1. Chapter 4: Types of Chemical Reactions & Solution Stoichiometry

2. WHY DO WE CARE?
Virtually all of the chemistry that makes life possible occurs in an aqueous environment
Medical tests (sugar, cholesterol, and Iron)
In our environment (nitrates, chloroform)
Aqueous solution: a solution in which water (H2O) is the dissolving medium, or solvent
Solvent: the dissolving medium in a solution
Solute: the dissolved substance in a solution
Water in salt water
Water in soda
Salt in salt water
Sugar in soda drinks
Carbon dioxide in soda drinks 3

3. WATER...The Universal Solvent
Most valuable property of water is its ability to dissolve many different substances
Bent shape with a bond angle (H-O-H) of 105º
Polar molecule 4

4. Solvation (hydration) : The process of attraction and association of molecules of a solvent with molecules or individual ions of a solute
When an ionic solid dissolves in water the ions become hydrated and are independently dispersed throughout the solution 5

5. What about Non-ionic substances???
Ethanol (C2H5OH) is very soluble in water
Wine, beer, and mixed drinks are aqueous solutions of ethanol and other substances 6

6. Electrical Conductivity
Electrolyte: a substance that will conduct electricity when in aqueous solution
Strong electrolyte
Weak electrolyte
Non-electrolyte
All depends directly on the number of ions present in the solution (more ions.....stronger electrical current.....bright light....strong electrolyte) 7

7. Strong electrolytes: substances that are completely ionized when they are dissolved in water
Soluble salts
Strong acids
Strong bases 8

8. Weak electrolytes: substances that slightly ionize when dissolved in water (few ions are produced)
Weak acids
Weak bases

9. Nonelectrolytes: substances that dissolve in water but do not ionize (no ions produced)
Ethanol (C2H5OH)
Sugar (C12H22O11) 9

10. Identify each of the following as a strong, weak or nonelectrolyte
Electrolyte Practice
Identify each of the following as a strong, weak or nonelectrolyte
NaCl
NH3
C6H12O6
HCN
KOH
H2SO3 HF
HCl
Li2SO3

11. M = moles of solute/ L of solution
In order to perform stoichiometric calculations we must know 2 things:
1. The form the chemical takes when dissolved
2. Amount of chemical present in the solution (usually expressed in concentration--Molarity)
Molarity (M) = moles of solute per volume of solution (L)
M = moles of solute/ L of solution 10

12. Example 1: Calculate the molarity of a solution prepared by dissolving 11.5 g of solid NaOH in enough water to make 1.50 L of solution 11

13. Example 2: Calculate the molarity of a solution prepared by dissolving 1.56 g of gaseous HCl in enough water to make 26.8 mL of solution 12

14. Example 3: How many grams of Sodium chloride are needed to prepare 1
Example 3: How many grams of Sodium chloride are needed to prepare 1.50 L of M NaCl solution?
Volumetric Flasks 13

15. Example 4: A chemist needs 1. 00 L of an aqueous 0
Example 4: A chemist needs 1.00 L of an aqueous M K2Cr2O7 solution. How much solid potassium dichromate must be weighed out to make this solution? 14

16. Steps Involved In Making a Standard Solution

17. B.) 1.0 M Iron III Perchlorate
Example 5: Determine the concentration of each type of ion in the following solutions
A.) 0.50 M Iron II Nitrate
B.) 1.0 M Iron III Perchlorate 15

18. Example 6: Calculate the number of moles of Cl- ions in 1. 75 L of 1
Example 6: Calculate the number of moles of Cl- ions in 1.75 L of 1.0 x 10-3 M ZnCl2 solution 16

19. Example 7: Typical blood serum is about 0. 14 M NaCl
Example 7: Typical blood serum is about 0.14 M NaCl. What volume of blood contains 1.0 mg NaCl? 17

20. moles of solute after dilution = moles of solute before dilution
Dilution: a specific (calculated) amount of water is added to a stock solution (concentrated solution) to achieve a solution of a desired concentration
Volumetric (transfer) pipet - designed to measure one volume accurately (5 mL, 10 mL, 15 mL)
Measuring pipets - various volumes of liquids
moles of solute after dilution = moles of solute before dilution
MstockVstock = MdiluteVdilute
M1V1 = M2V2 18

