1. Chemical Reactions & Stoichiometry
UNIT 4
Chemical Reactions & Stoichiometry

2. Chemical Equations
A chemical equation represents, with symbols and formulas, the reactants and products in a chemical reaction
Reactants are shown on the left hand side of the arrow
Products are shown on the right hand side of the arrow
The arrow means “produces”, “yields”, or “reacts to form”
Chemical equations may also include abbreviations to denote the state of matter of the reactants and products:
(s) for solid
(l) for liquid
(g) for gas
(aq) for aqueous, which means dissolved in water (in solution)

3. HC2H3O2(aq) + NaHCO3(s)  CO2(g) + H2O(l) + NaC2H3O2(aq)
N2(g) + H2(g)  NH3(g)
HC2H3O2(aq) + NaHCO3(s)  CO2(g) + H2O(l) + NaC2H3O2(aq)
The equations above are called formula equations because they represent the reactants and products of a reaction using their symbols and formulas
A word equation can be written using the names of the reactants and products

4. Why balance equations? 2H2O  2H2 + O2
In any chemical reaction, mass is conserved (meaning that the mass of the products must be equal to the mass of the reactants)
A balanced chemical equation shows how many moles of each element or compound react to exactly form the products
Coefficients are placed in front of formulas to show how many moles of that element/compound are needed in the equation
Coefficients must be whole numbers
Example:
2H2O  2H2 + O2

5. To balance a chemical equation
Steps to balance an equation:
Determine how many atoms of each element are present on the reactant and product sides
Use trial and error to add coefficients to balance out the atoms on both sides
When you think you are done, check your result to make sure that each element is balanced

6. Fe(OH)3 + BaCl2  FeCl3 + Ba(OH)2
Example #1
H2 + Cl2 HCl
Example #2
Fe(OH)3 + BaCl2  FeCl3 + Ba(OH)2

7. Example #3
Magnesium chloride reacts with silver nitrate to form solid silver chloride and magnesium nitrate

8. Types of Chemical Reactions
There are many different types of chemical reactions that will be studied throughout the year
In this unit, we will investigate 5 main types of reactions
Synthesis
Decomposition
Single Replacement
Double Replacement
Combustion

9. Synthesis Reaction Mg + O2  MgO
In a synthesis reaction, two or more reactants come together to form one product
EXAMPLE:
Magnesium and oxygen react to form magnesium oxide
Mg + O2  MgO

10. Decomposition Reaction
In a decomposition reaction, a single reactant breaks apart to form more than one product (the reverse of a synthesis reaction)
EXAMPLE:
Silver oxide decomposes into solid metallic silver and oxygen gas
Ag2O  Ag + O2

11. Decomposition of Nitrogen Triiodide

12. Single Replacement Reaction
In a single replacement reaction, one element replaces another element in a compound, leaving the element by itself
EXAMPLE:
Aluminum reacts with copper(II) chloride to form copper metal and aluminum chloride
Al + CuCl2  Cu + AlCl3

13. Double Replacement Reaction
In a double replacement reaction, two elements switch places, forming two new ionic compounds
EXAMPLE:
Silver nitrate solution reacts with sodium chloride solution to form solid silver chloride and aqueous sodium nitrate
AgNO3 + NaCl  AgCl + NaNO3

14. Demonstration Barium Nitrate with Potassium Sulfate
Video: -I

15. and the fifth type of reaction is…

16. COMBUSTION!!

17. Combustion Reaction hydrocarbon + O2  CO2 + H2O
In a combustion reaction, a hydrocarbon (or compound containing carbon, hydrogen, and/or oxygen) reacts with oxygen gas, forming carbon dioxide and water
hydrocarbon + O2  CO2 + H2O
Many different types of carbon compounds can provide the “fuel” for a combustion reaction…

18. Example __CH4 + __O2  __CO2 + __H2O
Methane reacts with oxygen gas, forming carbon dioxide and water
__CH4 + __O2  __CO2 + __H2O

19. Practice
Write and balance equations for the following combustion reactions:
Combustion of hexane (C6H14)
Combustion of methanol (CH3OH)

20. Predicting Products of Chemical Reactions

21. Aqueous Solutions and Dissolving
When ionic compounds are in the aqueous state, they have dissolved in water and their ions have disassociated (split apart from each other)
Dissolving Salt on a Molecular Level
NaCl (aq)  Na+ + Cl-
CuCl2 (aq)  Cu2+ + 2Cl-

22. Connection to Polarity
Polar water molecules are attracted to the positive and negative ions and those attractions are what causes a substance to dissolve
Some substances are too attracted to each other that water is not strong enough to pull the ions apart. These compounds are what we call insoluble and they are the precipitates in double replacement reactions

