UNIT 5: STOICHIOMETRY

UNIT 5: STOICHIOMETRY

Outline The Mole Molar Mass, Mass and atoms Molar Mass of Compounds Empirical Formula, Molecular Formula (Not Hydrates) Stoichiometry, Mole Ratios Limiting Reactants, Percent Yield Solution Concentrations

It s hard to count each molecule An atom or a molecule is SO SMALL There are about 100,000,000 atoms in a 1cm line. (100 billion) Chemists use mililiters and liters of chemicals in beakers and flasks. How can we count all these molecules?

Common Words for numbers A pair of gloves: individual gloves A couple of people: individual people A dozen eggs: individual eggs Two dozen eggs: individual eggs A ream of paper: sheets of paper

The Mole Just like a dozen means 12, the word mole means: 602,213,670,000,000,000,000,000 or abbreviated as 6.022 x 10 23 in chemistry. This number is referred to as Avogadro s number, in honor of the Italian scientist Avogadro who found the volume of 1 mol of gas.

The Mole So in chemistry, we use the mole, (abbreviated mol ) to count the number of atoms and molecules. Eg. I have 1 dozen apples. How many apples is this? Eg. I have 2 dozen apples. How many apples is this? Eg. I have 1 mol of apples. How many apples is this? Eg. You have 2 mol of apples. How many apples is this?

The Mole Eg. I have 0.5 mol of apples. How many apples do I have? Eg. There is 1 mol of C atoms in this beaker. How many C atoms are there? Eg. There is 0.5 mol of H 2 molecules in this balloon. How many H 2 molecules are in the balloon? How many H atoms are in this balloon?

The Mole Textbook p. 323 #1-4 p.324 #5-6

The Mass of a Mole How much mass does a dozen eggs have? (1 egg = 60g) How much mass does a dozen pencils have? (1 pencil = 20g) The mass of 12 eggs is not the same as the mass of 12 pencils.

The Mass of a Mole Similarly, the mass of 1 mole of carbon (as an example) is not the same as the mass of 1 mole of some other chemical. Eg. The mass of 1 mol of Helium gas is not the same as the mass of 1 mol of Uranium.

The Mass of a Mole: Molar Mass The mass (in grams) of one mole of a substance is called its molar mass. The mass numbers given in the periodic table (eg. Iron (Fe) has a mass number given as 55.845) is equal to the mass in grams of 1 mol of that atom.

he Mass of a Mole: Molar Mass The mass (in grams) of one mole of a substance is called its molar mass. n = m n = number of moles (units mol) m = mass (units grams, g) MM = molar mass (units g/mol) MM Textbook p. 328-331 #15-21

The Mass of a Mole: Molar Mass Same concept for compounds. A compound (two or more atoms joined by ionic or covalent bonding) is a pure substance. Eg. NaCl, Ca(OH) 2, H 2 O are compounds. Molar mass for a compound is the sum of the molar masses of each atom in the formula.

The Mass of a Mole: Molar Mass Textbook p.335 #29-36 p.336-339 #37-46

Percent by Mass How much of X is in 100g of the compound XY? If you knew there were 55g of X and 45g of Y in a 100g sample of XY, you would that there is 55% of X and 45% of Y in the compound. The percent by mass of each element in a compound is the percent composition.

Percent by Mass 1. Pretend you have 1 mol of the compound. 2. Look at the chemical formula, and write down how many moles you have of each element in the compound. Now calculate the mass of each of these amounts for each element.

Percent by Mass 3. Calculate the Molar Mass of the compound. 4. Now calculate the how much mass one of the elements in the compound has in relation to the mass of the whole compound.

Percent by Mass eg. NaHCO 3 Textbook p.344 #54-57

Empirical Formula

Empirical Formula Opposite to calculating the percent by mass If you have an unknown chemical, you run it through an analysis machine and you only get the percent composition. You, the chemist, need to determine what the chemical is. The formula you get from the percent composition is called the empirical formula.

