chapter 4 lec 1

Stoichiometry of Chemical Reactions

General Properties of Aqueous Solutions

Aqueous solutions are central to many chemical reactions and biological processes. Understanding their properties is essential for predicting reaction outcomes and behaviors.

• Solution: A homogeneous mixture of two or more components. The solvent is present in greater amount, while the solute is present in lesser amount.

• Types of Solutions: Solutions can be gaseous (e.g., air), liquid (e.g., salt dissolved in water), or solid (e.g., brass).

• Aqueous Solution: Water acts as the solvent, and the solute is a solid or liquid dissolved in water.

Example: Salt (NaCl) dissolved in water forms an aqueous solution.

Electrolytic Properties

Electrolytic properties describe the ability of a solution to conduct electricity, which depends on the presence of ions.

• Electrolyte: A substance that conducts electricity when dissolved in water due to the formation of ions. Example: NaCl in water.

• Nonelectrolyte: A substance that does not conduct electricity in water, typically molecular compounds that do not form ions. Examples: Sugar, ethylene glycol, ethyl alcohol.

• Strong Electrolyte: Dissociates completely in water, producing many ions and conducting electricity well. Examples: HCl, NaOH.

• Weak Electrolyte: Dissociates only partially in water, producing few ions and conducting electricity poorly. Examples: Acetic acid, ammonia.

Experiment: Immersing platinum electrodes in pure water does not light a bulb due to poor conductivity. Adding an electrolyte (e.g., NaCl) allows current to flow and lights the bulb.

Strong Electrolytes

• Strong Acids (100% dissociation): HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO4

• Strong Bases (100% dissociation): LiOH, NaOH, KOH, RbOH, CsOH, Ba(OH)2, Sr(OH)2, Ca(OH)2

• All soluble salts are strong electrolytes.

Weak Electrolytes

• Weak Acids: HF, acetic acid, citric acid, tartaric acid

• Weak Bases: Ammonia, pyridine, methylamine, diethylamine

• These substances only partially dissociate in water, resulting in equilibrium between ions and molecules.

Example Equation:

• $\mathrm{CH_3COOH~(acetic~acid)~+~H_2O~\rightleftharpoons~H_3O^+~+~CH_3COO^-}$

Only a small fraction of acetic acid molecules dissociate, making it a weak electrolyte.

Writing and Balancing Chemical Equations

Balancing Chemical Equations

Chemical equations must be balanced to obey the law of conservation of mass. This means the number of atoms of each element must be the same on both sides of the equation.

• Example: Decomposition of potassium chlorate:

  • Unbalanced: $\mathrm{KClO_3~\rightarrow~KCl~+~O_2}$

  • Balanced: $\mathrm{2KClO_3~\rightarrow~2KCl~+~3O_2}$

• Example: Combustion of octane:

  • Unbalanced: $\mathrm{C_8H_{18}~+~O_2~\rightarrow~CO_2~+~H_2O}$

  • Balanced: $\mathrm{2C_8H_{18}~+~25O_2~\rightarrow~16CO_2~+~18H_2O}$

Balancing equations often requires adjusting coefficients to ensure atom counts are equal.

Classifying Chemical Reactions

Precipitation Reactions

Precipitation reactions result in the formation of an insoluble product, called a precipitate, which separates from the solution. These reactions typically involve ionic compounds.

• Example: $\mathrm{AgNO_3~(aq)~+~HCl~(aq)~\rightarrow~AgCl~(s)~+~H_3O^+~(aq)~+~NO_3^-~(aq)}$

• Example: $\mathrm{Pb(NO_3)_2~(aq)~+~2KBr~(aq)~\rightarrow~PbBr_2~(s)~+~2K^+~(aq)~+~2NO_3^-~(aq)}$

Solubility Guidelines for Ionic Compounds in Water

Solubility rules help predict whether a precipitate will form when two solutions are mixed.

Soluble Ionic Compounds	Exceptions
All Group I compounds & NH4+	None
All nitrates (NO3-), bicarbonates (HCO3-), chlorates (ClO3-), acetates (C2H3O2-)	None
All halides (Cl-, Br-, I-)	Halides of Ag+, Hg22+, Pb2+
All sulfates (SO42-)	Sulfates of Sr2+, Ba2+, Hg22+, Pb2+

Insoluble Ionic Compounds	Exceptions
All carbonates (CO32-), phosphates (PO43-), chromates (CrO42-), sulfides (S2-)	Group I, NH4+, Ca2+, Sr2+, Ba2+
All hydroxides (OH-)	Group I & NH4+, Ca2+, Sr2+, Ba2+

Solubility: The maximum amount of solute that will dissolve in a given quantity of solvent at a given temperature.

Soluble: Large amount dissolves; slightly soluble: much less dissolves; insoluble: essentially none dissolves.

Molecular, Ionic, and Net Ionic Equations

Chemical reactions in solution can be represented in three ways:

• Molecular Equation: Shows compounds as whole units.

  • Example: $\mathrm{Pb(NO_3)_2~(aq)~+~H_2SO_4~(aq)~\rightarrow~PbSO_4~(s)~+~2HNO_3~(aq)}$

• Ionic Equation: Shows all ions present in solution.

