General Chemistry: Electrolytes, Stoichiometry, and Chemical Equations Study Guide

Electrolytes and Dissociation in Water

Definitions and Properties

Electrolytes are substances that, when dissolved in water, produce a solution that conducts electricity. This is due to the presence of ions in the solution.

• Strong Electrolytes: Substances that completely dissociate into ions in water (e.g., strong acids, strong bases, and most salts).

• Weak Electrolytes: Substances that partially dissociate into ions in water (e.g., weak acids and weak bases).

• Nonelectrolytes: Substances that do not produce ions in solution (e.g., most molecular compounds like sugar).

Example: Acetic acid (CH3COOH) is a weak acid and a weak electrolyte. It partially dissociates in water:

$\mathrm{CH_3COOH(aq) \rightleftharpoons CH_3COO^-(aq) + H^+(aq)}$

Key Point: The dissociation of a weak acid in water is a reversible process and does not go to completion.

Solubility and Precipitation

Solubility describes how well a substance dissolves in water. Some ionic compounds are soluble, while others are insoluble and form precipitates.

• Soluble: Dissolves in water to form ions.

• Insoluble: Does not dissolve significantly in water; may form a solid precipitate.

Example: Na3PO4 is soluble, while AgCl is insoluble in water.

Stoichiometry and Chemical Calculations

Mole-to-Mass Conversions

Stoichiometry involves calculations based on balanced chemical equations. One common calculation is converting moles of a substance to grams using its molar mass.

• Molar Mass: The mass of one mole of a substance, expressed in g/mol.

• Conversion Formula:

$\text{mass (g)} = \text{moles} \times \text{molar mass (g/mol)}$

Example: To find the mass of 5.71 moles of nickel (Ni):

$5.71\ \text{mol Ni} \times \frac{58.69\ \text{g Ni}}{1\ \text{mol Ni}} = 335\ \text{g Ni}$

Empirical and Molecular Formulas

The empirical formula gives the simplest whole-number ratio of atoms in a compound, while the molecular formula gives the actual number of atoms of each element in a molecule.

• To determine the empirical formula, use the masses of elements produced from combustion analysis to find the mole ratio.

• To find the molecular formula, divide the compound's molar mass by the empirical formula mass.

Example: A compound with a molar mass of 78.1 g/mol produces 2.311 g CO2 and 0.4729 g H2O upon combustion. Calculations show the empirical formula is CH, and the molecular formula is C6H6 (benzene).

Stoichiometry in Chemical Reactions

Stoichiometric calculations allow you to determine the amount of reactants or products in a chemical reaction.

• Use the coefficients from the balanced equation to set up mole ratios.

• Convert between grams and moles as needed.

Example: To find the mass of NO2 produced from a given mass of HNO3:

$4.30\ \text{g HNO}_3 \times \frac{1\ \text{mol HNO}_3}{63.02\ \text{g HNO}_3} \times \frac{2\ \text{mol NO}_2}{3\ \text{mol HNO}_3} \times \frac{46.01\ \text{g NO}_2}{1\ \text{mol NO}_2} = 2.1\ \text{g NO}_2$

Additional info: The actual calculation in the image gives 4.7 g NO2, which matches the correct stoichiometry for the reaction.

Chemical Equations and Ionic Equations

Complete Ionic Equations

A complete ionic equation shows all of the ions present in a reaction, especially for reactions in aqueous solution.

• Strong electrolytes are written as ions.

• Weak electrolytes and insoluble compounds are written as molecules or solids.

Example: The reaction between NaOH(aq) and CuSO4(aq):

$2\text{Na}^+(aq) + 2\text{OH}^-(aq) + \text{Cu}^{2+}(aq) + \text{SO}_4^{2-}(aq) \rightarrow \text{Cu(OH)}_2(s) + 2\text{Na}^+(aq) + \text{SO}_4^{2-}(aq)$

Additional info: The net ionic equation would be $\text{Cu}^{2+}(aq) + 2\text{OH}^-(aq) \rightarrow \text{Cu(OH)}_2(s)$.

Balancing Chemical Equations

Balanced chemical equations have the same number of atoms of each element on both sides. This is essential for stoichiometric calculations.

Example: Neutralization of phosphoric acid (H3PO4) with potassium hydroxide (KOH):

$\text{H}_3\text{PO}_4(aq) + 3\text{KOH}(aq) \rightarrow \text{K}_3\text{PO}_4(aq) + 3\text{H}_2\text{O}(l)$

Solubility Table (Inferred)

The following table summarizes the solubility of some common ionic compounds in water:

Compound	Solubility in Water
Na3PO4	Soluble
AgCl	Insoluble
Ca(OH)2	Slightly Soluble
HCl	Soluble (strong acid)

Additional info: Solubility rules can be used to predict whether a precipitate will form in a reaction.