Study Notes: Chemical Reactions - Chapter 5

Chemical Reactions - Chapter 5 Summary

5.1 Introduction to Chemical Reactions

  • A. General Features of Physical and Chemical Changes

    • Physical Change: Alters the physical state of a substance without changing its composition.

    • Chemical Change: Converts one substance into another. Chemical reactions involve:

    • Breaking bonds in the reactants (starting materials).

    • Forming new bonds in the products.

5.1B Writing Chemical Equations

  • 1. Definition of a Chemical Equation:

    • A chemical equation uses chemical formulas and other symbols to show what reactants are involved in a reaction and what products are formed.

    • Reactants are written on the left.

    • Products are written on the right.

    • Coefficients indicate the number of molecules of a given element or compound participating in the reaction.

  • 2. Law of Conservation of Mass:

    • States that atoms cannot be created or destroyed in a chemical reaction.

    • Coefficients ensure the equation is balanced; a balanced equation has the same number of atoms of each element on both sides.

  • 3. Symbols Used in Chemical Equations:

    • Reaction arrow (→)

    • Heat (Δ)

    • Solid (s)

    • Liquid (l)

    • Gas (g)

    • Aqueous solution (aq)

5.2 Balancing Chemical Equations

  • Step [1]: Write the unbalanced equation with correct formulas.

    • Example: For propane and oxygen: C3H8 + O2 → CO2 + H_2O.

    • Important: Subscripts cannot be changed to balance; modifying subscripts alters the identity of the compound.

  • Step [2]: Balance the equation with coefficients for one element at a time (C first, then H, then O).

  • Step [3]: Verify that the smallest set of whole numbers is used to balance the equation.

  • Torres’ Tips and Tricks for Balancing Reactions:

    • Start with metal atoms.

    • Leave oxygen and hydrogen till the end.

    • Keep polyatomic ions together.

    • If coefficients become large, start over.

    • If it's a back-and-forth process, restart the attempt.

5.3 Types of Reactions

  • The majority of chemical reactions can be classified into 7 categories:

    • Combination

    • Decomposition

    • Single Replacement

    • Double Replacement

    • Combustion

    • Oxidation and Reduction (Section 5.4)

    • Acid-Base (Chapter 9)

5.3A Combination and Decomposition

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

  • Decomposition Reaction: A single reactant converts to two or more products.

5.3B Replacement Reactions

  • Single Replacement Reaction: One element replaces another in a compound, forming a different element and a new compound as products.

  • Double Replacement Reaction: Two compounds exchange parts to form two new compounds.

    • Examples:

    • 2 NaCl + Br2 → 2 NaBr + Cl2 (Br replaces Cl).

    • Fe + CuSO4 → FeSO4 + Cu (Fe replaces Cu).

    • AgNO3 + NaCl → AgCl + NaNO3 (Ag and Na exchange).

    • HCl + NaOH → H_2O + NaCl (H and Na exchange).

5.3 Combustion Reactions

  • Definition: A combustion reaction occurs when a hydrocarbon reacts with oxygen, producing carbon dioxide and water.

    • General form: CxHy + O2 → CO2 + H_2O.

5.4 Oxidation and Reduction

  • A. General Features of Oxidation-Reduction Reactions:

    • Oxidation: Loss of electrons from an atom.

    • Reduction: Gain of electrons by an atom.

    • These processes occur together in a redox reaction, which transfers electrons from one element to another.

  • Oxidized and Reduced Components:

    • Example: In Zn + Cu^{2+} → Zn^{2+} + Cu, Zn is oxidized (loses electrons), Cu is reduced (gains electrons).

    • Reducing Agent: The compound that is oxidized while reducing another compound.

    • Oxidizing Agent: The compound that is reduced while oxidizing another compound.

  • Implications of Oxidation-Reduction:

    • Oxidation leads to a gain of oxygen atoms and loss of hydrogen atoms.

    • Reduction leads to loss of oxygen atoms and gain of hydrogen atoms.

5.5 The Mole and Avogadro’s Number

  • Definition of a Mole: A mole contains approximately 6.022 imes 10^{23} particles (Avogadro's Number).

    • Example: 1 mole of C atoms = 6.022 imes 10^{23} C atoms.

  • Molar Mass:

    • Molar mass is the mass of one mole of a substance expressed in grams/mole (g/mol).

    • For instance, 1 mole of water (H2O) has a molar mass of 18.015 g/mol, while 1 mole of carbon (C) weighs 12.01 g/mol.

  • Using Molar Mass in Calculations:

    • Molar mass can convert between grams and moles of a substance.

    • Example problem: How many molecules are in 5.0 moles of carbon dioxide? (Solution requires matching mole ratios and applications of Avogadro's number.)

5.7 Mole Calculations in Chemical Equations

  • A balanced equation indicates the number of moles for each reactant and product.

    • Coefficients serve as conversion factors for mole ratios to convert moles of one reactant to moles of another.

5.9 Percent Yield

  • Theoretical Yield: The expected amount of product based on reactants in a balanced chemical equation.

  • Actual Yield: The measured amount of product obtained from a reaction.

  • Percent Yield Calculation:

    • Formula: ext{Percent Yield} = rac{ ext{Actual Yield}}{ ext{Theoretical Yield}} imes 100.

5.10 Limiting Reactants

  • Limiting Reactant: The reactant that is completely consumed in a reaction, limiting the amount of product formed.

    • To determine, compare available quantities of reactants and calculate required amounts using mole ratios from the balanced equation.

  • Excess Reactant: The reactant present in a quantity greater than necessary for the reaction.

Conclusion

  • Understanding chemical reactions involves recognizing types, balancing equations, and applying the concepts of moles, yield, and limiting reactants in practical scenarios.