2. Bell Work “We’re Cookin’ Now”

3. Chocolate Chip Cookies!! 1 cup butter 1/2 cup white sugar 1 cup packed brown sugar 1 teaspoon vanilla extract 2 eggs 2 1/2 cups all-purpose flour 1 teaspoon baking soda 1 teaspoon salt 2 cups semisweet chocolate chips Makes 3 dozen How many eggs are needed to make 3 dozen cookies? 2 eggs How much butter is needed for the amount of chocolate chips used? 1 cup How many eggs would we need to make 9 dozen cookies? 6 eggs How much brown sugar would I need if I was cutting the recipe in half? 1/2 cup

4. Cookies and Chemistry…Huh!?!? Just like chocolate chip cookies have recipes, chemists have recipes as well Instead of calling them recipes, we call them chemical equations Furthermore, instead of using cups and teaspoons, we use moles Lastly, instead of eggs, butter, sugar, etc. we use chemical compounds as ingredients

5. Stoichiometry

6. Just what is stoichiometry? The word stoichiometry is derived from two Greek words: stoicheion (meaning element) and metron (meaning measure).

7. Jeremias Benjamin Richter (1762- 1807) was the first to lay down the principles of stoichiometry. In 1792 he wrote: “Stoichiometry is the science of measuring the quantitative proportions or mass ratios in which chemical elements stand to one another.” This was verified by Antoine Lavoisier with the Law of Conservation of Mass

8. Why do we use stoichiometry? To reduce waste To predict amounts of products To make money

9. How does it work? Using balanced equations we can calculate the amounts of substances in the chemical reaction. Since stoichiometry is used to predict amounts where there are no error, we can expect that the reaction in the lab won’t create the same amount of product predicted.

10. How do I use the chemical equation? Looking at a chemical equation tells us how much of something you need to react with something else to get a product (like the cookie recipe) The chemical equation is written like a mathematical expression. When balanced, the coefficients tell how many moles of each reactant or product is needed or expected. The important thing is that the equation must be BALANCED!!!!

11. Example: 2H 2 + O 2  2H 2 O One way to understand the information conveyed by a chemical equation is to convert the equation into an English sentence.

12. So: 2H 2 (g) + O 2 (g)  2H 2 O(l) (the coefficients tell how many moles, atoms or molecules of each chemical are needed in the “recipe”) Hydrogen gas reacts with oxygen gas to form liquid water. Now, we can get a bunch of relationships: molecules, atoms, moles, and masses The most important relationship is the mole relationship! We call that the MOLE RATIO

13. What is a mole ratio? Mole Ratios: The mole ratio is based on the balanced chemical equation. You will use the coefficients to get the mole ratio between two DIFFERENT substances

14. How do you use a mole ratio? The mole ratio is an equivalent which means that you can arrange the ratio in any way needed. The mole ratio is the KEY to calculations based upon a chemical equation. With it, you can calculate the amount of any other reactant in the equation and the maximum amount of product you can obtain.

15. Examples of Molar Ratios The expressing the ratio of moles in an equation using the coefficients Ex. 2H 2 + O 2 2H 2 O What is the molar ratio between H 2 and O 2 ? 2mol H 2 or 1mol O 2 1mol O 2 2mol H 2 What other mole ratios can we come up with?

16. Warning: The coefficients of a reaction only give the ratio in which substances react. They DO NOT in anyway tell you HOW MUCH is reacting.

17. Example of a mole ratio problem: N 2 + 3H 2  2NH 3 Write the molar ratios for N 2 and H 2. Write the molar ratios for NH 3 and H 2.

18. Stoichiometry has 5 basic steps: Step 1: Make sure the equation is balanced

19. Step 2: Write in all information given (make sure to identify what you are trying to find!)

20. Step 3: Convert everything to moles

21. Step 4: Use mole ratio to solve for what you are trying to find (you may also use a proportion to solve the problem.)

22. Step 5: Convert everything into the required unit if needed.

23. Moles to Moles One way to change ion ore, Fe 2 O 3, into metallic iron is to heat it together with hydrogen: Fe 2 O 3 + H 2  Fe + H 2 O How many moles of iron are made from 25 moles of Fe 2 O 3 ?

24. Moles to Moles Given: 25 moles Fe 2 O 3 Find: ? Moles Fe Ratio: What mole ratio will we use? 25 moles Fe 2 O 3 2mol Fe= 50 mol Fe 1mol Fe 2 O 3

25. Examples of a mole-mole stoichiometric problem: 2H 2 + O 2  2H 2 O 1. How many moles of H 2 O are produced when 5 moles of oxygen are used?

26. Examples of a mole-mole stoichiometric problem continued: 2H 2 + O 2  2H 2 O 2. If 3.00 moles of H 2 O are produced, how many moles of oxygen must be consumed?

27. Examples of a mole-mole stoichiometric problem continued: 2H 2 + O 2  2H 2 O 3. How many moles of hydrogen gas must be used, given the data in problem #2.

