Chemistry 121 Chapter 4 Notes - Stoichiometry of Chemical Reactions

CHEMISTRY 121 CHAPTER 4 NOTES: STOICHIOMETRY OF CHEMICAL REACTIONS

Chemical Equations

• Definition: Chemical equations represent chemical reactions.

• Components of Equations:

  1. Reactants: Substances present at the start of the reaction.

  2. Products: Substances formed during the reaction.

  3. Mole Ratios: Ratios of the number of particles involved in the reaction.

• Example:

  • $N<em>2(g) + 3H</em>2(g) \rightarrow 2NH_3(g)$

  • Here, $N2$ and $H2$ are reactants, while $NH_3$ is the product.

  • For every molecule of $N2$, 3 molecules of $H2$ are needed, yielding 2 molecules of $NH_3$.

  • For every mole of $N2$, 3 moles of $H2$ react to produce 2 moles of $NH_3$.

• Phase Indicators:

  • $(g)$: gas phase

  • $(l)$: liquid phase

  • $(s)$: solid phase

  • $(aq)$: aqueous solution

• Special Conditions: Conditions such as temperature, pressure, catalysts, or heating are indicated above or below the arrow in the equation.

The Law of Conservation of Matter

• Statement: Matter is neither created nor destroyed in a chemical reaction.

• Implications:

  • Atoms retain identity; number of atoms of each element must remain constant on both sides to balance the equation.

  • Balanced equations are shown using the smallest whole-number coefficients.

  • To balance an equation, change coefficients, not subscripts (changing subscripts changes the substance's identity).

  • Chemical formulas must reflect the substance in the form it exists, such as $H2$ for hydrogen and $O2$ for oxygen.

• Balancing Methodology:

  • Start with elements that appear in only one reactant and product.

  • Balance free elements last.

  • Polyatomic ions that do not change throughout the reaction can be treated as units.

Example Problems for Balancing Equations

• Sodium Metal and Water Reaction:

  • Reaction:

  • $Na + H<em>2O \rightarrow NaOH + H</em>2$

• Calcium Nitrate and Sodium Phosphate Reaction:

  • Reaction results in a precipitate, calcium phosphate, and an aqueous sodium nitrate solution.

Reversible Reactions

• Definition and Representation:

  • Reversible reactions can proceed in both forward (reactants to products) and reverse directions (products to reactants).

  • Represented by double arrow: $\rightleftharpoons$.

• Chemical Equilibrium: This state occurs when the forward and reverse reactions happen at equal rates.

• Equilibrium Conditions:

  1. The reaction must be reversible.

  2. The rate of the forward process equals the rate of the reverse process.

• Examples of Product-Favored and Reactant-Favored Reactions:

  • Product-favored: equilibrium heavily favors product formation with few reactants.

  • Reactant-favored: equilibrium favors reactant formation with few products.

  • Equilibrium can involve substantial amounts of both reactants and products.

Characteristics of Aqueous Solutions

• Classification of Solutes:

  • Electrolytes: Substances that conduct electricity in water.

  • Nonelectrolytes: Substances that do not conduct electricity in water.

• Electrolytes Types:

  1. Strong Electrolytes: Fully dissociate into ions, excellent conductors.

  2. Weak Electrolytes: Partially dissociate, poor conductors.

• Dissociation vs. Ionization:

  • Dissociation: Breakdown of ionic compounds into ions.

  • Ionization: Breakdown of molecular compounds into ions.

Strong and Weak Electrolytes

• Strong Electrolytes Include:

  1. Strong acids

  2. Strong soluble bases

  3. Soluble ionic compounds

• Example of Strong Acid Dissociation:

  • $HNO<em>3(aq) \rightarrow H^+(aq) + NO</em>3^-(aq)$

• Monoprotic Acid: Can lose one proton (e.g. $HNO_3$).

• Polyprotic Acid: Can lose more than one proton (e.g. $H<em>2CO</em>3$ which can dissociate as follows:

  • $H<em>2CO</em>3 \rightleftharpoons H^+ + HCO_3^-$

  • $HCO<em>3^- \rightleftharpoons H^+ + CO</em>3^{2-}$

• Properties of Acids:

  • Sour taste, turns litmus red, reacts with carbonates and metals.

• Common Strong Acids:

  • $HCl$, $HBr$, $HI$, $HNO3$, $H2SO4$, $HClO4$, $HClO_3$.

