General Chemistry: Chemical Quantities & Aqueous Reactions Study Notes

Chemical Quantities & Aqueous Reactions

Electrolytes and Nonelectrolytes

Electrolytes are substances that, when dissolved in water, produce a solution that conducts electricity due to the presence of ions. They are classified as strong or weak based on the extent of ionization.

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

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

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

Examples:

• HNO3 (Nitric acid) is a strong electrolyte.

• NH3 (Ammonia) is a weak electrolyte.

• CH3OH (Methanol) is a nonelectrolyte.

Solubility Rules and Precipitation Reactions

Solubility rules help predict whether an ionic compound will dissolve in water. Precipitation reactions occur when two aqueous solutions combine to form an insoluble product (precipitate).

• Soluble Compounds: Most salts of Na+, K+, NH4+, NO3-, and CH3COO- are soluble.

• Insoluble Compounds: Most carbonates (CO32-), phosphates (PO43-), sulfides (S2-), and hydroxides (OH-) are insoluble except with alkali metals and NH4+.

Example: Mixing solutions of AgNO3 and NaCl produces a white precipitate of AgCl.

Writing Chemical Equations in Aqueous Reactions

Reactions in aqueous solutions can be represented in three forms:

1. Molecular Equation: Shows all reactants and products as compounds.

2. Complete Ionic Equation: Shows all strong electrolytes as ions.

3. Net Ionic Equation: Shows only the species that actually change during the reaction.

Example: Reaction between Na2SO4 and BaCl2:

• Molecular: $\mathrm{Na_2SO_4(aq) + BaCl_2(aq) \rightarrow 2NaCl(aq) + BaSO_4(s)}$

• Complete Ionic: $2\mathrm{Na^+(aq) + SO_4^{2-}(aq) + Ba^{2+}(aq) + 2Cl^-(aq) \rightarrow 2Na^+(aq) + 2Cl^-(aq) + BaSO_4(s)}$

• Net Ionic: $\mathrm{Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)}$

Table: Predicting Precipitation and Representing Particles

The following table summarizes the prediction of precipitation and the representation of particles in solution:

Reactants	Products	Precipitate or not?	Representing particles
CaCl2(aq) + Na2SO4(aq)	CaSO4(s) + 2NaCl(aq)	Yes	Ca2+, SO42- (solid); Na+, Cl- (aqueous)
Pb(NO3)2(aq) + KI(aq)	PbI2(s) + 2KNO3(aq)	Yes	Pb2+, I- (solid); K+, NO3- (aqueous)
Na2CO3(aq) + HCl(aq)	2NaCl(aq) + H2O(l) + CO2(g)	No precipitate (gas evolution)	Na+, Cl- (aqueous); H2O (liquid); CO2 (gas)
AgNO3(aq) + NaCl(aq)	AgCl(s) + NaNO3(aq)	Yes	Ag+, Cl- (solid); Na+, NO3- (aqueous)

Additional info: Table entries inferred based on standard precipitation reactions and solubility rules.

Solubility of Compounds in Water

Some compounds are insoluble in water and form precipitates when their ions combine in solution.

• Examples of insoluble compounds: BaSO4, CaCO3, AgI

• Examples of soluble compounds: NaCl, KNO3

Writing Molecular, Complete Ionic, and Net Ionic Equations

For reactions in aqueous solution, it is important to distinguish between the different forms of equations:

• Molecular Equation: Shows all reactants and products as compounds.

• Complete Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only the ions and molecules directly involved in the reaction.

Example: Reaction between Pb(NO3)2(aq) and Na2SO4(aq):

• Molecular: $\mathrm{Pb(NO_3)_2(aq) + Na_2SO_4(aq) \rightarrow PbSO_4(s) + 2NaNO_3(aq)}$

• Complete Ionic: $\mathrm{Pb^{2+}(aq) + 2NO_3^-(aq) + 2Na^+(aq) + SO_4^{2-}(aq) \rightarrow PbSO_4(s) + 2Na^+(aq) + 2NO_3^-(aq)}$

• Net Ionic: $\mathrm{Pb^{2+}(aq) + SO_4^{2-}(aq) \rightarrow PbSO_4(s)}$

Oxidation Numbers and Redox Reactions

Oxidation numbers are assigned to atoms to keep track of electron transfer in redox reactions. The sum of oxidation numbers in a neutral compound is zero; in a polyatomic ion, it equals the ion's charge.

• Rules for Assigning Oxidation Numbers:

  • Elements in their standard state: 0

  • Group 1 metals: +1; Group 2 metals: +2

  • Fluorine: -1; Oxygen: usually -2; Hydrogen: +1 with nonmetals, -1 with metals

  • Sum of oxidation numbers in a compound = 0; in a polyatomic ion = ion charge

Example: Assign oxidation numbers to each atom in SO42-:

• Oxygen: -2 each, total -8

• Sulfur: x; x + (-8) = -2 → x = +6

Identifying Oxidizing and Reducing Agents

In redox reactions, the oxidizing agent is reduced (gains electrons), and the reducing agent is oxidized (loses electrons).

• Example: In the reaction $2NO + 2NO_2 + 2H_2O \rightarrow 3N_2O + 4H^+$, identify the oxidizing and reducing agents by assigning oxidation numbers and tracking changes.

Summary Table: Common Oxidation Numbers

Element/Ion	Common Oxidation Number
K	+1
LiF	Li: +1, F: -1
SO42-	S: +6, O: -2
CO	C: +2, O: -2
CO32-	C: +4, O: -2
P4	P: 0

Additional info: Table entries inferred based on standard oxidation number rules.