3.  I can define stoichiometry.  I can identify the number of moles required in a reaction based on the coefficients.  I can determine how many moles of product will be produced based on the molar ratios found in an equation.

4.  Note Expectations:  Cell phones and electronics are not in use.  You are taking the notes.  You are helping the people at your table to answer the questions.  You are prepared to answer the questions.

5. 2 H 2 + O 2  2 H 2 O What are the numbers in red called? In the equation above, what is the ratio of hydrogen atoms to oxygen atoms? How did you figure that out?

6. 2 H 2 + O 2  2 H 2 O If you have 6 hydrogen atoms, how many oxygen atoms do you need to react with all of the hydrogen atoms? Why? If you have 8 hydrogen atoms and 4 oxygen atoms, how many water molecules would you make? How did you figure out your answer?

7.  Stoichiometry is the study of measurable relationships that exist in chemical formulas and equations.  It involves the relationships between reactants and products in a chemical reaction.

8. 2 H 2 + O 2  2 H 2 O When we talk about chemistry in the lab, are atoms practical? Besides atoms, what unit could the coefficients represent?

9. 2 H 2 + O 2  2 H 2 O If you have 8 moles of hydrogen, how many moles of oxygen do you need to react with all of the hydrogen atoms? Why?

10.  Stoichiometry connects the mole and coefficients in a balanced equation.

11.  Example:  N 2 H 4 + 2H 2 O 2  N 2 + 4H 2 O  This means that 1 mole of N 2 H 4 reacts with 2 moles of H 2 O 2 to produce 1 mole of N 2 and 4 moles of H 2 O.

12.  There is now a relationship between the coefficients and moles.  Molar ratios allow you to determine the number of moles of any substance in the reaction.

13.  Example:  NH4NO3  N2O + 2H2O  How many moles of N2O and H2O are made from 2.5 moles of NH4NO3?

14.  You can convert from moles of one substance to moles of another substance using mole ratios.

15.  Example:  Example: 4Fe + 3O 2  2Fe 2 O 3  How many moles of Fe are needed to react with 2.3 moles of O 2 ?

16.  Balanced equations follow the law of conservation of matter.  The mass of the products equals the mass of the reactants.

17.  What is stoichiometry?  Stoichiometry connects ____________ and __________________.  What is a molar ratio?  How does the mass of the products in a reaction compare to the mass of the reactants? Why?

19.  I can determine a volume ratio based on coefficients in a chemical equation.  I can determine the volume required for a reaction or produced as a product based on volume ratios.

20.  Note Expectations:  Cell phones and electronics are not in use.  You are taking the notes.  You are helping the people at your table to answer the questions.  You are prepared to answer the questions.

21.  What is the volume of 1 mole of gas at STP?  Is this a constant or does it change based on the type of gas?

22.  If 3.8 L of H 2 reacts with chlorine gas, what volume of HCl gas will be produced? H 2 + Cl 2  2HCl

23.  One mole of gas occupies the same volume as any other mole of gas.  Therefore, molar ratios in a chemical reaction represent the ratio of the volumes of gases.

24.  Volume-volume problems give you a volume and ask you to find the volume of an unknown.  You will use the volume ratio from the equation to solve these problems.

25.  What unit is used to measure volume?

26.  The volume ratio is the same as a molar ratio, except it uses liters.

27.  Example: N2 + 3H2  2HCl  What volume of H2 is required to react with 15.5 L of N2?

28.  Example:  If.38 L of H2 reacts with chlorine gas, what volume of HCl gas will be produced?  H2 + Cl2  2HCl

29.  Is volume of 1 mole of gas at STP a constant or does it fluctuate?  What is volume measured in?  What is the volume ratio?  How does the volume ratio compare to the molar ratio?

31.  I can determine the mass required or produced during a chemical reaction based on the molar ratios found in a chemical equation.

32.  Note Expectations:  Cell phones and electronics are not in use.  You are taking the notes.  You are helping the people at your table to answer the questions.  You are prepared to answer the questions.

33.  When we talk about moles, is mass constant for every molecule? Why or why not?  If I give you the mass of a reactant and ask you to fine the mass of a product, is this possible? Why or why not?

34.  Since a chemical equation is related to moles, you can now do mole conversions using the information in an equation.

35.  If 20 g of H 2 reacts with chlorine gas, how many grams of HCl will be produced? H 2 + Cl 2  2HCl Talk with your table about what you think the first step in solving this problem would be.

