Quantitative Information from Chemical Equations Coefficients in a balanced equation number of molecules (formula units, etc) number of moles 2 H 2 + O 2  2 H 2 O 2 molecules 1 molecules 2 moles1 mole2 moles

Quantitative Information from Chemical Equations In industry, aspirin is prepared by: C 7 H 6 O 3 + C 4 H 6 O 3  C 9 H 8 O 4 + HC 2 H 3 O 2 Salicylic acid Acetic anhydride Acetyl- salicylic acid Acetic acid Chemists want to use the right amount of reactants in order to minimize “left over” reactants that contaminate the product.

Quantitative Information from Chemical Equations The coefficients in a balanced chemical equation can be used to relate the number of moles of each substance involved in a reaction. Molar ratios mol reactant mol reactant mol product mol reactant mol product mol product

3 different molar ratios (and their inverses) can be written: 1 mol N 2 1 mol N 2 3 mol H 2 3 mol H 2 2 mol NH 3 2 mole NH 3 Molar Ratios from Chemical Equations For the reaction: N H 2  2 NH 3

Molar Ratios from Chemical Equations MOLAR RATIOS = CONVERSION FACTORS Moles of product that can be formed from a certain number of moles of reactant(s) Moles of reactants needed to form a certain number of moles of product Moles of reactant 2 needed to completely react with reactant 1

Quantitative Information from Chemical Equations Example: Nitrogen and hydrogen react to form ammonia. If you have 1.0 mole of hydrogen, how many moles of ammonia can you produce? Write a balanced equation first, if I don’t give you one!

Quantitative Information from Chemical Equations Example: Nitrogen and hydrogen react to form ammonia. If you have 1.0 mole of H 2, how many moles of N 2 will be required to completely react all of the H 2 ? Balanced equation:

Quantitative Information from Chemical Equations Example: Nitrogen and hydrogen react to form ammonia. How many moles of N 2 are needed to produce 0.50 moles of NH 3 ? Balanced equation:

Quantitative Information from Chemical Equations Finding # of moles is great BUT You don’t measure out moles in the lab! Chemists use a balance to measure the mass of a substance used or produced in a reaction. How can you determine the mass of reactants or products?

Quantitative Information from Chemical Equations Use the molar mass to convert from moles to grams The number of grams of a substance per mole

Quantitative Information from Chemical Equations Mass (g) Compound A Moles Compound A Moles Compound B Mass (g) Compound B Molar mass Molar mass Molar ratio “The MAP”

Quantitative Information from Chemical Equations Example: How many grams of water will be produced by the complete combustion of 10.0 g of propane, C 3 H 8 ?

Combustion Reactions – Revisited! You should be able to write a balanced equation for the combustion of an organic compound or a metal. Organic compounds: Metals: Not bal.

Quantitative Information from Chemical Equations Example: Hydrofluoric acid can’t be stored in glass because it attacks the silicates in the glass: Na 2 SiO 3 (s) + 8 HF (aq)  H 2 SiF 6 (aq) + 2 NaF (aq) + 3 H 2 O How many grams of HF are needed to dissolve 55.0 g of Na 2 SiO 3 ?

Quantitative Information from Chemical Equations

Making Bologna Sandwiches Suppose you were going to make bologna sandwiches: ++ 2 Bread + 1 Bologna  1 sandwich

Making Bologna Sandwiches 2 Bread + 1 Bologna  1 sandwich + We can only make 4 sandwiches because we don’t have enough bologna! Bologna = limiting reagent or limiting reactant

Limiting Reagent or Limiting Reactants Similar situations occur in chemical reactions when one of the reactants is used up before the others. No further reaction can occur The excess reactant(s) are “leftovers.”

Limiting Reagent or Limiting Reactants 2 H 2 (g) + O 2 (g)  2 H 2 O (l) If we react 10 moles of H 2 with 7 moles of O 2, not all of the O 2 will react because we will run out of H 2 first!

Limiting Reagent or Limiting Reactants 10 H 2 7 O 2 10 H 2 O + 2 O 2

Limiting Reagent or Limiting Reactants Limiting reagent (limiting reactant): the reactant that is completely consumed in a reaction determines or limits the amount of product formed.

Limiting Reagent or Limiting Reactants Two approaches to identifying the limiting reactant: Compare the number of moles of each reactant needed with the number of moles of each reactant available Calculate the number of grams of product that each reactant could form Reactant that forms the least amount of product will be the limiting reagent. OR

Limiting Reagent or Limiting Reactants First Method: 1)Convert the mass of each reactant to moles. 2)Pick one of the reactants (it doesn’t matter which one) and calculate the number of moles of the other reagent needed to completely react with the one chosen. 3)Compare the # moles needed vs. # moles available

Limiting Reagent or Limiting Reactants Second Method: 1)Calculate the number of grams of product that each reactant could form. 2)Limiting reagent = reactant that forms the smallest mass of product.

Limiting Reagent or Limiting Reactants Example: If 10.0 grams of H 2 are mixed with 75.0 grams of O 2, which reactant is the limiting reagent? 2 H 2 (g) + O 2 (g)  2 H 2 O (l)

Limiting Reagent or Limiting Reactants Method 1

Limiting Reagent or Limiting Reactants Method 2

Limiting Reagent or Limiting Reactants Once you find the limiting reagent, you can also find the amount of product that can be formed in the reaction. If you used the second method for identifying the limiting reagent, you’ve already done this! Moles limiting reagent Moles product grams product Molar ratio Molar mass

Limiting Reagent or Limiting Reactants Example: How many grams of water are produced by burning 5.00 g of methane in the presence of 5.00 g of oxygen? CH 4 (g) + 2 O 2 (g)  CO 2 (g) + 2 H 2 O (l)

Example Since the mass of both reactants is given, you need to check to see if one of them is a limiting reagent.

Example

Yield Theoretical Yield: the quantity of product that is calculated to form when all of the limiting reagent reacts Actual yield: the amount of product actually obtained in a reaction

Yield Often, the actual yield is less than the theoretical yield: reactants may not react completely t i.e. the reaction does not go to completion “by-products” may form t unwanted side reactions (competing reactions) difficulty isolating and purifying the desired product

Yield Percent Yield: relates the actual yield and the theoretical yield % Yield = actual yield x 100% theoretical yield

Yield Example: In a certain reaction between H 2 and CO to form methanol, the theoretical yield is 83.3 g of CH 3 OH. If the actual yield of the reaction was 81.5 g, what was the % yield?