Ch4

Interactive General Chemistry 2.0: Chemical Reactions and Aqueous Solutions

Chapter Outline

• Section 4.1: Chemical Equations

  • Identify the parts of a balanced chemical equation, including phases.

  • Write complete and balanced chemical equations based on either chemical symbols or word descriptions.

• Section 4.2: Types of Chemical Reactions

  • Recognize and describe the five basic types of chemical reactions.

  • List and describe the driving forces for reactions in aqueous solution.

• Section 4.3: Compounds in Aqueous Solution

  • Describe and write equations representing the dissociation of ionic compounds upon dissolution in water.

  • Identify compounds as being strong electrolytes, weak electrolytes, or nonelectrolytes from chemical formulas.

• Section 4.4: Precipitation Reactions

  • Apply solubility guidelines to predict the formation of a precipitate in reactions of ionic compounds in aqueous solution.

  • Write and interpret ionic and net ionic equations for precipitation reactions.

• Section 4.5: Acid-Base Reactions

  • Predict the products of acid-base reactions in aqueous solution.

  • Write and interpret ionic and net ionic equations for acid-base reactions.

  • Predict the products of gas-forming aqueous reactions.

• Section 4.6: Oxidation States and Redox Reactions

  • Assign oxidation states to elements within compounds and polyatomic ions.

  • Identify redox reactions using oxidation states.

  • Define and apply the terms oxidizing agent and reducing agent.

• Section 4.7: Predicting the Products of Redox Reactions

  • Write balanced equations for synthesis and decomposition reactions.

  • Apply the activity series to predict the product of single-replacement reactions.

Section 4.1: Chemical Equations

• Parts of a Balanced Chemical Equation

  • A balanced chemical equation includes reactants, products, and their physical states.

• Writing Balanced Chemical Equations

  • Chemical equations can be represented in words (word equations) or symbols (chemical formulas).

Information in a Chemical Equation

• The notation for physical states of reactants and products is indicated in parentheses:

  • Solid (s), Liquid (l), Gas (g), Aqueous solution (aq).

  • Example:

  • 1 mol aqueous magnesium iodide, MgI₂(aq), reacts with 2 mol aqueous silver nitrate, AgNO₃(aq), to produce 2 mol solid silver iodide, AgI(s), and 1 mol aqueous magnesium nitrate, Mg(NO₃)₂(aq).

Reaction Conditions

• Conditions necessary for reactions may be listed above or below the arrow in a chemical equation.

  • Example: Heat is needed to decompose magnesium carbonate into solid magnesium oxide and carbon dioxide gas. Without heat, the reaction does not occur.

Balancing Chemical Equations

• Law of Conservation of Mass: All atoms present at the beginning of a reaction must be present at the end.

  • In a balanced equation, the quantity of each type of atom must equal the corresponding quantity on the opposite side of the equation.

• Steps to Balance Chemical Equations

  1. If polyatomic ions are present as reactants and products, balance them as units (e.g., SO₄²⁻).

  2. Start balancing the first element that appears in one reactant and one product and proceed left to right.

  3. Balance any remaining elements and change previously determined coefficients if necessary.

  4. Verify the balanced equation to ensure equal numbers of each type of atom on both sides.

  5. Ensure coefficients are in the smallest whole-number ratio; adjust them if any coefficients are fractions.

Section 4.2: Types of Chemical Reactions

• Five Basic Types of Chemical Reactions

  • Combination (synthesis) reactions

  • Decomposition reactions

  • Single-replacement reactions

  • Double-replacement reactions

  • Combustion reactions

Combination Reactions

• These reactions involve simple reactants combining to form a more complex product.

  • Example: Iron (Fe) reacting with oxygen (O₂) to form iron(III) oxide (Fe₂O₃).

Decomposition Reactions

• Occurs when a single reactant breaks down into less complex products.

  • Example: Water (H₂O) decomposes into hydrogen gas (H₂) and oxygen gas (O₂) through electrical energy.

Single-Replacement Reactions

• Involves an element replacing another in a compound, resulting in a new compound and a new element.

  • Example: Zinc (Zn) reacting with hydrochloric acid (HCl), where Zn replaces H.

Double-Replacement Reactions

• Two ionic compounds exchange ions forming two new compounds.

  • Example: Lead(II) nitrate (Pb(NO₃)₂) reacting with potassium iodide (KI).

Acid-Base Reactions

• A specific case of double-replacement reactions where an acid (often starting with H) reacts with a base (often ending with OH).

Combustion Reactions

• Involves the rapid combination of a substance (usually hydrocarbon) with oxygen, producing carbon dioxide and water.

Summary of Reaction Types

Section 4.3: Compounds in Aqueous Solution

• Dissociation of Ionic Compounds

  • Ionic compounds break apart into ions when dissolved in water.

  • Strong electrolytes dissociate completely, while weak electrolytes only partially.

Aqueous Solutions

• A compound that dissolves in water is termed soluble; if it remains solid, it is insoluble.

  • Example notations:

  • NaCl(aq) signifies sodium chloride dissolves in water to form an aqueous solution.

  • AgCl(s) indicates silver chloride does not dissolve and remains solid when mixed with water.

Electrolytes

• Substances that conduct electricity when dissolved in water due to mobile hydrated ions are known as electrolytes.

  • Strong electrolytes (e.g., ionic compounds) dissociate 100%.

Molecular Compounds in Water

• Most molecular compounds form solutions that do not conduct electricity (nonelectrolytes).

  • Weak acids ionize in water to produce ions, thus acting as electrolytes.

Strong Acids

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• Strong acids ionize completely in water, producing solutions that are strong electrolytes.

