Chapter 9: Chemical Reactions in Aqueous Solutions

General Properties of Aqueous Solutions

A solution is defined as a homogeneous mixture of two or more substances.

• Solvent: The substance present in the largest amount (measured in moles).
• Solute: All other substances present in the mixture.
• Solubility: A substance that dissolves in a specific solvent is described as being soluble in that solvent.

Unless specified otherwise, within the context of this study material, "solution" refers specifically to an aqueous solution, where the solvent is water.

Electrolytes and Nonelectrolytes

Electrolytes

An electrolyte is a substance that, when dissolved in water, yields a solution capable of conducting electricity. This property results from the process of dissociation, where the substance breaks apart into its constituent ions.

• Ionization: This term specifically refers to molecular compounds that form ions when they dissolve in water.

Nonelectrolytes

A nonelectrolyte is a substance that dissolves in water to yield a solution that does not conduct electricity. In these cases, the molecules (such as sucrose, $C_{12}H_{22}O_{11}$) remain completely intact upon dissolving and do not form ions.

Strong and Weak Electrolytes

Strong Electrolytes

A strong electrolyte is a compound that dissociates completely ($100\%$) into ions in solution. Categories include:

• Water-soluble ionic compounds.
• Strong Acids (listed in Table 9.1).
• Strong Bases.

Table 9.1: The Strong Acids

*Note: Sulfuric acid has two ionizable hydrogen atoms. However, only the first ionization occurs completely producing one $H^{+}$ and one $HSO_4^{-}$ ion per molecule. The second ionization occurs only to a very small extent.

Weak Electrolytes

A weak electrolyte produces ions upon dissolving but exists in the solution predominantly as unionized molecules.

• Weak Acids: Such as acetic acid or ascorbic acid.
• Weak Bases: Such as ammonia.

Dynamic Chemical Equilibrium: Denoted by a double arrow ($\rightleftharpoons$), this state occurs when the forward and reverse reaction rates are equal, and the concentration of reactants and products remains constant.

Classification Logic (Slide 11 Flowchart)

1. Ionic: The compound is a strong electrolyte.
2. Molecular:
   • Acid: If it is one of the seven strong acids, it is a strong electrolyte. If not, it is a weak electrolyte.
   • Base: Generally a weak electrolyte (unless it is a metal hydroxide, which makes it ionic/strong).
   • Neither: The compound is a nonelectrolyte.

Worked Example 9.1:

• Sucrose ($C_{12}H_{22}O_{11}$) and Fructose ($C_{6}H_{12}O_{6}$): Molecular compounds, neither acids nor bases; classified as nonelectrolytes.
• Sodium Citrate ($Na_3C_6H_5O_7$) and Potassium Citrate ($K_3C_6H_5O_7$): Contain metal cations; ionic compounds classified as strong electrolytes.
• Ascorbic Acid ($H_2C_6H_6O_6$): An acid not found on the list of seven strong acids; classified as a weak electrolyte.

Precipitation Reactions and Solubility Guidelines

Precipitation Reactions

A precipitation reaction is a chemical reaction that results in the formation of an insoluble product known as a precipitate.

Polarity and Hydration

Water is an effective solvent for ionic compounds because it is a polar molecule. Due to electron distribution, the oxygen atom is partially negative ($\delta^{-}$) as it attracts the hydrogen atoms' electrons, while the hydrogen atoms are partially positive ($\delta^{+}$).

Hydration is the process where water molecules surround and remove individual ions from an ionic solid to allow it to dissolve.

Solubility Guidelines (Table 9.2 & 9.3 Summary)

Solubility is the maximum amount of solute that dissolves in a given quantity of solvent at a specific temperature.

