Chemical Stoichiometry

Stoichiometry - The study of quantities of materials consumed and produced in chemical reactions.

Atomic MassesElements occur in nature as mixtures of

isotopes

Carbon = 98.89% 12C1.11% 13C<0.01% 14C

Average atomic mass = (% of each isotope)(atomic mass of each isotope)

100

Carbon atomic mass = 12.01 amu

The MoleThe number equal to the number of carbon atoms in exactly 12 grams of pure 12C.

1 mole of anything = 6.022 1023 units of that thing

Molar MassA substance’s molar mass (molecular weight) is the mass in grams of one mole of the compound.

CO2 = 44.01 grams per mole

Percent Composition

Mass percent of an element:

For iron in iron (III) oxide, (Fe2O3)

massmass of element in compound

mass of compound% 100%

Formulas

molecular formula = (empirical formula)n

[n = integer]

molecular formula = C6H6 = (CH)6

empirical formula = CH

Empirical Formula Determination

1. Base calculation on 100 grams of compound.

2. Determine moles of each element in 100 grams of compound.

3. Divide each value of moles by the smallest of the values.

4. Multiply each number by an integer to obtain all whole numbers.

Percent Composition/ Empirical Formula Problem

1. An ion containing only oxygen and chlorine is 31% oxygen by mass. What is its empirical formula?

2. What is the percent composition by mass of the elements in the compound NaNO3?

Chemical Equations

Chemical change involves a reorganization of the atoms in

one or more substances.

A representation of a chemical reaction:

C2H5OH + 3O2 2CO2 + 3H2O

reactants products

Chemical Equation

C2H5OH + 3O2 2CO2 + 3H2O

The equation is balanced.

1 mole of ethanol reacts with 3 moles of oxygen

to produce2 moles of carbon dioxide and 3

moles of water

Hydrogen and Nitrogen React to Form Ammonia According to the

Equation N2 + 3H22NH3

Schematic Diagram of the Combustion Device Used to

Analyze Substances for Carbon and Hydrogen

Combustion Stoichiometry Problem

A gaseous hydrocarbon sample is completely burned in air, producing 1.80 liters of carbon dioxide at standard temperature and pressure and 2.16 grams of water.

a. What is the empirical formula for the hydrocarbon?b. What was the mass of the hydrocarbon consumed?c. The hydrocarbon was initially contained in a closed

1.00 liter vessel at a temperature of 32oC and a pressure of 760 mmHg. What is the molecular formula of the hydrocarbon?

d. Write the balanced equation for the combustion of the hydrocarbon.

Calculating Masses of Reactants and

Products1. Balance the equation.2. Convert mass to moles.3. Set up mole ratios.4. Use mole ratios to calculate

moles of desired substituent.5. Convert moles to grams, if

necessary.

Solving a Stoichiometry Problem with a Limiting Reactant

1. Balance the equation.

2. Convert masses to moles.

3. Determine which reactant is limiting.

4. Use moles of limiting reactant and mole ratios to find moles of desired product.

5. Convert from moles to grams.

The limiting reactant is the reactant that is consumed first, limiting the amounts of products formed.

Limiting Reactant Problem

2 Mg(s) + 2CuSO4 (aq) + H2O (l) 2 MgSO4 (aq) + Cu2O(s) + H2 (g)

a. If 1.46 grams of Mg(s) are added to 500. Milliliters of a 0.200 molar solution of CuSO4, what is the maximum molar yield of H2 (g)?

b. When all of the limiting reactant has been consumed in (a), how many moles of the other reactant (not water) remain?

c. What is the mass of the Cu2O produced in (a)?

d. What is the value of [Mg2+] in the solution at the end of the experiment? (Assume that the volume of the solution remains unchanged.)

Percent Yield- the actual yield of a product as a percentage of theoretical yield

A chemist runs the reaction described below

2 Mg(s) + 2CuSO4 (aq) + H2O (l) 2 MgSO4 (aq) + Cu2O(s) + H2 (g)

The expected yield of Cu2O from a previous problem was 4.29 grams. The chemist is only able to collect 3.96 grams. What is the chemist’s percent yield?

