1. Types of Reactions and Solution Stoichiometry
Chapter 4
Types of Reactions and
Solution Stoichiometry

2. Objectives Describe why water is a good solvent
Describe the process dissolving
Compare electrolytes
Calculate molarity
Perform dilution problems

3. Aqueous Solutions
Aqueous Solution – Homogenous mixture where water is the solvent
Solvent – Substance doing the dissolving
Solute – Substance that is dissolved
Water is a very common solvent

4. Water As A Solvent Structure of water allows it to be a solvent
Bent geometry
Oxygen has a high electronegativity
Hydrogen has a low electronegativity
Water is polar
Positive and negative ends
Positive is attracted to negative
Negative is attracted to positive

6. Dissolving Process is called hydration “Likes Dissolve Likes”
Water molecules surrounding particles
“Likes Dissolve Likes”
Polar solvents dissolve polar solutes
Water, Alcohols, Ionic, Sugars
Nonpolar solvents dissolve nonpolar solutes
Hydrocarbons, Petroleum, Fats

7. Electrolytes Substance that when dissolved conducts electrical current
Means there are ions in solution
Strong electrolytes – Conduct current well
Strong Acids/Bases and Salts
Weak electrolytes – Do not conduct well
Weak Acids/Bases
Nonelectrolytes – Does not conduct
Covalent compounds

8. Molarity Moles of solute per liter of solution
Molarity = moles / liters
Units mol/L
3.0 mol/L or 3.0 M
Also called Molar

9. Example
What is the molarity of a solution made by dissolving 8.1 grams of NaCl in 500. mL of solution?

10. Dissolving Reactions
When ionic compounds dissolve they dissolve into ions based on their formulas
HCl  H+ + Cl-
NaNO3  Na+ + NO3-
Li2SO4  2Li+ + SO4-2
Mg(NO3)2  Mg+2 + 2NO3-

11. Ion Concentration
The concentration of ions in solution is not always the same as the concentration of the solution
Ca(NO3)2  Ca+2 + 2NO3-
If Calcium Nitrate is 2.0 M then
Calcium ion is 2.0M
Nitrate ion is 4.0M

12. Example
How many moles of Nitrate ions are present in 500. mL of 2.0M Calcium Nitrate?

13. Preparing Solutions
A way of determining what mass of a solute is required to prepare a solution.
What volume x molarity x molar mass

14. Example
What mass of sodium chloride is required to prepare 250. mL of a .150M solution?

15. Preparing the solution

16. Dilution Process of making a solution less concentrated Equation
C1V1 = C2V2
C = Concentration V = Volume
Volume can be in L or mL (Just be the same)
Each side is equal to moles

17. Example
What volume of 15.4 M Nitric Acid is required to prepare 750. mL of 1.0M HNO3?

18. Homework
p. 181 #'s 15,18a,20,21,24abe,26,27

19. Objectives
Describe the three types of chemical reactions

20. Reaction Types Precipitation Neutralization
Formation of an insoluble compound
Neutralization
Acid and base form a salt and water
Oxidation Reduction Reaction (Redox)
Reaction involving the transfer of electrons

21. Precipitation Reactions
An insoluble compound (and a soluble compound) is formed from two compounds that are usually in solution.
Sometimes called double replacement reactions
You need to determine the insoluble compound by solubility rules

22. Solubility Rules Table 4.1 in text Absolutely must be memorized
Nitrate, Acetate, Ammonium and Group 1 salts are soluble
Most halide salts are soluble. Except for Ag+,Pb+2, and Hg2+2
Most sulfates salts are soluble. Except for Ba+2,Pb+2,Hg2+2,Ca+2
All Acids are soluble

23. Cont.
Most hydroxide salts are insoluble. Ca(OH)2, Ba(OH)2, Mg(OH)2 are slightly soluble
Most sulfides, carbonates, chromates, phosphates are insoluble
Memorize, memorize, memorize!!!!

24. Test Yo’ Knowledge Calcium Nitrate Ammonium Phosphate
Soluble
Ammonium Phosphate
Lead (II) Chloride
Insoluble
Iron (III) Sulfate

25. Reaction Prediction Predict the products of the following reactions
LiNO3 (aq) + BaI2 (aq) 
Cu(C2H3O2) (aq) + NaOH (aq) 
Ba(NO3)2 (aq) + Na2SO4 (aq) 

26. Types of Equations Molecular Equations Complete Ionic Equations
Normal equations including all formulas
Complete Ionic Equations
Includes ions of all dissociated compounds
Net Ionic Equation
Removes spectator ions
Only includes species that undergo change
Very handy equations

27. Dissociation Compounds dissolving into ions HCl  H+ + Cl-
NaNO3  Na+ + NO3-
Li2SO4  2Li+ + SO4-2
Mg(NO3)2  Mg+2 + 2NO3-

28. Net Ionic Equations Dissociate all compounds that are soluble
Leave all solids, liquids, and gases together
Remove species that are exactly the same on both sides
Rewrite equation with remaining species

29. Example
Predict the products when solutions of potassium iodide and lead (II) nitrate are mixed. Then write the net ionic equation.