36.  If 20 g of H 2 reacts with chlorine gas, how many grams of HCl will be produced? H 2 + Cl 2  2HCl Now that we have established what to do in the first step, where are we going to go next?

37.  If 20 g of H 2 reacts with chlorine gas, how many grams of HCl will be produced? H 2 + Cl 2  2HCl What will the final step in this problem be?

38.  The first step is to always convert mass to moles, and then you can go anywhere from there.  The last step is to convert back to moles of the substance you are looking for.

39.  In a mass-mass problem, you are given the mass of one substance and asked to find the mass of another substance in a chemical reaction.  Remember the coefficients are not masses, they represent moles.

40.  What mass of water is produced from 1.5 grams of glucose (C 6 H 12 O 6 )? C 6 H 12 O 6 + 6O 2  6CO 2 + 6H 2 O

41.  When 3.8 g of Al 2 O 3 decomposes, how many grams of oxygen are produced? 2Al 2 O 3  4Al + 3O 2

42.  Is mass a constant for every molecule? Why or why not?  In a mass-mass problem what are you looking for?

44.  I can define a limiting reactant.  I can determine the limiting reactant for a reaction based on the amount of reactants used in a reaction.

45.  Note Expectations:  Cell phones and electronics are not in use.  You are taking the notes.  You are helping the people at your table to answer the questions.  You are prepared to answer the questions.

46.  Can chemical reactions go on and on forever? Why or why not?

47.  In chemical reactions, the amount of product formed is dependent on the amount of reactants available.  The amount of gasoline in your car limits how far you can drive.

48.  Predict the definition of a limiting reactant.  What does it mean if you have a reactant in excess?

49.  Often, in a chemical reaction, there is a reactant that limits the amount of product that can be formed.  A limiting reactant is one that limits the amount of product formed in a reaction.

50.  Which reactant will be used up in a chemical reaction, the limiting reactant or the reactant in excess?  Which reactant will determine how much product is made, the limiting reactant or the reactant in excess? Why?

51.  The limiting reactant will be completely used up in the reaction.  The other reactant or reactants will have some left over at the end of a reaction.

52.  The leftover reactant is said to be in excess.  The amount of product is always determined by the limiting reactant.

53.  If you have 3.5g of Cu and 6.0 g of AgNO3, which reactant is the limiting reactant in the production of silver? Cu + 2AgNO3  Cu(NO3)2 + 2Ag Predict how you would solve this problem. What are you going to look for?

54.  The reactant that produces the least amount of product is the limiting reactant.

55.  Identify the limiting reactant when 5.87 grams of Mg(OH) 2 reacts with 12.84 grams of HCl to produce MgCl 2 and water.  We are going to break these problems into three steps.

56. How much water is produced if you start with 5.87 grams of Mg(OH) 2 ? Mg(OH) 2 + 2HCl  MgCl 2 + 2H 2 O

57.  How much water is produced if you start with 12.84 grams of HCl? Mg(OH) 2 + 2HCl  MgCl 2 + 2H 2 O

58.  Circle the limiting reactant: Mg(OH) 2 or HCl.

59.  What is a limiting reactant?  What does it mean if a reactant is in excess?  Which reactant will determine how much product is made, the limiting reactant or the reactant in excess? Why?

61.  I can define expected yield.  I can define actual yield.  I can explain why the expected yield and actual yield are not the same.  I can define the percent yield.  I can calculate the percent yield for a reaction.

62.  The amount of product that is produced in a reaction is usually less than the predicted amounts.  The amount of product that should be produced based on calculations is the expected yield.

63.  The amount of product really obtained is the actual yield.  Product can be lost for several reason.  The reactants may not react.

64.  The reactants may be used up in side reactions that should not be occurring.  Some of the product may be lost due to transfer between containers.

65.  It is useful to determine what percent of the expected yield was actually obtained.  This percentage is called percent yield.

66.  Formula:  % yield = actual expected X 100

67.  Example: A piece of copper, with a mass of 5.00 g is placed in silver nitrate. The silver metal produced has a mass of 15.2 g but the expected yield was 17.0 g of Ag. What is the percent yield of silver?

68.  You know how to calculate the expected yield of product based on the calculations we have been doing.

69.  Example: Determine the percent yield for the reaction between 2.8 g Al(NO3)3 and excess NaOH if 0.966 g AL(OH)3 is recovered.  Al(NO3)3 + NaOH  Al(OH)3 + 3NaNO3 First calculate how much Al(OH)3 should have been produced, then calculate the percent yield.