28. Mole-Mass Conversions Most of the time in chemistry, the amounts are given in grams instead of moles We STILL go through moles and use the mole ratio, but now we also use molar mass to get to grams Example: How many grams of chlorine are required to react completely with 5.00 moles of sodium to produce sodium chloride? 2 Na + Cl 2  2 NaCl 5.00 moles Na 1 mol Cl 2 70.90g Cl 2 2 mol Na 1 mol Cl 2 = 177g Cl 2

29. Practice Calculate the mass in grams of Iodine required to react completely with 0.50 moles of aluminum. 2 Al + 3 I 2  2 AlI 3

30. Mass-Mole We can also start with mass and convert to moles of product or another reactant We use molar mass and the mole ratio to get to moles of the compound of interest – Calculate the number of moles of ethane (C 2 H 6 ) needed to produce 10.0 g of water – 2 C 2 H 6 + 7 O 2  4 CO 2 + 6 H 2 0 10.0 g H 2 O 1 mol H 2 O 2 mol C 2 H 6 18.0 g H 2 O 6 mol H 2 0 = 0.185 mol C 2 H 6

31. Practice Calculate how many moles of oxygen are required to make 10.0 g of aluminum oxide 4 Al + 3 O 2  2 Al 2 O 3

32. Mass-Mass Conversions Most often we are given a starting mass and want to find out the mass of a product we will get (called theoretical yield) or how much of another reactant we need to completely react with it (no leftover ingredients!) 1. We must go from grams to moles by using molar mass of the compound given 2. Go from moles of compound given to moles of compound needed using the mole ratio from the coefficients 3. Go back to grams of compound we are interested in by using molar mass of the compound you are looking for

33. Mass-Mass Conversion Ex: Calculate how many grams of ammonia are produced when you react 2.00g of nitrogen with excess hydrogen. N 2 + 3 H 2  2 NH 3

34. Practice How many grams of calcium nitride are produced when 2.00 g of calcium reacts with an excess of nitrogen? 3 Ca + N 2  Ca 3 N 2

35. Gas Stoichiometry (Volume-Volume) 1. Take volume of gas given. 2. Convert to moles using the fact that 22.4 L=1 mole 3. Change from moles of compound given to moles of new compound using molar ratio from the coefficients. 4. Change moles of new compound to volume using 22.4 L = 1 mole

36. Gases and Stoichiometry 2 H 2 O 2 (l)  2 H 2 O (g) + O 2 (g) 3.5 L of H 2 O 2 was decomposed in a flask. What is the volume of O 2 made at STP? Bombardier beetle uses decomposition of hydrogen peroxide to defend itself.

37. Another Example Find the number of liters of oxygen necessary for the combustion of 6.7 L of magnesium, assuming the reaction occurs at STP. 2 Mg + O 2  2 MgO

38. LIMITING REACTANT

39. Limiting Reactant: Cookies 1 cup butter 1/2 cup white sugar 1 cup packed brown sugar 1 teaspoon vanilla extract 2 eggs 2 1/2 cups all-purpose flour 1 teaspoon baking soda 1 teaspoon salt 2 cups semisweet chocolate chips Makes 3 dozen If we had the specified amount of ingredients listed, could we make 4 dozen cookies? What if we had 6 eggs and twice as much of everything else, could we make 9 dozen cookies? What if we only had one egg, could we make 3 dozen cookies?

40. Limiting Reactant Most of the time in chemistry we have more of one reactant than we need to completely use up other reactant. That reactant is said to be in excess (there is too much). The other reactant limits how much product we get. Once it runs out, the reaction s. This is called the limiting reactant.

41. Limiting Reactant To find the correct answer, we have to try all of the reactants. We have to calculate how much of a product we can get from each of the reactants to determine which reactant is the limiting one. The lower amount of a product is the correct answer. The reactant that makes the least amount of product is the limiting reactant. Once you determine the limiting reactant, you should ALWAYS start with it! Be sure to pick a product! You can’t compare to see which is greater and which is lower unless the product is the same!

42. Limiting Reactant: Example 10.0g of aluminum reacts with 35.0 grams of chlorine gas to produce aluminum chloride. Which reactant is limiting, which is in excess, and how much product is produced? 2 Al + 3 Cl 2  2 AlCl 3 Start with Al: Now Cl 2 : 10.0 g Al 1 mol Al 2 mol AlCl 3 133.5 g AlCl 3 27.0 g Al 2 mol Al 1 mol AlCl 3 = 49.4g AlCl 3 35.0g Cl 2 1 mol Cl 2 2 mol AlCl 3 133.5 g AlCl 3 71.0 g Cl 2 3 mol Cl 2 1 mol AlCl 3 = 43.9g AlCl 3

43. LR Example Continued We get 49.4g of aluminum chloride from the given amount of aluminum, but only 43.9g of aluminum chloride from the given amount of chlorine. Therefore, chlorine is the limiting reactant. Once the 35.0g of chlorine is used up, the reaction comes to a complete.

44. Limiting Reactant Practice 15.0 g of potassium reacts with 15.0 g of iodine. Calculate which reactant is limiting and how much product is made. 2K + I 2  2KI

45. Percent Yield Theoretical Yield: The maximum quantity of product that can be obtained, according to the reaction stoichiometry, from a given quantity of reactant actually Actual Yield: The quantity of product actually obtained from experimentation Percentage Yield = Actual Yield Theoretical yield X 100%

46. What if I get an answer over 100%? Percent yield can NEVER be over 100% This would conflict with the Law of Conservation of Matter If you obtain a yield over 100% you either: – A. Did your calculation incorrectly by putting theoretical yield over percent yield OR – B. Made a mistake in the experiment

47. Percent yield problems: We calculated that 19.6g of KI forms from 15.0g K and 15.0g of I 2. If we ran the experiment and only obtained 18.5g of KI what is the percent yield?

49. Limiting Reactant & Percent Yield: Recap 1. You can recognize a limiting reactant problem because there is MORE THAN ONE GIVEN AMOUNT. 2. Convert ALL of the reactants to the SAME product (pick any product you choose.) 3. The lowest answer is the correct answer. 4. The reactant that gave you the lowest answer is the LIMITING REACTANT. 5. The other reactant(s) are in EXCESS. 6. To find the percent yield, divide the mass of what you get from the experiment from mass of the product from the limiting reactant and multiply by 100%.