evolution Reaction)</p></li></ul></li><li><p><strong>Precipitation Reaction</strong>: Characterized by the formation of an insoluble solid (precipitate) when two aqueous solutions are mixed. This is a common and observable driving force.</p><ul><li><p>Often a <strong>double-displacement reaction</strong> (also known as a <strong>metathesis reaction</strong> or <strong>ion-exchange reaction</strong>):$AB + CD
        ightarrow AD + CB$. In this type of reaction, the cations and anions of two ionic compounds effectively exchange partners.</p></li><li><p>Example:$K2CrO4(aq) + Ba(NO3)2(aq)
        ightarrow BaCrO4(s) + 2KNO3(aq)$</p></li></ul></li><li><p><strong>Acid-Base Reaction</strong>: Involves an$H^+$ion (from an acid) reacting directly with an$OH^-$ion (from a base) to form water. The net result is a neutralization of acidic and basic properties.</p><ul><li><p>Net ionic equation for strong acid-strong base:$H^+(aq) + OH^-(aq)
        ightarrow H_2O(l)$</p></li><li><p>Example:$HCl(aq) + KOH(aq)
        ightarrow H_2O(l) + KCl(aq)$</p></li></ul></li><li><p><strong>Oxidation-Reduction Reaction</strong>: Characterized by the transfer of electrons between reactants, leading to changes in the oxidation states of the involved atoms. Many chemical reactions, especially those involving elements, fall into this category.</p><ul><li><p>Example:$2Li(s) + F_2(g) ightarrow 2LiF(s)$</p><ul><li><p>Here, lithium (a metal) loses electrons to become$Li^+$(oxidized), and fluorine (a nonmetal) gains electrons to become$F^-$(reduced).</p></li></ul></li></ul></li><li><p><strong>Formation of a Gas</strong>: Can be a distinct driving force, often visible as bubbles. While many gas-forming reactions are also oxidation-reduction reactions, not all are. Some involve decomposition or acid-carbonate reactions.</p><ul><li><p>Often seen in a <strong>single-replacement reaction</strong>:$A + BC
        ightarrow B + AC$. In this type of reaction, a more reactive element replaces another element in a compound.</p></li><li><p>Example:$Zn(s) + 2HCl(aq) ightarrow H2(g) + ZnCl2(aq)$. In this reaction, solid zinc replaces hydrogen from hydrochloric acid.</p><ul><li><p>This is an oxidation-reduction reaction:$Zn$(oxidation state 0) loses electrons to become$Zn^{2+}$(oxidized), while$H^+$ions (oxidation state +1) gain electrons to form uncharged$H_2$gas (oxidation state 0, reduced).</p></li></ul></li></ul></li></ul><h5 id="ea58cb0b-ac22-4243-8120-1f4a4f369743" data-toc-id="ea58cb0b-ac22-4243-8120-1f4a4f369743" collapsed="false" seolevelmigrated="true">Other Ways to Classify Reactions</h5><ul><li><p><strong>Combustion Reactions</strong></p><ul><li><p>Involve a rapid reaction with oxygen ($O_2$), producing energy (often in the form of heat and light) so quickly that a flame results. They are always a special class of oxidation-reduction reactions, commonly involving organic compounds.</p></li><li><p>Example: Complete combustion of methane:$CH4(g) + 2O2(g)
        ightarrow CO2(g) + 2H2O(g)$. Here, carbon and hydrogen are oxidized by oxygen.</p></li></ul></li><li><p><strong>Synthesis (Combination) Reactions</strong></p><ul><li><p>Involve the formation of a more complex compound from two or more simpler, often elemental, materials. These reactions usually combine elements or simple compounds to form a more complex one.</p></li><li><p>Example: Formation of carbon dioxide from its elements:$C(s) + O2(g) ightarrow CO2(g)$. This is also an oxidation-reduction reaction because the oxidation states of C and O change.</p></li></ul></li><li><p><strong>Decomposition Reactions</strong></p><ul><li><p>Occur when a single compound is broken down (decomposed) into two or more simpler substances. This process often requires an input of energy, such as heat, light, or electricity.</p></li><li><p>Example: Electrolysis of water:$2H2O(l) ightarrow 2H2(g) + O_2(g)$. This is an oxidation-reduction reaction where the compound water is broken down into its elemental components, involving changes in oxidation states.</p></li></ul></li><li><p><strong>Concept Check 7</strong>: The reaction$2H2 + O2
        ightarrow 2H_2O$$ can be classified as:

        • a) oxidation-reduction reaction (Hydrogen is oxidized, Oxygen is reduced)

        • b) synthesis reaction (Two simpler substances combine to form a more complex one)

        • c) combustion reaction (Reaction with oxygen, often producing heat/light)

        • d) All of the above (This reaction fits all three classifications).