Premed · Premed · General Chemistry 1
Lecture 4: Reactions in Aqueous Solution
General Chemistry I
Learning Objectives
By the end of this lecture, students will be able to:
- Define electrolytes and nonelectrolytes and predict electrical conductivity of solutions
- Write molecular, complete ionic, and net ionic equations
- Use solubility rules to predict whether a precipitate forms
- Identify and balance acid-base neutralization reactions
- Assign oxidation numbers and identify oxidation-reduction (redox) reactions
- Identify the oxidizing and reducing agents in a redox reaction
- Perform solution stoichiometry calculations using molarity
Lecture Content
I. Aqueous Solutions and Electrolytes
An aqueous solution forms when a solute dissolves in water, the solvent. Water is an excellent solvent because of its polarity, which enables it to dissolve many ionic and polar compounds. Substances that dissociate into ions when dissolved in water are called electrolytes, and their solutions conduct electricity. Strong electrolytes dissociate completely (100%) into ions and include strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4), strong bases (Group 1A hydroxides such as NaOH and KOH, along with Ca(OH)2, Ba(OH)2, and Sr(OH)2), and soluble ionic compounds like NaCl and KNO3. Weak electrolytes dissociate only partially, establishing an equilibrium between intact molecules and ions; examples include weak acids (HF, CH3COOH, H2CO3, H3PO4) and weak bases (NH3 and amines). Nonelectrolytes dissolve without producing ions, so their solutions do not conduct electricity. Common nonelectrolytes include molecular compounds such as glucose (C6H12O6), ethanol (C2H5OH), and sucrose.
<image>A three-panel diagram illustrating electrolyte behavior. Panel A (Strong electrolyte -- NaCl): a beaker of water with fully separated Na+ and Cl- ions scattered throughout, a lightbulb circuit connected to electrodes in solution glowing brightly. Panel B (Weak electrolyte -- CH3COOH): a beaker with a mixture of intact CH3COOH molecules and a few CH3COO- and H+ ions, lightbulb glowing dimly. Panel C (Nonelectrolyte -- C6H12O6): a beaker with intact glucose molecules dispersed in water, lightbulb not glowing. Each panel labeled with dissociation equation below.</image>
II. Solubility Rules for Ionic Compounds
A set of solubility rules allows predictions about whether a given ionic compound will dissolve in water. All compounds containing Group 1A cations (Li+, Na+, K+) or NH4+ are soluble, as are all nitrates (NO3-) and acetates (CH3COO-). Chlorides, bromides, and iodides are generally soluble, except those of Ag+, Pb2+, and Hg2^2+. Sulfates are generally soluble, except for those of Ba2+, Pb2+, Sr2+, and (partially) Ca2+. On the insoluble side, hydroxides are insoluble except for those of Group 1A, Ca2+, Ba2+, and Sr2+. Carbonates (CO3^2-) and phosphates (PO4^3-) are insoluble except when paired with Group 1A cations or NH4+. Sulfides (S^2-) follow the same pattern, being insoluble except with Group 1A, Group 2A, and NH4+.
III. Precipitation Reactions
A precipitation reaction occurs when two aqueous ionic compounds are mixed and produce an insoluble product, called a precipitate. The solubility rules guide these predictions. Such reactions can be written in three forms. The molecular equation shows the complete formulas of all reactants and products: AgNO3(aq) + NaCl(aq) -> AgCl(s) + NaNO3(aq). The complete ionic equation breaks all strong electrolytes into their constituent ions: Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) -> AgCl(s) + Na+(aq) + NO3-(aq). The net ionic equation eliminates spectator ions -- those that appear unchanged on both sides -- leaving only the species that participate in the reaction: Ag+(aq) + Cl-(aq) -> AgCl(s). In this example, Na+ and NO3- are the spectator ions.
IV. Acid-Base Reactions (Neutralization)
Under the Arrhenius definition, an acid produces H+ (or H3O+) in aqueous solution, while a base produces OH-. When an acid reacts with a base, the result is a neutralization reaction that produces a salt and water: HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l). The net ionic equation for a strong acid-strong base neutralization is simply H+(aq) + OH-(aq) -> H2O(l). When a strong acid reacts with a strong base, the reaction goes to completion. If a weak acid or weak base is involved, it should be written in molecular form (undissociated) in ionic equations. Some acid reactions also produce gases. For instance, carbonates react with acids to release carbon dioxide: CaCO3(s) + 2 HCl(aq) -> CaCl2(aq) + H2O(l) + CO2(g). Similarly, sulfites react with acids to produce sulfur dioxide: Na2SO3 + 2 HCl -> 2 NaCl + H2O + SO2(g).