  • Example: $\mathrm{Pb^{2+}~(aq)~+~2NO_3^-~(aq)~+~2H_3O^+~(aq)~+~SO_4^{2-}~(aq)~\rightarrow~PbSO_4~(s)~+~2H_3O^+~(aq)~+~2NO_3^-~(aq)}$

• Net Ionic Equation: Shows only the species that change during the reaction; spectator ions are omitted.

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

Acid-Base Reactions

Arrhenius Definition of Acids and Bases

The Arrhenius theory defines acids and bases based on their behavior in water.

• Acid: Substance that ionizes in water to produce H+ (or H3O+) ions.

• Base: Substance that ionizes in water to produce OH- ions.

General Properties of Acids

• Sour taste

• Change litmus from blue to red

• React with metals (Zn, Mg, Fe) to produce hydrogen gas:

  • $\mathrm{Zn~(s)~+~2HCl~(aq)~\rightarrow~ZnCl_2~(aq)~+~H_2~(g)}$

• React with carbonates and bicarbonates to produce CO2:

  • $\mathrm{2HCl~(aq)~+~BaCO_3~(s)~\rightarrow~BaCl_2~(aq)~+~H_2O~(l)~+~CO_2~(g)}$

• Aqueous solutions conduct electricity

General Properties of Bases

• Bitter taste

• Slippery feel (e.g., soaps)

• Change litmus from red to blue

• Aqueous solutions conduct electricity

Strong Acids and Bases

• Strong Acids: Monoprotic acids (HCl, HBr, HI, HNO3, HClO3, HClO4), and one diprotic acid (H2SO4).

• Polyprotic Acids: Acids with more than one proton (e.g., H2CO3, H3PO4), but only H2SO4 is strong.

• Strong Bases: LiOH, NaOH, KOH, CsOH, RbOH, Ba(OH)2, Sr(OH)2, Ca(OH)2 (the last three are dibasic).

Brønsted-Lowry Theory of Acids and Bases

The Brønsted-Lowry theory expands the definition of acids and bases beyond aqueous solutions.

• Brønsted Acid: Proton donor

• Brønsted Base: Proton acceptor

• These definitions apply in any solvent system, not just water.

Example:

• $\mathrm{HNO_3~(aq)~+~H_2O~(l)~\rightleftharpoons~H_3O^+~(aq)~+~NO_3^-~(aq)}$ Here, H2O acts as a base, accepting a proton.

• $\mathrm{NaOH~(s)~+~H_2O~(l)~\rightleftharpoons~Na^+~(aq)~+~OH^-~(aq)}$ Here, H2O acts as an acid, donating a proton.

Polyprotic Acids

• Diprotic Acid: Can donate two protons in two steps.

  • Step 1: $\mathrm{H_2SO_4~(aq)~+~H_2O~(l)~\rightleftharpoons~H_3O^+~(aq)~+~HSO_4^-~(aq)}$

  • Step 2: $\mathrm{HSO_4^-~(aq)~+~H_2O~(l)~\rightleftharpoons~H_3O^+~(aq)~+~SO_4^{2-}~(aq)}$

• Triprotic Acid: Can donate three protons in three steps.

  • Step 1: $\mathrm{H_3PO_4~(aq)~+~H_2O~(l)~\rightleftharpoons~H_3O^+~(aq)~+~H_2PO_4^-~(aq)}$

  • Step 2: $\mathrm{H_2PO_4^-~(aq)~+~H_2O~(l)~\rightleftharpoons~H_3O^+~(aq)~+~HPO_4^{2-}~(aq)}$

  • Step 3: $\mathrm{HPO_4^{2-}~(aq)~+~H_2O~(l)~\rightleftharpoons~H_3O^+~(aq)~+~PO_4^{3-}~(aq)}$

In these reactions, water acts as the base, accepting protons.

Brønsted Base Example

• $\mathrm{NH_3~(aq)~+~H_2O~(l)~\rightleftharpoons~NH_4^+~(aq)~+~OH^-~(aq)}$ Ammonia is a Brønsted base, accepting a proton from water.

Acid-Base Neutralization

Neutralization reactions occur when an acid reacts with a base, producing a salt and water.

• General Equation: $\mathrm{acid~+~base~\rightarrow~salt~+~water}$

• Example: $\mathrm{HClO_4~(aq)~+~KOH~(aq)~\rightarrow~KClO_4~(aq)~+~H_2O~(l)}$

• Ionic Equation: $\mathrm{H^+~(aq)~+~ClO_4^-~(aq)~+~K^+~(aq)~+~OH^-~(aq)~\rightarrow~K^+~(aq)~+~ClO_4^-~(aq)~+~H_2O~(l)}$

• Net Ionic Equation: $\mathrm{H^+~(aq)~+~OH^-~(aq)~\rightarrow~H_2O~(l)}$

For weak acids, the net ionic equation includes the undissociated acid:

• Example: $\mathrm{HF~(aq)~+~LiOH~(aq)~\rightarrow~LiF~(aq)~+~H_2O~(l)}$

• Net Ionic Equation: $\mathrm{HF~(aq)~+~OH^-~(aq)~\rightarrow~F^-~(aq)~+~H_2O~(l)}$

Neutralization always produces water and a salt.

Additional info: The notes provide foundational concepts for stoichiometry, solution chemistry, electrolytes, and acid-base reactions, including solubility rules and the classification of acids and bases. These are essential for understanding chemical reactions in aqueous solutions and predicting reaction outcomes.