• Table 9.2: Soluble Compounds:

  • Compounds containing Alkali metal ions ($Li^{+}$, $Na^{+}$, $K^{+}$, $Rb^{+}$, $Cs^{+}$) and Ammonium ($NH_4^{+}$).
  • Nitrates ($NO_3^{-}$), Acetates ($CH_3COO^{-}$), Bicarbonates ($HCO_3^{-}$), Chlorates ($ClO_3^{-}$), and Perchlorates ($ClO_4^{-}$).
  • Halides ($Cl^{-}$, $Br^{-}$, $I^{-}$) except those containing $Ag^{+}$, $Hg_2^{2+}$, and $Pb^{2+}$.
  • Sulfates ($SO_4^{2-}$) except those containing $Ag^{+}$, $Ca^{2+}$, $Sr^{2+}$, $Ba^{2+}$, $Hg_2^{2+}$, and $Pb^{2+}$.

• Table 9.3: Insoluble Compounds:

  • Carbonates ($CO_3^{2-}$), Phosphates ($PO_4^{3-}$), Chromates ($CrO_4^{2-}$), and Sulfides ($S^{2-}$) except those containing alkali metal ions or ammonium.
  • Hydroxides ($OH^{-}$) except those containing alkali metal ions and the $Ba^{2+}$ ion.

Writing Equations for Aqueous Reactions

1. Molecular Equations

Compounds are represented by their chemical formulas as if they exist as intact molecules or formula units.

• Metathesis (Double Replacement): Reactions where cations in two ionic compounds exchange anions.

2. Ionic Equations

Compounds that exist predominantly as ions are represented as dissociated ions.

• Example: Aqueous $Na_2SO_4$ and $Ba(OH)_2$:     $2Na^{+}(aq) + SO_4^{2-}(aq) + Ba^{2+}(aq) + 2OH^{-}(aq) \rightarrow BaSO_4(s) + 2Na^{+}(aq) + 2OH^{-}(aq)$

3. Net Ionic Equations

Includes only the species participating in the reaction.

• Spectator Ions: Ions appearing on both sides of the equation that do not participate in the reaction. These are canceled out.

Steps to determine equations:

1. Write and balance the molecular equation, predicting products by cation exchange.
2. Write the ionic equation by separating strong electrolytes into constituent ions.
3. Identify and cancel spectator ions for the net ionic equation.
4. If all reactants and products are strong electrolytes, no reaction occurs (no net ionic equation).

Acid-Base Reactions

Definitions

• Arrhenius Acid: Ionizes in water to produce $H^{+}$.
• Arrhenius Base: Dissociates in water to produce $OH^{-}$.
• Brønsted Acid: A proton ($H^{+}$) donor.
• Brønsted Base: A proton ($H^{+}$) acceptor.

Strong Bases (Table 9.4)

Strong bases are strong electrolytes (hydro-oxides of Group 1A and heavy Group 2A):

• $LiOH$, $NaOH$, $KOH$, $RbOH$, $CsOH$
• $Ca(OH)_2$, $Sr(OH)_2$, $Ba(OH)_2$

Acids

• Monoprotic: One ionizable proton (e.g., $HCl$).
• Polyprotic: More than one acidic hydrogen atom (e.g., $H_2SO_4$ is diprotic). These lose protons stepwise.
  • Step 1: $H_2SO_4(aq) \rightarrow H^{+}(aq) + HSO_4^{-}(aq)$
  • Step 2: $HSO_4^{-}(aq) \rightleftharpoons H^{+}(aq) + SO_4^{-}(aq)$

Bases

• Monobasic: Produce one mole of $OH^{-}$ per mole of compound ($NaOH$).
• Dibasic: Produce two moles of $OH^{-}$ per mole of compound ($Ba(OH)_2$).

Neutralization

A reaction between an acid and a base, typically producing water and a salt. The common net ionic equation for strong acid-strong base reactions is: $H^{+}(aq) + OH^{-}(aq) \rightarrow H_2O(l)$

Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons between reactants.

• Oxidation: Loss of electrons.
• Reduction: Gain of electrons.
• Half-reactions: A redox reaction is the sum of an oxidation half-reaction and a reduction half-reaction.

Oxidation Numbers (Oxidation States)

The oxidation number is the hypothetical charge an atom would have if electrons were transferred completely.

