Particles (Atoms, Molecules, and Ions) Can Be Counted Three Ways by Using a Unit Called

Stoichiometry

What are some ways we can measure matter? o Counting o Weighing o Volume

Particles (atoms, molecules, and ions) can be counted three ways by using a unit called the mole (mol).

Number of particles

Mole Mass of particles

Volume of particles A) Avogadro’s Number

Representative particles: the smallest pieces of a substance Substance Representative Particle Covalent (molecular) compound Molecule Ionic compound Formula Unit (ionic compound) Element Atom

 one mole of any substance (element or compound) always contains the same number of particles.  1 mole = 6.02 x 1023 particles = Avogadro’s number

Note: 1 mole of an element = 6.02 x 1023 atoms 1 mole of a molecular compound = 6.02 x 1023 molecules 1 mole of an ionic compound = 6.02 x 1023 formula units

Mole: the number equal to the number of carbon atoms in exactly 12 grams of pure 12C.

So, in 12 grams of 12C, there are 6.02 x 1023 atoms of carbon. Techniques such as mass spectroscopy, which count atoms very precisely, have been used to determine this number. 1. How many atoms are in 1.75 mol of calcium?

2. How many moles are in 2.42 x 1021 molecules of water?

3. How many atoms are present in 1.69 mol of PCl3? B) Molar Mass  remember that the atomic mass of an atom is given in the periodic table and has the unit amu. Similarly, the molar mass of an element is the number of grams of that element that is equivalent to its atomic mass. Carbon

12.0 amu 12.0 g (Atomic mass) (Molar mass)

 in a compound, the molar mass is the sum of all the elements.

H2SO4 molecular mass molar mass H x 2 = 2.0 amu H x 2 = 2.0 g S x 1 = 32.1 amu S x 1 = 32.1 g O x 4 = 64.0 amu O x 4 = 64.0 g 98.1 amu 98.1 g

one molecule of H2SO4 one mole of H2SO4 has a mass of 98.1 amu has a mass of 98.1g

***** 1 mole = molar mass in grams its units are g/mol Try these:

1. How many grams of each element should you measure out (to the nearest tenth) in order to get one mole of a) oxygen (O) b) carbon (C) c) calcium (Ca) d) phosphorus (P) e) nitrogen (N) f) chlorine (Cl)

2. How many grams of each compound should you measure out (to the nearest tenth) in order to get one mole of

a) calcium chloride (CaCl2)

b) calcium carbonate (CaCO3)

c) calcium nitrate (Ca(NO3)2)

d) calcium phosphate (Ca3(PO4)2)

3. How many grams are in 6.95 mol of magnesium oxide (MgO)?

4. How many moles in 375 g of nitric acid (HNO3)?

5. How many grams in 4.2 mol of silver nitrate?

6. How many moles in 2.15 g of lithium oxalate? INTERPRETING CHEMICAL EQUATIONS

 the starting point for any problem involving quantities of chemicals in a reaction is the balanced equation.  analyze the meaning of this balanced equation:

N2(g) + 3H2(g)  2NH3(g)

Reactants Products Conserved? Atoms 2 + 6 8 yes Molecules 1 + 3 2 no Moles 1 + 3 2 no Mass (g) 28.0 + 6.0 34.0 yes

 therefore, we have the law of conservation of matter (mass). Matter (mass) is conserved during any chemical change.  calculations using balanced equations are called stoichiometric calculations.

TRY THIS: Is mass conserved when magnesium chloride reacts with calcium hydroxide? MOLE RELATIONSHIPS A) MOLE RATIOS  balanced equations tell us the number of moles of each chemical taking part in the reaction. The ratio that gives moles of one chemical to moles of another chemical is called a mole ratio.

There are 6 mole ratios that can be made from this reaction:

N2(g) + 3H2(g)  2NH3(g)

i) 1 mol N2 ii) 3 mol H2 iii) 1 mol N2

3 mol H2 1 mol N2 2 mol NH3

iv) 2 mol NH3 v) 3 mol H2 vi) 2 mol NH3

1 mol N2 2 mol NH3 3 mol H2

 these mole ratios are used as conversion factors to help solve mole-mole problems.

moles of A  moles of B

Ex. Methanol fuel burns according to this equation:

2CH3OH + 3O2  2CO2 + 4H2O

If 3.50 mol of methanol are burned in lots of oxygen, a) how many moles of oxygen are used? b) how many moles of water are produced? B) MASS-MASS  the procedure for calculating the masses, in grams, of reactants or products from a balanced chemical equation is called gravimetric stoichiometry. One example of this is the mass-mass problem. The known mass of one chemical (A) is used to find the unknown mass of a second chemical (B). 1) mass of A  moles of A

2) moles of A  moles of B

3) moles of B  grams of B

Examples: 1) If 2.85 g of hydrogen gas is reacted with oxygen gas in the

reaction 2H2(g) + O2(g)  2H2O(l)

a) how many grams of oxygen gas will be needed? b) how many grams of water will be produced? 2) The equation for the roasting of iron(II) sulfide is

2FeS + 3O2  2FeO + 2SO2

If we started with 95 g of iron(II) sulfide, then, a) what mass of oxygen gas will be needed? b)what mass of sulfur dioxide gas will be produced? 3) One type of antacid is magnesium hydroxide:

Mg(OH)2 + 2HCl  MgCl2 + 2H2O

If an antacid tablet has a mass of 4.56 g, then a) what mass of hydrochloric acid will it react with? b) what mass of water will it make? Limiting Reagents

When chemicals are mixed together so they can undergo a reaction, they are often mixed in stoichiometric quantities.

They are used in exactly the correct amounts so that all the reactants "run out" or are used up at the same time.

Hydrogen is produced for use in the manufacture of methane gas.

CH4 (g) + H2O (g)  3 H2 (g) + CO (g)

If 100 g of water is used to react with 249 g of methane, how much hydrogen gas is produced? Because the amount of water limits the amount of products that can be formed, it is called the limiting reagent.

Where reactants are not mixed in the stoichiometric quantities, it is important to determine the limiting reagent in order to calculate correctly the amounts of products that will be formed. Steps for Solving Stoichiometry Problems Involving Limiting Reagents

1. Write a balanced equation for the reaction.

2. Convert known masses of reactants to moles.

3. Using the numbers of moles of reactants and the appropriate mole ratios, determine which reactant is limiting.

4. Using the amount of the limiting reactant and the appropriate mole ratios, compute the number of moles of the desired product.

5. Convert from moles of product to grams of product, using the molar mass. Ex 1. Suppose 25.0 kg of nitrogen gas and 5.00 kg of hydrogen gas are mixed and reacted to form ammonia. Calculate the mass of ammonia produced when this reaction is run to completion.

Ex 2. Nitrogen gas can be prepared by passing gaseous ammonia over solid copper II oxide at high temperatures. The other products of the reaction are solid copper and

water vapor. How many grams of N2 are formed when

18.1 g of NH3 are reacted with 90.4 g of CuO? The Yield of a Chemical Reaction theoretical yield - the maximum amount of a product that can form in a chemical reaction. It is calculated by assuming that all of the limiting reagent has reacted to form product.

This assumption may not be correct for a number of reasons. 1. Many reactions do not go to completion. 2. There many be other reactions (side reactions) which occur in addition to the main reaction. 3. Some product may be lost during the process of separating it from the reaction mixture. actual yield - the amount of product that is actually obtained from a chemical reaction.

Theoretical - calculated quantity Actual - experimentally determined quantity

Percent yield - the actual yield expressed as a percentage of the theoretical yield.