information given by the chemical equation

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Information Given by theChemical Equation

• Balanced equations show the relationship between the relative numbers of reacting molecules and product molecules.

2 CO + O2 2 CO2

2 CO molecules react with 1 O2 molecule to produce

2 CO2 molecules.


• The information given is relative:

2 CO + O2 2 CO2

Therefore:

200 CO molecules react with 100 O2 molecules to produce 200 CO2 molecules.

2 moles CO molecules react with 1 mole O2 molecules to produce 2 moles CO2 molecules.

12 moles CO molecules react with 6 moles O2 molecules to produce 12 moles CO2 molecules.



• The coefficients in the balanced chemical equation also show the mole ratios of the reactants and products.

• Since moles can be converted to masses, we can also determine the mass ratios of the reactants and products.



2 CO + O2 2 CO2

2 moles CO; 1 mole O2; 2 moles CO2

Since 1 mole of CO = 28.01 g (12.01g + 16g), 1 mole O2 = 32.00 g (2 x 16.00 g), and 1 mole CO2 = 44.01 g (12.01 g + 32 g), then we have the following quantities:

2(28.01) g CO; 1(32.00) g O2; 2(44.01) g CO2



Example #1

Determine the number of moles of carbon monoxide required to react with 3.2 moles oxygen, and determine the moles of carbon dioxide produced.


Example #1 (cont.)

• Write the balanced equation: 2 CO + O2 2 CO2

• Use the coefficients to find the mole relationship:

2 moles CO = 1 mol O2 = 2 moles CO2


• Use dimensional analysis:

Example #1 (cont.)


Example #2

Consider the following reaction:

CH4(g) + 4Cl2(g) CCl4(l) + 4HCl(g)

How many moles of Cl2 would be needed to react with 0.3 mol of CH4?


Example #3

Determine the number of grams of carbon monoxide required to react with 48 g of oxygen, and determine the number of grams of carbon dioxide produced.


Example #3 (cont.)

• Write the balanced equation:

2 CO + O2 2 CO2

• Use the coefficients to find the mole relationships:

• Determine the molar mass of each:


• Use the molar mass of the given quantity to convert it to moles.

• Use the mole relationships to convert the moles of the given quantity to the moles of the desired quantity:

Example #3 (cont.)


• Use the molar mass of the desired quantity to convert the moles to mass:

Example #3 (cont.)


Example #4

Consider the reaction below:

Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

If you start with 0.056 mol of Zn, how much HCl in grams is needed?

How much ZnCl2 in grams can be produced?


Limiting and Excess Reactants

• Limiting reactant: a reactant that is completely consumed when a reaction is run to completion

• Excess reactant: a reactant that is not completely consumed in a reaction

• Theoretical yield: the maximum amount of a product that can be made by the time the limiting reactant is completely consumed


+

+

+

Consider the following reactions:

What is the limiting reactant and what is the excess reactant?


Example #5

Determine the number of moles of potassium oxide produced when 8 moles of potassium metal reacts with 3 moles of oxygen.


Example #5 (cont.)

• Write the balanced equation:

• Use the coefficients to find the mole relationships:


• Use dimensional analysis to determine the # of moles of reactant A needed to react with reactant B:

Example #5 (cont.)


• Compare the calculated # of moles of reactant A to the # of moles of reactant A given.

• If the calculated # of moles is greater, then A is the limiting reactant; if the calculated # of moles is less, then A is the excess reactant.

• What is the limiting reagent in this reaction?

Example #5 (cont.)


• Use the limiting reactant to determine the maximum # of moles of product that can form:

Example #5 (cont.)


Example #6

Consider the reaction between zinc metal and iodine (I2) to form zinc(II) iodide. If you begin with 50g of Zn and 50 g of I2, what is the limiting reagent? How much zinc(II) iodide in grams can be formed in this reaction?


Example #7

Consider the following unbalanced equation:

SF4 + I2O5 IF5 + SO2

Calculate the maximum number of grams of IF5 that can be produced from 10 g of SF4 and 10 g of I2O5.


Percent Yield

• Most reactions do not go to completion.

• Theoretical yield: The maximum amount of product that can be made in a reaction

• Actual yield: the amount of product actually made in a reaction

Percent Yield = Actual Yield

Theoretical Yieldx 100%


Example #8

If in the last problem 9.0 g of IF5 is actually made, what is the percent yield of IF5?


Example #9


N2(g) + 3H2(g) 2NH3(g)

When 10g of N2 and 10 g of NH3 react, 8 g of NH3 is formed.

What is the percent yield of NH3?


Example #10


CS2(l) + 3O2(g) CO2(g) + 2SO2(g)

When 1 g of CS2 reacts with 1 g of O2, 0.33 g of CO2 results. What is the percent yield of CO2?

information given by the chemical equation

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