Chapter 3 Stoichiometry Deals with quantities of materials consumed and produced in chemical...

Chapter 3Stoichiometry

Deals with quantities of materials consumed and produced in chemical reactions

Stoichiometry Terms Atomic Mass

-Average value of the isotopes in a naturally occurring element-Sometimes called the atomic weight

Mole- the number equal to the number of carbon atoms in exactly 12 grams of pure Carbon-12- symbol is mol

Terms con. Avogadro’s Number

- one mole of anything contains 6.022 x 1023 units of that substance

- Mass of 1 mol of an element= its atomic mass in grams

- One mole of Al= 6.022 x 1023 atoms of Al

Examples Determine the number of moles of atoms

and the number of atoms I a 10.0 g sample of aluminum.

Calculate the number of moles in a sample of cobalt containing 5.00 x 1020 atoms and calculate the mass of the sample.

Molar Mass Mass in grams of 1 mole of the compound Example: Calcium carbonate(CaCO3) is

found in limestone, marble, and chalk. Determine its molar mass. If a sample contains 4.86 moles, determine its mass and the mass of carbonate ions present?

Another example Calculate the number of molecules in 1.058

moles of H20.

Percent Compostion of Compounds

Percent by mass of each element present in a compound1. Determine the molar mass of the compound2. Divide the mass of each element by the total molar mass

Example: Determine the mass percent of each element in glucose, C6H12O6.

Determining the Formula of a Compound

Two Types of Formulas

1. Emperical formula

- the simplest whole number ration of atoms in a compound

2. Molecular Formula

- exact formula of a compound

- molar mass of the compound is needed

Examples of Emperical and Molecular

Emperical Molecular

H2O H2O

CH2O C6H12O6

Determining the Emperical formula

You need to convert % composition to empirical formula

Percent to mass Mass to moles Divide by small Multiply ‘til whole

Example of Emperical Formula A compound consists of the following:

71.65 % Cl 24.27% C 4.07% H

Determine its emperical formula.

Determining the Molecular Formula

1. Determine the emperical formula

2. Calculate the emperical formula mass

3. Divide the molar mass of the compound by the emperical formula mass

4. Multiply the emperical formula subscripts by this factor

Molecular formula example From a previous example, the empirical

formula was ClCH2. The molar mass of the compound is 98.96 g/mol. What is the molecular formula?

Chemical Equations Represents a chemical reaction showing the #’s of

reactants and products Types of Info found in an Equation

1. Nature of the reactants/products Solid(s) Liquid(l) Gas(g) Dissolved in water(in aqueous solution)(aq)

2. Relative # of each reactant and product

Balancing Chemical Equations CH4 + O2 CO2 + H20

Coefficient- used to balance equations 1. Tells how many moles of

reactants/products 2. Tells how many molecules involved in the

reaction

Balancing examples Solid ammonium dichromate when ignited

produces solid chromium(III) oxide, nitrogen gas, and water vapor. Write a balance equation.

Ammonia gas, NH3, reacts with oxygen gas to form gaseous nitric oxide, NO, and water vapor. Write a balanced equation.

Stoichiometric Calculations What mass of oxygen will react with 96.1 g

of propane, C3H8, to produce carbon dioxide and water vapor?

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

Calculations involving a Limiting Agent

Limiting Reactant Reactant that is consumed first and limits the number

of products

Nitrogen gas can be prepared by passing gaseous ammonia over solid copper (II) oxide at high temperatures. The other products are solid copper and water vapor. If a sample containg 18.1 g of NH3 is reacted with 90.4 g of CuO, which is the limiting agent? How many grams of N2 will be formed?

Calculating Percent Yield Theoretical Yield- amount of a product

formed when a limiting reactant is completely consumed

% yield= actual yield/theoretical yield • 100

Percent yield example Methanol(CH3OH) can be manufactured by

a combination of gaseous carbon monoxide and hydrogen. Suppose 68.5 kg CO(g) is reacted with 8.60 kg H2(g). Calculate the theoretical yield of methanol. If 3.57 x 104 g of methanol is actually produced, what is the % yield of methanol?