23. Products for Double Replacement
Two aqueous solutions come together.
Often a precipitate or water is formed.
The ions keep the same charges! They just switch partners
Example:
NaOH + CuCl2  ?
Steps:
Determine the charges for the cations and anions in the reactant compounds
Swap partners
Balance the charges of the new partners to see what the formulas for the products are
Balance the entire reaction (if it doesn’t work- check your formulas)

24. Practice
For the following, predict the products formed, and balance the overall equation.
CuSO4 + Ba(NO3)2 
Mg(OH)2 + HCl 
CaCl2 (aq) + Na3PO4 (aq) 

25. Products for Single Replacement
Usually one aqueous reactant and one metal or nonmetal reactant.
Example:
Mg (s) + CuCl2 (aq)  ?
Steps:
1.Find the charges of the species that are aqueous.
2. Replace a metal with metal or nonmetal with nonmetal and give them the appropriate charge based on the periodic table
3. Determine the formula for the compound being formed based on the charges of the cation and the anion
4. Determine if the species becoming an element is diatomic or not (HOFBrINCl)

26. Practice
K + ZnCO3 
Cl2 + AlBr3 
H2O + Li 

27. Products for Synthesis and Decomposition
In a synthesis reaction, reactants combine to form the product
If the reactants include 1 metal and 1 nonmetal- you should be able to predict the product based on what ionic compound would be formed. Determine the charges that those elements have as ions, and balance them to determine the formula.
ex. zinc reacts with oxygen gas
In a decomposition reaction, the reactant breaks apart into elements or common, stable compounds
You will often be told what elements or compounds the decomposition results in.
ex. Calcium carbonate decomposes, forming calcium oxide and carbon dioxide.

28. Products for Combustion
In a combustion reaction, the other reactant is always oxygen gas, and the products are carbon dioxide and water
The combustion of butane (C4H10)

29. Element reactivity
In single replacement reactions, the reaction will only occur if the single element that is reacting is more reactive than the element it is replacing in the compound
For example:
Al + CuCl2  Cu + AlCl3 - this reaction occurs (you saw it)
Cu + AlCl3  Al + CuCl2 - this reaction does not occur
This is because aluminum is more reactive than copper
If the reaction is:
then the striped atom must be more reactive than the blank one

30. Element reactivity (cont.)
An activity series is a list of elements according to how reactive they are, from most reactive to least reactive
In a single replacement reaction, elements that are higher on the list will replace elements that are lower on the list in compounds
Example: Which reaction will occur?
Ag + CoBr2  Co + AgBr OR
Co + AgBr  Ag + CoBr2
Answer: the second reaction will occur, because Co is more reactive than Ag

31. Practice
For each set of reactions below, use the activity series to determine which reaction will occur.
ZnSO4 + Na  Na2SO4 + Zn OR
Na2SO4 + Zn  ZnSO4 + Na
CoCl2 + Sn  SnCl2 + Co OR
SnCl2 + Co  CoCl2 + Sn
CaF2 + I2  CaI2 + F2 OR
CaI2 + F2  CaF2 + I2

32. Practice
For each reaction below, use the activity series to determine if the reaction will occur. If it will occur, write the formulas for the products and balance the equation.
ZnO + Sr 
Pb + Mg(NO3)2 
Fe2(SO4)3 + Ca 
Ni(OH)2 + Na 
NaCl + F2 
MgBr2 + I2 

33. Stoichiometry

34. What is Stoichiometry?
Using a balanced equation to predict relative amounts of reactants and products
EXAMPLE:
3H2 + N2  2NH3
How many molecules of hydrogen and nitrogen have to react to create 10 molecules of ammonia (NH3)?

35. Mole Ratios
We can think about the coefficients as ratios of molecules, but we can also think about them in terms of ratios of moles (since a mole is 6.022x1023 molecules)
3H2 + N2  2NH3
How many moles of nitrogen gas and hydrogen gas will it take to create 1 mole of ammonia gas?

36. Practice
How many moles of water can be formed by reacting 3 moles of hydrogen gas with an excess of oxygen gas?
Hint: Write out a balanced equation!

37. Limiting Reactant/Reagent
Lab reactions are rarely carried out in exact amounts that react with each other; usually, one reactant is used up first, causing the reaction to stop occurring
The limiting reactant is the reactant that limits the amount of the other reactants that can combine and the amount of products formed in a chemical reaction
The substance(s) that is not used up completely in a reaction is called the excess reactant.

