Chemistry 6.0

I. Balanced Chemical Equations A. Provide qualitative and quantitative information B. Supports the Law of Conservation of MatterConservation of Matter

2H2 + O2 2H2O

The above equation is interpreted in terms of particles as follows:

1. 2 molecules of H2 react with 1 molecule of O2

to produce 2 molecules of water.

The ratio of H2 to O2 is 2:1.

or 2 moles of H2 react with 1 mole of O2 to produce 2 moles of water.

The ratio of H2 to O2 is 2:1.

2. It is more convenient to interpret the coefficients as number of moles, because we measure amounts of substances by massing.

C. Stoichiometry 1. The study of the quantitative relationships that exist in a formula or a chemical reaction. 2. Importance

a. Provides for the safe, economical and reproducible

manufacture of chemicals.b. Provides for the safe administration of

pharmaceuticals.

D. Proof of the conservation of matter in a balanced equation 1. Convert all reactants and products to their mass equivalents. 2. Sum up the mass of reactants and compare the sum of the mass of products.

II. Stoichiometry Problems

A. Steps to Solve Problems

1. Write a balanced equation.

2. Identify the given ( ) and the unknown or required substance (?).3. Convert mass of given into moles.

4. Use the mole (molar) ratio to convert from given to required substance.5. If needed, convert moles of required into mass of required substance.

B. Examples1. How many moles of oxygen are required

to react with 16 moles of hydrogen in the production of water?

H2 + O2 H2O 2 2 ?

1 mole O2 2 moles H2

= 8.0 moles O2

Mole ratio links 1 substance to another in a reaction. Required in problem solving.

16 moles H2

2. Antimony reacts with water to produce antimony(III) oxide

and hydrogen. How many moles of hydrogen are produced

from 7.5 moles of antimony?

Sb + H2O Sb2O3 + H22 3 3

11 mol H2

3. What mass of aluminum oxide can be prepared by the reaction of 67.5 g of aluminum in a synthesis reaction?

Al + O2 Al2O34 3 2

128 mol Al2O3

4. Sodium bicarbonate, a.k.a. baking soda, can be used to extinguish a fire. When heated, it decomposes to give carbon dioxide gas which smothers the fire. It also produces sodium carbonate and water. If a sample contains 4.0 g of sodium bicarbonate, what mass of carbon dioxide is produced?

2 NaHCO3 Na2CO3 + H2O + CO2

III. Percent Yield A. Expected Yield: the amount of product

that should be produced (theoretical)

B. Actual Yield: the amount of product that is

actually produced (experimental)

C. Percent Yield: percent of expected yield

that was obtained

% Yield = (actual yield/expected yield) x 100

D. Steps to Solving Percent Yield Problems1. Write a balanced equation2. Identify the given () which is the

mass of reactant, and identify the actual yield.

3. Solve for the expected mass of product using the given mass of reactant.

4. Calculate the % yield.% Yield = actual yield x 100

expected yield

E. Examples1. A reaction between 2.80 g aluminum nitrate and

excess sodium hydroxide produced 0.966 g of

aluminum hydroxide in this double replacement

reaction. Calculate the % yield.

Al(NO3)3 + 3NaOH Al(OH)3 + 3NaNO3

1.03 g Al(OH)3

2. Determine the % yield for the reaction between 3.74 g of

sodium and excess oxygen if 4.24 g of sodium oxide is recovered in the direct combination reaction.

4Na + O2 2Na2O

IV. Limiting Reactants

A. Definition: the reactant that determines, or limits, the amount of product(s) formed in a chemical reaction

B. Problem Solving Tips1. The limiting reactant is not necessarily the

reactant present in the smallest amount 2. When you are given the amounts of 2 or

more reactants, you should suspect that you are dealing with a limiting reactant problem.

C. Steps1. Write a balanced equation2. Calculate the number of moles of each

reactant3. Compare the mole ratios of the reactants as

available ratio (from the given masses) and the required ratio (from the coefficients)

4. Identify the limiting reactant, and use it to calculate the mass of product formed.

D. Examples1. What mass of CO2 could be formed by the combustion of 16.0 g CH4 with 48.0 g O2?

CH4 + 2O2 CO2 + 2H2O

33.0 g CO2

Oxygen is the limiting reactant

2. What is the maximum mass of nickel(II) hydroxide that could be prepared by mixing 25.9 g nickel(II) chloride with 10.0 g sodium hydroxide?

NiCl2 + 2NaOH Ni(OH)2 + 2NaCl

11.6 g Ni(OH)2

Sodium Hydroxide is the LR

V. Sequential Reactions

A. Definition: A chemical process in which several reactions are required to convert starting materials into product(s).

B. Main Concept: The amount of desired product from each reaction is taken as the starting material for the next reaction. Consideration is given to steps that produce less than 100%.

C. Examples1. Given the following sequence of equations, calculate the mass of Ni(CO)4 produced from 75.0 g of carbon. Assume 100% yields.

C + H2O CO + H2

Ni + 4CO Ni(CO)4

2. Hydrogen, obtained by the electrical decomposition of water, is combined with chlorine to produce 84.2 g of hydrogen chloride. Calculate the mass of water decomposed. Assume 100% yield.

2H2O 2H2 + O2

H2 + Cl2 2 HCl

VI. Solution StoichiometryA. Many reactants are introduced to a reaction chamber as a solution.B. The most common solution concentration is molarity.

molarity = mol/liter

C. Examples1. Excess lead(II) carbonate reacts with 27.5 mL of

3.00M nitric acid. Calculate the mass of lead(II) nitrate formed

PbCO3 + 2HNO3 Pb(NO3)2 + H2CO3

2. Calculate the volume, in mL, of a 0.324 molar solution of sulfuric acid required to react completely with 2.792 g of sodium carbonate according to the equation below.

H2SO4 + Na2CO3 Na2SO4 + CO2 + H2O

ProblemsC3H8 + 5O2 → 3CO2 + 4H2O ΔH = -2.22x103

kJ

1. How much heat is released when 22.0g of propane is burned?

2. How much carbon dioxide is produced, in grams, when 2,500 kJ of energy is released?

-1.11x103 kJ (released)

150 g CO2

ΔΔH from H from ΔΔHHff

The standard enthalpy change of a reaction is equal to the sum of the standard molar enthalpies of formation of the products multiplied by its coefficient, n, in the balanced equation, minus the corresponding sum of standard molar enthalpies of formation of reactants.

Hrxn = ∑ nHf,products - ∑ nHf,reactants

***By definition, the standard enthalpy of formation of the most stable form of any element is zero

because there is no formation reaction needed when the element is already in its standard

state.***

ΔH from ΔHf Problem:Using the Heats of Formation Table, calculate

the H for the following reaction:SF6(g) + 3H2O(l) 6HF(g) + SO3(g)

Write the thermochemical equation for this reaction:

SF6 + 3H2O + 45.1kJ 6HF + SO3

45.1 kJ

ΔH from ΔHf Problem:Using the Heats of Formation Table,

calculate the standard heat of combustion for propane.C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g)

Write the thermochemical equation for this reaction:

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g) + 2043.9kJ

-2043.9 kJ