Unit 4 Key Concepts: Chemical Reactions

Topic 1 – Introduction for Reactions

• Type of Change
  • Physical:
    • Description: Change in appearance or state.
    • Examples: Phase changes, separating a mixture.
    • Equation Examples:
      • H2O(s) \rightarrow H2O(l)
      • H2O(l) \rightarrow H2O(s)
  • Chemical:
    • Description: Transformation of substances; substances are changed into new substances.
    • Example: Mg(s) + H2O(l) \rightarrow Mg(OH)2(aq) + H_2(g)
      • (aq) aqueous means “dissolved in water”
    • Symbols:
      • \rightleftharpoons Reversible Reaction
      • \triangle Add Heat
      • \lightning Add Electricity
      • \rightarrow^{Pt} Add a catalyst (Platinum in this case)
• Signs that a Chemical Reaction is Taking Place
  • Production of light
  • Large change in temperature
  • Production of a precipitate
  • Production of a gas
  • Change in color
  • Change in odor or texture

Topic 2 – Net Ionic Equations

• Na(s) + Cl_2(g) \rightarrow NaCl(s)
• Solubility Rules:
  • Compounds containing alkali metal ions (Li^+, Na^+, K^+, etc.) or the ammonium ion (NH_4^+) will always be soluble in water.
  • Compounds containing the nitrate ion (NO_3^-) will always be soluble in water.
• Example Reactions:
  • Solutions of barium nitrate and potassium sulfate are mixed together:
    • Ba(NO3)2 + K2SO4 \rightarrow BaSO4 + 2 KNO3
  • Solutions of silver nitrate and zinc chloride are mixed together:
    • 2 AgNO3 + ZnCl2 \rightarrow 2 AgCl + Zn(NO3)2
  • Complete molecular equation
  • Complete ionic equation
  • Zn^{2+} and NO_3^- are spectator ions.
  • Net ionic equation

Topic 3 – Representations of Reactions

• 2 CO(g) + O2(g) \rightarrow 2 CO2(g)
  • O_2 is the limiting reactant
  • CO is the excess reactant
  • Reactants transform into Products

Topic 4 – Physical and Chemical Changes

• The dissolving of an ionic compound is a physical change, but it has chemical aspects, as well.

Topic 5 – Stoichiometry

• Al(s) + Cu(NO3)2(aq) \rightarrow Al(NO3)3(aq) + Cu(s)
• Example Problem:
  • If we place a 5.00 gram piece of aluminum into excess copper(II) nitrate solution, how many grams of copper metal can be produced?
    1. Convert to moles.
    2. Mole ratio.
    3. Convert to final unit.
• Percent Yield Calculation:
  • Percent Yield = \frac{Actual Yield}{Theoretical Yield} \times 100
• Limiting Reactant Example:
  • If we add 1.00 g Ca to 0.100 mol H2O, how many moles of H2 gas should be produced?
    • Ca(s) + 2 H2O(l) \rightarrow Ca(OH)2(aq) + H_2(g)

Topic 6 – Introduction to Titration

• Titrant in the Buret (usually a base)
• Analyte in the Flask (usually an acid)
• Equivalence Point:
  • moles acid = moles base
• Endpoint:
  • moment of color change
• Titration Calculation Example:
  • In a titration, 10.00 mL of acetic acid are placed into an Erlenmeyer flask and titrated with 0.1245 M NaOH solution. If it takes 21.45 mL of the NaOH solution to reach the equivalence point, what is the concentration of the acetic acid?
    • MA VA = MB VB

Topic 7 – Types of Chemical Reactions

• Precipitation Reactions
  • Two aqueous solutions are mixed together
  • The product is an insoluble solid precipitate
  • Spectator ions do not react, and they remain in solution
    • Example: Sodium carbonate and nickel(II) chloride solutions are mixed.
• Redox Reactions
  • Involve the transfer of electrons
  • Gaining electrons is reduction; losing electrons is oxidation
    • LEO the lion goes GER (losing electrons oxidation, gaining electrons reduction)
    • 2 Ca(s) + O_2(g) \rightarrow 2 CaO(s)
      • Ca goes oxidation from 0 to +2
      • O goes reduction from 0 to -2
      • Ca is reducing agent
      • O is oxidizing agent
• Rules for Determining Oxidation State
  • All free-standing elements have an oxidation state of 0.
  • When in compounds, alkali metals are +1 and alkaline earth metals are +2.
  • Hydrogen is +1 when bonded to a nonmetal. Hydrogen is –1 when bonded to a metal.
  • Oxygen in a compound is always –2. Only exception: Oxygen in a peroxide is –1
  • If the elements are not obvious, the oxidation state of the most electronegative atom can be predicted from the periodic table.
  • All compounds have a total oxidation state of 0, and a polyatomic ion has a total oxidation state equal to its charge.
  • Use the oxidation states you know to solve for the ones you don’t.
    • Example: Cr2O7^{2-}

Topic 8 – Introduction to Acid-Base Reactions

• HBr(aq) + H2O(l) \rightarrow H3O^+(aq) + Br^–(aq)
  • acid + base --> (hydronium) acid + base
  • Brønsted-Lowry acid = proton (H^+) donor; conjugate acid-base pair
  • Brønsted-Lowry base = proton (H^+) acceptor
  • HBr = very strong acid / Br – = very weak base
  • HBr = very strong acid / H_2O = very strong base

Topic 9 – Oxidation-Reduction (Redox) Reactions

• A strip of aluminum is placed into a solution of lead(II) nitrate.