Weak Acids and Their Ionization

• Example of Weak Acid Ionization:

  • $CH<em>3COOH(aq) \rightleftharpoons CH</em>3COO^-(aq) + H^+(aq)$

  • Only a small fraction ionizes in solution; e.g. vinegar is about 0.5% ionized.

• Organic acids commonly contain the carboxylic acid group (-COOH).

Bases

• Definition: Substances that yield $OH^-$ ions in water or accept $H^+$ ions during a reaction (Arrhenius and Bronsted-Lowry definitions, respectively).

• Classification:

  • Strong Bases: Soluble metal hydroxides (e.g. $LiOH, NaOH$).

  • Weak Bases: Soluble molecular substances that produce $OH^-$ in water (e.g. $NH_3$).

• Example of Strong Base Reaction:

  • $BaO + H<em>2O \rightarrow Ba(OH)</em>2$.

Solubility Guidelines for Ionic Compounds

1. Common inorganic acids are soluble in water.

2. All Group IA metal compounds and ammonium ($NH_4^+$) are soluble.

3. Compounds with nitrate, acetate, chlorate, and perchlorate are soluble.

4. Common chloride compounds are soluble except $AgCl$, $Hg2Cl2$, and $PbCl_2$.

5. Common sulfate compounds are soluble except for $PbSO4$, $BaSO4$, and $HgSO_4$.

6. Most metal hydroxides are insoluble, with some exceptions like $Ca(OH)_2$.

7. Compounds containing carbonate, phosphate, and arsenate are largely insoluble except with Group IA metals and $NH_4^+$.

Different Methods to Represent Chemical Reactions

1. Formula Unit/Molecular Equations: Show complete formulas for all compounds.

   • Example: $Zn(s) + CuSO<em>4(aq) \rightarrow Cu(s) + ZnSO</em>4(aq)$.

2. Total Ionic Equations: Depict predominant forms in aqueous solution; strong electrolytes dissociate into ions.

3. Net Ionic Equations: Show only those substances involved in the reaction, omitting spectator ions.

Classifying Chemical Reactions

• Oxidation-Reduction (Redox) Reactions:

  • Oxidation: Loss of electrons.

  • Reduction: Gain of electrons.

  • The substance that loses electrons is oxidized (reducing agent), and the one that gains electrons is reduced (oxidizing agent).

Oxidation Numbers Rules

1. Pure elements have oxidation number 0.

2. Monatomic ion's oxidation number equals its charge.

3. Total oxidation number in a compound is zero; for polyatomic ions, it equals the ion's charge.

4. Fluorine always has an oxidation number of -1.

5. Group IA metals are +1; Group IIA metals are +2.

6. Hydrogen's number varies, being -1 with metals and +1 with nonmetals.

Balancing Redox Reactions Using Half-Reaction Method

1. Write unbalanced reaction & assign oxidation numbers.

2. Determine oxidation and reduction substances.

3. Write half-reactions, balancing until oxygen/hydrogen through water and $H^+$ in acidic or $OH^-$ in basic solutions.

4. Combine balanced half-reactions.

Types of Chemical Reactions

Combination Reactions

• Definition: Two or more reactants combine to form a single product.

• Forms:

  1. Element + Element → Compound

  2. Compound + Element → Compound

Decomposition Reactions

• Definition: One reactant breaks apart into two or more products.

• General Layout: $A
  ightarrow B + C$.

Displacement Reactions

• Definition: An element displaces another from a compound.

Activity Series of Metals

• Metals higher in the activity series can displace metals lower in series from their compounds.

• Example Activity Series (from most to least active):

  • K, Na, Ca, Mg, Al, C, Zn, Fe, Sn, Pb, H, Cu, Ag, Au.

Percent Yields and Limiting Reactant Concept

• Percent Yield Calculation:

  • $Percent~Yield = \left(\frac{Actual~Yield}{Theoretical~Yield}\right) \times 100$

• Limiting Reactant: Reactant that limits the amount of product formed in a reaction.

Examples in Quantitative Analysis and Combustion Analysis

• Example of Gravimetric Analysis to find mass percentage of analyte:

  • AgNO3 reacts with Cl^− to form AgCl precipitate; calculate mass percentage of Cl.

• Combustion Example: Hydrocarbon combustion yields $CO2$ and $H2O$; empirical formula can be derived from masses of combustion products.