Rules for Assigning Oxidation Numbers (Table 9.5):

1. Elemental form: Oxidation number is $0$.
2. Sum of oxidation numbers: Must equal the overall charge of the species (neutral molecule = $0$; polyatomic ion = its charge).
3. Reliable assignments:
   • Fluorine: Always $-1$ in compounds.
   • Group 1 Metals: $+1$; Group 2 Metals: $+2$.
   • Hydrogen: Always $+1$ (except in metal hydrides like $LiH$ or $CaH_2$, where it is $-1$).
   • Oxygen: Usually $-2$ (except in peroxides like $H_2O_2$, where it is $-1$, or superoxides like $KO_2$, where it is $-1/2$).
   • Group 17 (excluding F): Usually $-1$ unless combined with oxygen or fluorine.

The Activity Series (Table 9.6)

This is a list of metals (and hydrogen) arranged by decreasing ease of oxidation.

• Active Metals: Located at the top (e.g., $Li$, $K$, $Ba$).
• Noble Metals: Located at the bottom (e.g., $Ag$, $Pt$, $Au$).
• Displacement Rule: An element in the series will be oxidized by the ions of any element appearing below it in the table.

Balancing Redox Equations

Equations must achieve both mass and charge balance.

• Half-Reaction Method: Balance individual half-reactions for atoms and charge, then multiply by integers to equalize electron counts before adding the reactions together.

Types of Redox Reactions

• Combination: e.g., $N_2(g) + 3H_2(g) \rightarrow 2NH_3(g)$.
• Decomposition: e.g., $2NaH(s) \rightarrow 2Na(s) + H_2(g)$.
• Disproportionation: One element is simultaneously oxidized and reduced (e.g., decomposition of $H_2O_2$).
• Combustion: e.g., $CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l)$.

Concentration of Solutions

Molarity ($M$)

Molar concentration is the number of moles of solute per $dm^3$ of solution. $M = \frac{\text{moles of solute}}{\text{liters of solution}}$$\text{moles} = M \times V \text{ (in } dm^3)$

Dilution

The process of preparing a less concentrated solution from a stock solution. The moles of solute remain constant. $M_1V_1 = M_2V_2$

Serial Dilution: A series of sequential dilutions used to prepare increasingly dilute solutions.

The pH Scale

Acidity depends on the concentration of hydronium ions ($H_3O^{+}$). pH is the negative base-10 logarithm of that concentration: $pH = -\log_{10}[H_3O^{+}]$ At $25^\circ\text{C}$, pure water has a pH of $7.0$.

Benchmark pH Values:

• $[H_3O^{+}] = 1.0 \times 10^{-1}\,mol\,dm^{-3} \rightarrow pH = 1.00$
• $[H_3O^{+}] = 1.0 \times 10^{-7}\,mol\,dm^{-3} \rightarrow pH = 7.00$
• $[H_3O^{+}] = 1.0 \times 10^{-13}\,mol\,dm^{-3} \rightarrow pH = 13.00$

Chemical Analysis

Gravimetric Analysis

An analytical technique based on the measurement of mass. It is highly accurate but only applicable to reactions that go to completion ($100\%$ yield).

• Example: Determining percent chloride in a sample by precipitating it with excess $AgNO_3$ to form $AgCl$.

Acid-Base Titrations

A volumetric technique used to determine quantities in neutralization reactions using burets.

• Equivalence Point: The point where the acid has been completely neutralized.
• Indicator: A substance that changes color in response to acidic or basic media.
• End Point: The point where the indicator color change occurs; chosen to be as close to the equivalence point as possible.
• Standardization: The process of determining the exact concentration of a solution (e.g., using Potassium Hydrogen Phthalate, KHP, to standardize $NaOH$).

Worked Example 9.15 (Standardization):

• $25.49\,cm^3$ of $NaOH$ neutralizes $0.7137\,g$ of KHP.
• Molar mass of KHP = $204.2\,g\,mol^{-1}$.
• $\text{moles of KHP} = \frac{0.7137\,g}{204.2\,g\,mol^{-1}} = 0.003495\,mol$.
• Since the ratio is $1:1$, moles of $NaOH = 0.003495\,mol$.
• $M = \frac{0.003495\,mol}{0.02549\,dm^3} = 0.1371\,mol\,dm^{-3}$.