CHEM 1311 Chapter 4 part 3

Main Ideas

1. Oxidation-Number Rules

• Any element in its elemental form has an oxidation number of 0. This includes elements that exist as single atoms, such as sodium or magnesium, and elemental diatomic molecules, such as H₂, O₂, and N₂.
• In a compound, assign an oxidation number first to the more electronegative element. Its oxidation number is often the charge it would have as an ion.
• The oxidation numbers in a neutral compound must add up to 0. In a polyatomic ion, they must add up to the ion’s overall charge.
• Subscripts matter because they indicate how many atoms contribute to the total. Coefficients in front of a formula do not change the oxidation number of an atom in that formula.
• Examples:
  • In NaCl, chlorine is assigned −1, so sodium is +1.
  • In magnesium nitride, nitrogen is assigned −3; the nitrogen atoms’ combined oxidation number must be balanced by magnesium.

2. Limiting Reactants

The instructor introduces limiting reactants through a recipe analogy: one cup of water plus two cups of flour makes six tortillas. With three cups of water and four cups of flour, the flour runs out first, so it limits production. The recipe can make 12 tortillas, with one cup of water left over.

A limiting reactant is used up first and determines the maximum amount of product. The other reactant is the excess reactant.

Procedure for Limiting-Reactant Problems

1. Balance the chemical equation first.
2. Identify the amounts of the reactants. If amounts are given in grams, convert them to moles before applying the balanced-equation mole ratios.
3. Determine the limiting reactant using either method:
   • Compare possible product amounts: Convert each reactant amount to the same product using the mole ratios. The reactant that produces less product is limiting.
   • Compare reactants directly: Use the mole ratio to calculate how much of one reactant is needed to consume the other. Compare the required amount with what is available. The reactant in short supply is limiting.
4. If there is more than one product, compare amounts of the same product for both reactants.
5. Once the limiting reactant is identified, use it as the starting point for later stoichiometry calculations.

Calculating Product and Excess Remaining

• To find how much product forms, convert from the limiting reactant to the desired product using the balanced equation’s mole ratio. Convert to grams if the question asks for grams.
• To find how much excess reactant remains:
  1. Convert the limiting reactant consumed into the amount of excess reactant used.
  2. Subtract that amount from the excess reactant’s initial amount.

Examples

• For the reaction of cobalt(II) oxide with oxygen to form cobalt(III) oxide, the equation is balanced before calculations. Given 0.327 mol CoO and 0.0806 mol O₂, the instructor identifies O₂ as the limiting reactant. The calculation gives 0.1612 mol of product and 0.0046 mol CoO left over.
• For magnesium reacting with oxygen to form magnesium oxide, the equation is balanced first. With 2.2 g Mg and 1.9 g O₂, the instructor identifies Mg as limiting, estimates 0.45 g O₂ left over, and calculates a theoretical yield of 3.6 g MgO.
• In a lead nitrate and sodium chloride example, the reaction is balanced, and the instructor concludes that lead nitrate is limiting and sodium chloride is in excess.

3. Theoretical Yield and Percent Yield

• Theoretical yield is the maximum amount of product calculated from the limiting reactant, assuming the reaction proceeds perfectly.
• Actual yield is the amount of product actually obtained in an experiment or process. It may be lower because of transfer losses, imperfect procedures, or a reaction that does not proceed efficiently or completely.
• The formula is:

Percent yield = (actual yield ÷ theoretical yield) × 100

• In the cobalt oxide example, an actual yield of 0.0976 mol compared with a theoretical yield of 0.1612 mol gives a percent yield of 60.5%.
• In a multistep synthesis, the yields at each stage compound. A low yield at one or more stages can leave very little final product. The instructor notes that this matters in chemical synthesis, including the production of drugs.

Speakers and Sources Featured

• Instructor: The primary speaker, teaching oxidation numbers, limiting reactants, stoichiometry, and percent yield. The instructor is unnamed in the subtitles.
• Students: Several students ask questions and make clarifying comments. They are not identified by name in the subtitles; one is addressed with a name captioned as “Del Crete,” though the auto-generated transcription may be inaccurate.
• External sources: None are cited or featured.