CHM2045 Class 4- Balancing Equations and Quantitative Relations

Page 1: Chemical Reactions

  • Reactants: Zinc (Zn) + Iodine (I2)

  • Product: Zinc Iodide (ZnI2)

Page 2: Introduction to Chemical Reactions

  • Chemical reactions occur when bonds between atoms are formed or broken.

  • These reactions involve:

    • Changes in matter

    • Formation of new materials with new properties

    • Accompanying energy changes.

  • Notation used in chemistry:

    • Symbols: Represent elements.

    • Formulas: Describe compounds.

    • Chemical Equations: Describe a chemical reaction.

Page 3: Parts of a Reaction Equation

  • Chemical Equations: Show conversion of reactants to products.

    • Reactants are on the left side of the arrow.

    • Products are on the right side of the arrow.

    • A plus sign (+) separates different molecules on the same side.

    • The arrow (→) is read as "yields".

    • Example: C + O2 → CO2

    • This is read as “carbon plus oxygen react to yield carbon dioxide.”

Page 4: Example Reaction

  • Charcoal (carbon) reacts with oxygen to yield carbon dioxide:

    • Equation: C + O2 → CO2

    • Provides both qualitative and quantitative meaning.

Page 5: Symbols Used in Chemical Equations

  • Common Symbols:

    • Solid (s)

    • Liquid (l)

    • Gas (g)

    • Aqueous solution (aq)

    • Catalyst (e.g., H2SO4)

    • Escaping gas (g)

    • Precipitate (s)

    • Change of temperature (Δ)

Page 6: Representing Reactions

  • Represents H2 reacting with O2 to form H2O:

    • A chemical equation uses chemical symbols to show changes in a chemical reaction.

Page 7: Conservation of Matter

  • A chemical equation must be balanced, preserving the number of atoms of each element on both sides.

  • historical reference: Antoine Lavoisier, 1788.

Page 8: Balancing Equations

  • Only coefficients can be added to balance equations, never change subscripts as they define compounds' identities.

    • Subscripts are determined by valence electrons.

Page 9: Subscripts vs. Coefficients

  • Subscripts: Indicate the number of atoms of an element in a compound.

  • Coefficients: Indicate the quantity (number of molecules) of the compound.

Page 10: Steps to Balancing Equations

  1. Write correct formulas for reactants and products without balancing.

  2. Count atoms for each element on both sides.

  3. Add coefficients to balance atoms for each element.

  4. Verify if all atoms are balanced and coefficients are in lowest ratios.

Page 11: Helpful Hints for Balancing Equations

  • Tackle one element at a time, left to right but save H and O for last.

  • If Oxygen is unbalanced with no whole number solution, double all coefficients.

  • For polyatomic ions appearing on both sides, balance them as independent units.

Page 12: Example of Balancing Equations

  • Example: H2(g) + O2(g) → H2O(l)

    • Initial equation is unbalanced and requires adjustments.

Page 13: Balancing Example Explained

  • Balanced Equation:

    • H2(g) + O2(g) → 2 H2O(l)

    • Two hydrogen atoms combine with one oxygen atom to produce two water molecules, achieving balance.

Page 14: Practice Balancing Equations

  • Balance the following reactions:

    • Al(s) + Br2(l) → Al2Br6(s)

    • C3H8(g) + O2(g) → CO2(g) + H2O(g)

Page 15: Balancing Continued

  • Balanced Reactions:

    • Al(s) + 3 Br2(l) → Al2Br6(s)

    • C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g)

Page 16: Balancing Complex Reactions

  • Example: B4H10(g) + O2(g) → B2O3(g) + H2O(g)

    • Initial coefficients: 2 B4H10 + 5 O2 → 2 B2O3 + 10 H2O

    • Multiply coefficients for the final balance.

Page 17: Using Polyatomic Ions

  • Approach reactions involving polyatomic ions such as PO4 as single units when balancing.

    • Example: Na3PO4 + Fe2O3 → Na2O + FePO4

Page 18: Interpreting Chemical Equations

  • Example Equation: 4 Al(s) + 3 O2(g) → 2 Al2O3(s)

    • It indicates that 4 aluminum atoms combined with 3 oxygen molecules yield 2 Al2O3 units.

Page 19: Understanding Chemical Equations

  • Breakdown described:

    • 2 Mg + O2 → 2 MgO

    • 2 atoms of Mg and 1 molecule of O2 yield 2 formula units of MgO.

    • Proper mass translation is important during balance.

Page 20: Steps for Calculating Mass Changes

  1. Write a balanced chemical equation.

  2. Convert known quantities into moles.

  3. Calculate sought moles using coefficients.

  4. Convert moles to desired units.

Page 21: Practical Example of Reaction

  • When methanol (CH3OH) burns in air:

    • 2CH3OH + 3O2 → 2CO2 + 4H2O

    • Using 209 g of methanol produces a calculated 235 g of water based on stoichiometry.

Page 22: Limiting Reagents

  • After reaction analysis:

    • Limiting reagent details demonstrated with reactant references.

Page 23: Understanding Limiting Reactants

  • Example reaction: 124 g Al with 601 g Fe2O3 leading to formation of Al2O3.

    • Comparison and calculations reveal Fe2O3 is in excess; Al is the limiting reagent.

Page 24: Yield Concepts

  • Theoretical Yield: The max amount of product if limiting reagent reacts fully.

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

  • Percent Yield Formula:

    • % Yield = (Actual Yield / Theoretical Yield) x 100.