Empirical Formula The empirical formula is the smallest whole-number mole ratio of the elements in a compound. An empirical formula IS NOT the real formula of a chemical, but it is important because it gives you a good guide to find the real formula of the chemical.

Empirical Formula STEPS Example: A 250g sample of a substance is made up of 40. % S and 60.% O. What is the empirical formula of the substance? 1. Calculate the mass of each element with the given percentages. 2. Calculate the number of moles of each element from the masses. 3. Divide the moles by the smallest mole number you calculated. 4. If there are still some that aren t whole numbers, multiply everything by a factor that will make it a whole number. Textbook p.346 #58-61

Molecular Formula The molecular formula is the real formula of a substance. Textbook p.350 #62-66 The molecular formula is always a factor of the empirical formula. Can get the molecular formula if you know the unknown substance s molar mass and its empirical formula. 1. Divide molar mass given by the empirical formula to get a factor. 2. Multiply the factor to each element in the empirical formula.

STOICHIOMETRY (Chapter 11) Have you ever wondered how much Oxygen and Carbon dioxide we need to form 500 grams of sugar? What about 1000grams? 700 grams?

STOICHIOMETRY The study of quantitative (numerical) relationships between the amounts of reactants used and the amounts of products formed by a chemical reaction is called stoichiometry. Law of conservation of mass: You are not a wizard. You cannot create or destroy mass.

STOICHIOMETRY Law of conservation of mass: In any chemical reaction, the amount of matter present at the end of the reaction is the same as the amount of matter present at the beginning. Example: Iron metal + oxygen gas makes iron oxide.

STOICHIOMETRY Mole ratio: A mole ratio is a ratio between the number of moles of any two substances in a balanced chemical equation. Example: what is the mole ratio of iron to iron oxide in prev. q? What is the mole ratio of iron to oxygen gas? Oxygen to iron oxide? etc.

STOICHIOMETRY Using mole ratios: You can use these ratios to find out how many moles of something is formed, or how many moles of something is needed!

STOICHIOMETRY Using mole ratios: Example 2K(s) + 2H 2 O(l) 2KOH(aq) + H 2 (g) How much hydrogen is produced if only 0.0400 mol of potassium is dropped in the water? https://www.youtube.com/watch?v=oqmn3y8k9so Set up mole ratio equation or unit cancelling equation. Textbook p. 375, 376, 377

STOICHIOMETRY Limiting Reactants Why do reactions stop? Because one or all of the reactants have been used up in the reaction! Need butter, eggs, flour to bake a cake. You can t bake more cake if you run out of flour, for example.

STOICHIOMETRY Limiting Reactants The limiting reacting is the reactant that is used up completely first. The reactants that are left over are called the excess reactants. Cake example again: lets say you need 2 eggs, 1 butter stick, and 1 cup of flour for one cake. You have 2 boxes of eggs (60 eggs), 50 butter sticks and 5 cups of flour. What is the limiting reactant? Which are the excess?

STOICHIOMETRY Limiting Reactants So if you want to know how much product you will get, you have to find out which reactant is the limiting reactant. You can only find this by comparing the mole ratios in the balanced equation to the amount of moles that you are given in the question (that you calculate)

STOICHIOMETRY Limiting Reactants: Steps 1. Calculate mole ratio of the reactants. 2. Choose one reactant and find out how many moles of the other reactant you need to completely react with your chosen reactant. 3. The reactant that you do not have enough of (in the question) is the limiting reactant.

STOICHIOMETRY Limiting Reactants: FAST METHOD Divide the number of moles of each reactant by the reactant s coefficient in the balanced chemical equation. The smaller answer is the limiting reactant!

STOICHIOMETRY Theoretical Yield Once you know the limiting reactant, you will know how much product you will get. You do this by calculating the number of moles of product using the mole ratio of the limiting reactant and the product.

STOICHIOMETRY Theoretical Yield The theoretical yield is the maximum amount of product that can be produced from a given amount of reactant. 1. Write the mole ratios of the limiting reactant and the desired product. 2. Calculate the number of moles of the product that will form from the given number of moles of the limiting reactant.