Molarity Molarity (M) = moles of solute per volume of solution in liters:

What’s one intrepretation of the label- “3 M HCl?”

Common Terms of Solution Concentration

Stock - routinely used solutions prepared in concentrated form.

Concentrated - relatively large ratio of solute to solvent. (5.0 M NaCl)

Dilute - relatively small ratio of solute to solvent. (0.01 M NaCl)

Solution ChemistryThe Water Molecule is Polar

Polar Water Molecules Interact with the Positive

and Negative Ions of a Salt

BaCI2 Dissolving

Preparation of a Standard Solution

A chemist whishes to prepare 1.00L of a 0.200 M sodium hydroxide solution. Describe the steps, with calculations, necessary to complete this task starting with solid sodium hydroxide and distilled water.

Dilution of Solutions- M1V1 = M2V2

You’ve been asked to prepare 150 ml of a 0.035M solution of sodium hydroxide from the 0.200M stock sodium hydroxide solution prepared earlier. Detail the steps necessary to complete this task.

(a) A Measuring Pipet

(b) A Volumetric (transfer) Pipet

Beer- Lambert Law Beer’s Law

Relates the amount of light being absorbed to the concentration of the substance absorbing the light

A=abcA = measured absorbancea =

b =

c =

Beer’s Law Sample Problems

1. A solution with a concentration of 0.14M is measured to have an absorbance of 0.43. Another solution of the same chemical is measured under the same conditions and has an absorbance of 0.37. What is its concentration?

2. The following data were obtained for 1.00 cm samples of a particular chemical. What is the concentration of a 1.00 cm sample that has an absorbance of 0.60?

Conc. (M)

Abs.

0.50 0.69

0.40 0.55

0.30 0.41

0.20 0.27

3. The absorptivity of a particular chemical is 1.5/M·cm. What is the concentration of a solution made from this chemical if a 2.0 cm sample has an absorbance of 1.20?

Beer’s Law Sample Problems

4. Using the data from the graphing example in question #2, what are the concentrations of solutions with absorbances of 0.20, 0.33, and 0.47? 

5. A solution is prepared to be 0.200M. A sample of this solution 1.00 cm thick has an absorbance of 0.125 measured at 470nm and an absorbance of 0.070 measured at 550nm. Calculate the concentrations of the following solutions:

Sample Absorbance

Wavelength

Path length1 0.055 470nm 1.00cm

2 0.155 470nm 1.00cm3 0.120 550nm 1.00cm4 0.048 550nm 5.00cm

Stoichiometry Problem Set

1. Aluminum oxide is to be made by combining 5.00 g of aluminum with oxygen gas. How much oxygen is needed in moles? In grams? In liters?

2. During its combustion, ethane (C2H6) combines with oxygen gas to give carbon dioxide and water. A sample of ethane was burned completely and the water that formed had a mass of 1.61 g. How much ethane, in moles and in grams, was in the sample?

3. Chloroform, CHCl3, reacts with chlorine gas to form carbon tetrachloride and hydrogen chloride. In one experiment the reactants were initially presented in a ratio of 1 to 1 by mass; specifically, 25.0 g of CHCl3 was mixed with 25.0 g of Cl2 (g). Which is the limiting reactant? What is the maximum yield of carbon tetrachloride in moles and in grams?

Stoichiometry Problem Set

4. One of the steps in one industrial synthesis of sulfuric acid (H2SO4) from sulfur is the conversion of sulfur dioxide (SO2) into sulfur trioxide (SO3) by this reaction:

2SO2+O2 2SO3

In one “run,” 1.75 kg of SO2 was used and 1.72 kg of SO3 was isolated from the mixture of products. What was the percent yield?

5. A student needs 0.250 mol of NaCl and all that is available is a solution labeled “0.400 M NaCl.” What volume of the solution should be used? Give your answer in milliliters.

6. Describe how to prepare 250 mL of 0.200 M NaHCO3.

Stoichiometry Problem Set

7. How many milliliters of 0.114 M H2SO4 solution are necessary to completely neutralize 32.2 mL of 0.122 M NaOH?

8. Describe how to make 500 mL of 0.20 M NaOH from 0.50M NaOH.