V. Oxidation-Reduction (Redox) Reactions
Oxidation-reduction reactions involve the transfer of electrons between species. Oxidation is defined as the loss of electrons, corresponding to an increase in oxidation number, while reduction is the gain of electrons, corresponding to a decrease in oxidation number. Two popular mnemonics capture this: OIL RIG (Oxidation Is Loss, Reduction Is Gain) and LEO GER (Lose Electrons Oxidation, Gain Electrons Reduction). The oxidizing agent is the substance that is reduced -- it accepts electrons and thereby causes another species to be oxidized. The reducing agent is the substance that is oxidized -- it donates electrons and thereby causes another species to be reduced.
A. Rules for Assigning Oxidation Numbers
A systematic set of rules governs oxidation number assignments. Free (uncombined) elements have an oxidation number of 0 (for example, Fe, O2, S8). Monatomic ions have oxidation numbers equal to their charge (Na+ = +1, Cl- = -1). Hydrogen is +1 in most compounds but -1 in metal hydrides such as NaH. Oxygen is -2 in most compounds, -1 in peroxides (H2O2), and -1/2 in superoxides (KO2). Fluorine is always -1. The sum of oxidation numbers in a neutral compound must equal 0, and in a polyatomic ion it must equal the ion's charge.
<image>A diagram illustrating electron transfer in a redox reaction between zinc metal and copper(II) sulfate solution. Left side: a shiny zinc strip being placed into a blue CuSO4 solution. Right side: the zinc strip now coated with reddish-brown copper metal, solution color faded. Below: the half-reactions written out -- Zn(s) -> Zn2+(aq) + 2e- (oxidation, labeled "reducing agent") and Cu2+(aq) + 2e- -> Cu(s) (reduction, labeled "oxidizing agent"). Arrows show the flow of electrons from Zn to Cu2+. Oxidation numbers are shown changing: Zn from 0 to +2, Cu from +2 to 0.</image>
VI. Types of Redox Reactions
Redox reactions come in several forms. Combination (synthesis) reactions take the form A + B -> AB, as in 2 Mg + O2 -> 2 MgO. Decomposition reactions reverse this pattern: AB -> A + B, exemplified by 2 HgO -> 2 Hg + O2. Single displacement (replacement) reactions have an element replacing another in a compound (A + BC -> AC + B). A metal can displace another metal, as when Zn(s) + CuSO4(aq) -> ZnSO4(aq) + Cu(s), or displace hydrogen from an acid, as in Zn(s) + 2 HCl(aq) -> ZnCl2(aq) + H2(g). Whether a displacement occurs is predicted by the activity series -- a more active metal displaces a less active one. Combustion reactions involve a substance reacting with O2 to produce heat and light; hydrocarbon combustion follows the general pattern CxHy + O2 -> CO2 + H2O.
VII. The Activity Series of Metals
The activity series ranks metals by their tendency to be oxidized (lose electrons). A metal higher in the series can displace a metal lower in the series from a compound. The series, from most active to least active, runs: Li > K > Ba > Ca > Na > Mg > Al > Zn > Cr > Fe > Ni > Sn > Pb > H2 > Cu > Ag > Pt > Au. Metals above H2 in the series can displace hydrogen from acids, while metals below H2 cannot react with common acids to produce hydrogen gas.
VIII. Molarity and Solution Stoichiometry
Molarity (M) expresses concentration as moles of solute per liter of solution: M = n / V. Dilution -- the process of adding solvent to reduce concentration -- preserves the number of moles of solute, giving the relationship M1V1 = M2V2. For solution stoichiometry problems, use molarity and volume to calculate moles (moles = M x V in liters), then apply the mole ratios from the balanced equation. For example, to determine how many mL of 0.100 M AgNO3 are needed to react with 25.0 mL of 0.150 M NaCl, first calculate the moles of NaCl: 0.0250 L x 0.150 mol/L = 0.00375 mol. Since the reaction has a 1:1 mole ratio, 0.00375 mol of AgNO3 is required, which corresponds to a volume of 0.00375 / 0.100 = 0.0375 L, or 37.5 mL.
IX. Titrations
A titration is a technique for determining the concentration of an unknown solution by reacting it with a solution of known concentration, called the standard or titrant. The equivalence point is reached when stoichiometrically equivalent amounts of acid and base have been mixed. An indicator -- a substance that changes color at or near the equivalence point -- signals this moment visually. At the equivalence point of an acid-base titration, the relationship moles acid x stoichiometric factor = moles base holds true. For a 1:1 acid-base reaction, this simplifies to M_acid x V_acid = M_base x V_base.
<image>A labeled diagram of an acid-base titration setup. A burette filled with NaOH solution (labeled with concentration) is clamped to a ring stand, with a stopcock at the bottom. Below, an Erlenmeyer flask contains the acid solution with a few drops of phenolphthalein indicator (initially colorless). A magnetic stir plate beneath the flask shows the stir bar. An inset shows the moment of the endpoint: the solution turning pink as the last drop of NaOH is added, with the equation HCl + NaOH -> NaCl + H2O written alongside.</image>