38. An analogy, if you will…
A manufacturer of bicycles has 4815 wheels, 2305 frames, and handlebars.
How many bicycles can be manufactured using these parts?
How many parts of each kind are left over?
Which part limits the production of bicycles?

39. Practice 3H2 + N2  2NH3 If 7 moles of H2 react with 2.5 moles of N2…
Which is the limiting reactant?
How many moles of ammonia are formed?

40. Demonstration: Limiting Balloons!
Record the chemical reaction being performed:
Record observations (a picture can be helpful) to document what happens in the demonstration:

41. Demonstration, cont.
How did the five trials of this experiment differ?
Identify the limiting and excess reagents in each flask.
From this demo, could you determine the exact amount of baking soda to react with 100 mL? Why or why not?

42. Example 1
Silicon dioxide (quartz) is usually quite stable but reacts with hydrogen fluoride according to the following equation: SiO2 + 4HF  SiF4 + 2H2O If 2.0 mol of HF is combined with 4.5 mol of SiO2, which is the limiting reactant?

43. Using Mole Ratios in Stoichiometry
Use dimensional analysis with mole ratios as a conversion factor to convert between moles of different substances
Mole ratio = coefficients in balanced equation
2Al(s) +3 CuCl2(aq)  3Cu(s) + 2AlCl3(aq)
How many moles of copper chloride will react with 0.45 moles of aluminum?
0.45 moles Al x 3 moles CuCl2 = moles CuCl2
2 moles Al

44. Example
Note: excess just means that there is plenty of it to react with the Aluminum. Aluminum must be the limiting reactant.
2Al(s) +3 CuCl2(aq)  3Cu(s) + 2AlCl3(aq)
How many moles of Cu will be formed if 3.29 moles of Aluminum react with an excess of copper chloride? Show work with dimensional analysis.
How many moles of Copper chloride reacted if moles of Aluminum reacted? Show work with dimensional analysis.

45. Practice
Hydrochloric acid (HCl) can react with oxygen from the air to produce dangerous chlorine gas and water.
Write the balanced equation
If moles of Hydrochloric acid completely react with oxygen gas, how many moles of chlorine gas are produced? Show work with dimensional analysis

46. Stoichiometry with grams
We don’t usually measure the number of moles in a lab. We have scales that measure in grams. In order to compare relative amounts of reactants and products- we need to be in moles or molecules. 2 H2 + O2  2H2O Why can’t we say that 2 grams of Hydrogen react with 1 gram of oxygen to form 2 grams of water?

47. Converting from grams to moles
Use the Molar Mass!
Molar mass is the number of grams 1 mole of that element or compound weighs.
How many grams of hydrogen react with how many grams of oxygen in 2 H2 + O2  2H2O ?

48. General Rules of stoichiometry
If in grams- convert to moles using the molar mass of that substance
Use mole ratios (the coefficients in the balanced equation) to covert from moles of one substance to moles of another
If final answer needs to be in grams- convert back to grams using the molar mass of that substance. If it asks for moles- leave it in moles!
(Coefficients in Balanced equation)

49. Example 1
Hydrogen and oxygen gases can combine to form water when electricity is added.
Write the balanced equation
64 grams of hydrogen gas will react with how many moles of oxygen gas when forming water?
64 grams of hydrogen gas will produce how many grams of water?

50. Example 2
In the production of ammonia (NH3) nitrogen and hydrogen gases combine.
Write the balanced equation
If 5.3 grams of hydrogen gas react with an excess of nitrogen gas, how many grams of ammonia are formed?
How many moles of nitrogen gas react with this 5.3 grams of hydrogen when producing ammonia?

51. Theoretical Yield
The theoretical yield of a reaction is how many grams of the product is mathematically supposed to be formed assuming a complete reaction
Every time we do a calculation to calculate the number of grams of a product- we are finding the theoretical yield.
Why is it called theoretical? Because it didn’t actually happen yet- it is just what is mathematically supposed to happen.

52. Actual Yield or Observed Yield
The actual (or observed) yield is what was actually created when the reaction occurred in the lab.
The actual yield might be too high (has contaminants)
The actual yield might be too low (some of the product was lost while carrying out the lab, or the reaction didn’t go to completion)

53. Percent Yield
We can calculate the percent yield to see what percent of the theoretical yield was actually made.
If percent yield is below 100%- did not create the entire product
If percent yield is above 100%- something else must be in your product!

54. Example
Aluminum was reacted with copper (II) chloride according to the following equation:
___Al + ___CuCl2  ___AlCl3 + ___Cu
If 5.3 grams of aluminum react with excess copper chloride, what is the theoretical yield of copper?
In the lab, you filter and dry the product, and find that you have produced 10.2 grams of Copper. What is the percent yield?