21. Example 1: What volume of 17
Example 1: What volume of 17.4 M stock solution of acetic acid must be used to prepare 500. mL of 1.00 M acetic acid?
***HOW TO PREPARE THE SOLUTION:
Using a measuring pipet, deliver 28.7 mL of 17.4 M acetic acid into a 500. mL volumetric flask containing distilled water then dilute it with distilled water to a total volume of 500. mL

22. Example 2: What volume of 16 M Sulfuric acid must be used to prepare 1
Example 2: What volume of 16 M Sulfuric acid must be used to prepare 1.5 L of a 0.10 M H2SO4 solution? How would you prepare this solution? 20

23. 3 Main Types of Solution Reactions:
Precipitation
Acid-Base
Oxidation-Reduction

24. Precipitation Reactions
2 solutions are mixed, an insoluble substance sometimes forms known as a precipitate
Solid that forms and separates from solution
Example: When an aqueous solution of potassium chromate is added to a aqueous solution containing Barium Nitrate, a yellow solid forms.
What is the identity of the yellow solid.
What is the equation that describes the chemical change?

27. Exceptions
Salts containing

28. SOLUBILITY RULES YOU NEED TO KNOW!
Compounds with an alkali metal cation or an ammonium cation (NH4+) are ALWAYS SOLUBLE
Compounds with a nitrate (NO3-) anion are ALWAYS SOLUBLE
Strong Bases -- Alkali metal (Group IA) hydroxides & Ca(OH)2, Sr(OH)2 and Ba(OH)2

29. Describing Reactions in Solution
Formula equation
Gives the overall reaction stoichiometry but not necessarily the actual forms of the reactants and products in solution
2. Complete ionic equation
Represents as ions all reactants and products that are strong electrolytes
3. Net ionic equation
Includes only those solution components undergoing a change (NO spectator ions)
Spectator Ion(s) : Ions that are present in the solution before and after the reaction. They do not participate directly in the reaction

30. Example 1: Aqueous Lead II Nitrate is added to aqueous Potassium Iodide
Formula Equation
Complete Ionic Equation
Net Ionic Equation

31. Example 2: Aqueous Potassium chromate is added to aqueous Barium nitrate
Formula Equation
Complete Ionic Equation
Net Ionic Equation

32. Example 3: Aqueous Potassium chloride is added to aqueous Silver nitrate
Formula Equation
Complete Ionic Equation
Net Ionic Equation

33. Part II. #1-4 do a/b/c……#4-8 ONLY a
TIME TO PRACTICE!
Part II. #1-4 do a/b/c……#4-8 ONLY a

34. Chapter 4.7: Stoichiometry of Precipitation Reactions

35. Stoichiometry of Precipitation Reactions
Always write down the species present in the solution (ions)
To obtain moles of reactants we must use the volume and Molarity (M)
Steps to Solving:
Write the balanced net ionic equation
Determine moles of reactants
Identify limiting reactant (if necessary)
Determine moles of product

36. Example 1: Calculate the mass of solid NaCl that must be added to 1
Example 1: Calculate the mass of solid NaCl that must be added to 1.50 L of a M AgNO3 solution to precipitate all the Ag+ ions in the form of AgCl
1. Write the balanced net ionic reaction
2. Determine moles of reactants
3. Determine moles of product
4. Convert to desired units (mass, volume….)

37. Example 2: When an aqueous solution of Na2SO4 and Pb(NO3)2 are mixed, PbSO4 precipitates. Calculate the mass of PbSO4 formed when 1.25 L of M Pb(NO3)2 and 2.00 L of M Na2SO4 are mixed
1. Write the balanced net ionic equation
2. Calculate the moles of reactants
3. Determine the limiting reactant
4. Calculate the moles of product
5. Convert to grams of product

38. TRY THESE…
What mass of Na2CrO4 is required to precipitate all of the Ag+ ions from 75.0 mL of M solution of AgNO3?
What mass of barium sulfate can be produced when mL of a M solution of barium chloride is mixed with mL of a M solution of Iron III sulfate?

39. LAB NOTEBOOK SET-UP
Table of Contents
Objective
Procedure
Calculations & Data
Analysis Questions (if applicable)
Conclusion