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• Table 4.3: Strong Acids
  

  Name	Formula	Ions
  Hydrochloric acid	HCl	H⁺(aq) + Cl⁻(aq)
  Hydrobromic acid	HBr	H⁺(aq) + Br⁻(aq)
  Hydroiodic acid	HI	H⁺(aq) + I⁻(aq)
  Nitric acid	HNO₃	H⁺(aq) + NO₃⁻(aq)
  Perchloric acid	HClO₄	H⁺(aq) + ClO₄⁻(aq)
  Chloric acid	HClO₃	H⁺(aq) + ClO₃⁻(aq)
  Sulfuric acid	H₂SO₄	H⁺(aq) + HSO₄⁻(aq)
  Bases

  • Strong Bases are ionic compounds with hydroxide ions that also dissociate completely in water to form strong electrolytes.

  • Weak Bases are usually molecular compounds like ammonia, which react with water to produce OH⁻ ions to a small extent, acting as weak electrolytes.

  Section 4.4: Precipitation Reactions

  • Use solubility guidelines to predict which ionic compounds will precipitate in aqueous solution.

  • Write and interpret ionic and net ionic equations.

  Solubility Guidelines (1 of 2)

  • Group 1 elements and ammonium (NH₄⁺) compounds are soluble.

  • Nitrates (NO₃⁻), chlorates (ClO₃⁻), perchlorates (ClO₄⁻), and acetates (C₂H₃O₂⁻) are soluble.

  • Chlorides (Cl⁻), bromides (Br⁻), and iodides (I⁻) are generally soluble, except for those of Ag⁺, Pb²⁺, and Hg₂²⁺.

  • Carbonates (CO₃²⁻), sulfites (SO₃²⁻), phosphates (PO₄³⁻), and chromates (CrO₄²⁻) are generally insoluble, except for those with Group 1 cations.

  Solubility Guidelines (2 of 2)

  • Hydroxides (OH⁻) and sulfides (S²⁻) are mostly insoluble except with Group 1 elements and Ba²⁺.

  • Most sulfates (SO₄²⁻) are soluble except for those with Ca²⁺, Sr²⁺, and Ba²⁺.

  Steps to Determine Precipitate Formation

  1. Determine possible new ionic compounds that could form and write their formulas based on ion charges.

  2. Use solubility guidelines to check which products are insoluble.

  3. If both products are soluble, no reaction occurs. If one is insoluble, then a precipitation reaction occurs.

  Ionic Equations

  • Ionic equations describe reactions in electrolyte solutions more accurately than total equations showing complete formulae of reactants and products.

  Net Ionic Equations

  • Spectator ions that do not participate in the reaction are removed from the ionic equation to yield the net ionic equation.

  Section 4.5: Acid-Base Reactions

  • Predict product formation for acid-base reactions in aqueous solution, and write respective ionic and net ionic equations.

  Completing and Balancing Acid-Base Reactions

  • The hydrogen from the acid combines with hydroxide from the base to produce water, while the remaining ions form a salt.

  • These reactions generally show no visible change in the mixture and release heat.

  Gas-Forming Reactions

  • Certain unstable compounds spontaneously react in aqueous solution to form gases, such as carbonic acid (H₂CO₃) decomposing into CO₂ and H₂O, with CO₂ escaping as bubbles.

  Section 4.6: Oxidation States and Redox Reactions

  • Oxidation States: Also known as oxidation numbers used to track electron transfers in redox reactions.

  Rules for Assigning Oxidation States

  • Free elements have an oxidation state of 0.

    • Example: Na = 0, Cl₂ = 0 in 2 Na(s) + Cl₂(g).

  • Monatomic ions have oxidation states equal to their charge (e.g., Na = +1, Cl = -1 in NaCl).

  • The total oxidation states of all atoms in a compound must sum to 0, while in polyatomic ions, they must equal the ion's charge.

  Additional Rules for Nonmetals

  • Group I metals have oxidation state +1; Group II metals have +2.

  • Nonmetals have specific oxidation states (e.g., F: -1, O: -2, etc., based on their group).

  Identifying Redox Reactions

  • Redox reactions involve electron transfer: oxidation refers to losing electrons (increase in oxidation state), while reduction involves gaining electrons (decrease in oxidation state).

  • Example: In the synthesis of aluminum chloride (AlCl₃), Al is oxidized from 0 to +3, and Cl is reduced from 0 to -1.

  Summary of Redox Terminology

  Term	Definition
  Oxidation	Increase in oxidation state
  Reduction	Decrease in oxidation state
  Reducing agent	The substance that is oxidized
  Oxidizing agent	The substance that is reduced

  Section 4.7: Predicting the Products of Redox Reactions

  • Balanced equations for synthesis and decomposition redox reactions must be written.

  • Culminates in applying the activity series to predict outcomes for single-replacement reactions.

  Single-Replacement Reactions

  • Also redox reactions where reactive metals can easily oxidize and give up electrons.

  Activity Series

  Reducing Agent	Activity as Element	Oxidizing Agent	Activity as Ion
  Groups 1–2 metals	Most active	Ions of Groups 1–2 metals	Least active
  Al	Al³⁺	Mn	Mn²⁺
  Zn	Zn²⁺	Cr	Cr³⁺
  Fe	Fe²⁺	Ni	Ni²⁺
  Sn	Sn²⁺	Pb	Pb²⁺
  H	H⁺	Cu	Cu²⁺
  Ag	Ag⁺	Au	Au³⁺
  	Least active		Most active

  Reactions of Metals with Acids

  • Acids release H⁺ in aqueous solutions, where neutral metals above H in activity series can displace H⁺ to produce H₂(g).

  • Metals below H in the activity series will not react with acids.