STOICHIOMETRY Percent Yield The theoretical yield is the maximum amount of product that can be produced from a given amount of reactant. (Calculated amount of product from the amount of Limiting Reactant.) But the actual yield is the amount of product that is actually produced when the chemical reaction is carried out in an experiment.

STOICHIOMETRY Percent Yield The percent yield is the ratio of the actual yield to the theoretical yield, expressed as a percent. P.387 #28-30 P. 388 #34-35

STOICHIOMETRY: Solutions Ch.14 Expressing Solution Concentration (14.2) The Solvation Process (14.3) Colligative Properties (14.4)

STOICHIOMETRY: Solutions Ch.14 A solution is a homogeneous (uniform) mixture. Air is a solution Coffee is a solution The concentration of a solution is a measure of how much solute is dissolved in a specific amount of solvent.

STOICHIOMETRY: Solutions Ch.14 Key words: Solute: Substance present in smaller numbers in the solution Solvent: Substance present in greater numbers in the solution

STOICHIOMETRY: Solutions Ch.14 Concentrated tea Dilute tea

STOICHIOMETRY: Solutions Ch.14 What is the problem with using only the words dilute and concentrated to describe the concentration of a solution?

STOICHIOMETRY: Solutions Ch.14 So to describe the concentration of a solution accurately, we can use the following: Percent by mass Percent by volume Molarity Molality

STOICHIOMETRY: Solutions Ch.14 Molarity, M Moles of solute per liter of solution. P.483 #16-19 Molarity problems p.484 #20-23 Dilutions p.486 #24-26

STOICHIOMETRY: Solutions Ch.14 Molality Moles of solute per kg of solution. P.487 #27-28

STOICHIOMETRY: Solutions Ch.14 The Solvation Process When a solute gets separated from its other solute particles, and gets surrounded by the solvent particles to form a solution, this process is called solvation.

STOICHIOMETRY: Solutions Ch.14 Factors Affecting Solvation 1. Agitation: stirring moves dissolved solute particles away from the contact surfaces more quickly, and allows new collisions between solute and solvent particles.

STOICHIOMETRY: Solutions Ch.14 Factors Affecting Solvation 2. Surface Area: Breaking solute into small pieces increases surface area, which allows more collisions to occur.

STOICHIOMETRY: Solutions Ch.14 Factors Affecting Solvation 3. Temperature: Increasing temperature increases the rate of solvation (dissolves more quickly).

STOICHIOMETRY: Solutions Ch.14 Solubility Amount of solute that dissolves in the solvent. Units are often g/l or mol/l. Saturated solution: contains maximum amount of dissolved solute for the amount of solvent at the T and P.

STOICHIOMETRY: Solutions Ch.14 Colligative Properties When a solute is dissolved in a solvent, the physical properties of the solution changed. It didn t matter what the solute was.

STOICHIOMETRY: Solutions Ch.14 Colligative Properties 1. Vapor Pressure Lowering: Less of the solvent particles exist in the gas state above the liquid. (Some solute particles are at the surface of the liquid, so less solvent particles become gas)

STOICHIOMETRY: Solutions Ch.14 Colligative Properties 2. Boiling Point Elevation: The solution must be heated to a higher temperature to boil.

STOICHIOMETRY: Solutions Ch.14 Colligative Properties 3. Freezing Point Depression: Freezing point is lower than that of the pure solvent.

STOICHIOMETRY: Solutions Ch.14 Phase Diagrams (12.4) Most substances exist in three states, depending on the temperature and pressure. Phase diagram is a graph of pressure versus temperature that shows the phases a substance exists for different T and P.

STOICHIOMETRY: Solutions Ch.14

What is the critical temperature of compound X? If you were to have a bottle containing compound X in your closet, what phase would it most likely be in? At what temperature and pressure will all three phases coexist? If I have a bottle of compound X at a pressure of 45 atm and temperature of 100 C, what will happen if I raise the temperature to 400 C? Why can t compound X be boiled at a temperature of 200 C? If I wanted to, could I drink compound X?