30. On a molecular level draw the contents of the beakers containing the potassium iodide, lead (II) nitrate, and resulting products.
KI(aq)
Pb(NO3)2 (aq)
Products

31. Solution Stoichiometry
Same as other types of stoichiometry except molarity is used in calculations
General method
Volume x molarity x mole ratio x molarity or
molar mass
Limiting reactants still apply

32. Example
What volume of .15M KI is required to react with 25mL of .10M Pb(NO3)2?

33. Example
What mass of AgCl is produced when 47mL of 0.33M AgNO3 reacts with 86mL of 0.48M NaCl?

34. Homework
p. 182 #'s 30,32,34,38,39,42

35. Acid Base Reactions
The reaction of an acid and base to produce a salt and water
Acid – Proton (H+) donor
Acids formulas start with H
Base – Proton acceptor
Usually have OH- at the end
Also ammonia (NH3)

36. Example
Solutions of hydrobromic acid and sodium hydroxide are mixed. Write the molecular equation. Write the net ionic equation
The net ionic equation for ALL neutralization reactions is always
H+ + OH-  H2O

37. Example
What volume of .100M HCl will react with mL of .200M NaOH?

38. Titrations Volumetric Analysis – Procedure with liq.
Titrant – Solution of known concentration
Analyte – Chemical under study
Equivalence Point – Mole ratio equality
End Point – Visual change point in a reaction
Indicator – Chemical that marks end point
End point and equivalence point should agree

39. Cont. Required data Volumes or masses of titrant and analyte
Concentrations if known
Types of calculations
Molarity
Molar Mass
Mass Percent

40. Example
38.20 mL of KOH was titrated with mL of .75 M standard HNO3. What was the concentration of the KOH?

41. Homework
p. 183 #'s 48,50c,52,54

42. Objectives Describe redox reactions Calculation oxidation states
Balance redox reactions

43. Oxidation Reduction Reactions
A reaction that involves a transfer of electrons
Also called redox reactions
Examples
Combustion
Respiration
Single replacement

44. Oxidation Numbers
A number assigned to an atom that assumes the complete transfer of electrons in compounds
A way of keeping track of electrons
Evidence for redox reactions
If oxidation numbers change then it is a redox reaction.

45. Assigning Oxi. Numbers Any free element = 0
Examples, Al, P, O2, H2
Any monatomic ion = its charge
Example Na+ = +1
Oxygen in compounds = -2
Except for peroxides then = -1
Hydrogen in compounds = +1
Fluorine in compounds = -1

46. Cont. The sum of all oxidation numbers in a compound = 0
The sum of all oxidation numbers in an ion equals the charge of the ion
Find the oxidation number for all atoms in the compound Ca(ClO3)2
Ca +2, O -2, Cl +5

47. Examples – Find the oxidation number for each atom
NaNO3
Na +1, N +5, O -2
Li3PO4
Li +1, P +5, O -2
Cr2O7-2
Cr +6, O -2

48. Redox Vocabulary Reduced – Gain electrons Oxidized – Lost electrons
Leo says Ger
Lose electrons Oxidized, Gain electrons reduced
Oxidizing agent – chemical that is reduced
Reducing agent- chemical that is oxidized
Electrons lost equals electrons gained

49. Balancing Redox Equations
Separate the reaction in the oxidation and reduction half reactions
Balance each half reaction
Elements
Balance O with Water
Balance H with H+
Charge with electrons
Equalize electrons
Add reactions/Cancel like terms

50. Cont. If the solution is basic there cant be H+ left
Add the same number of OH- to both sides as there are H+ left
H+ and OH- make water

51. Example
Balance the equation and determine what species is reduced and oxidized.
Au+3 + I-  Au + I2

52. Example
Balance the equation and determine what species is the reducing agent. (Acidic)
Cr2O7-2+ I-  Cr+3 + I2

53. Cr(OH)3+ ClO3-  CrO4-2 + Cl-
Example
Balance the equation and determine what species is reduced and oxidized. (Basic)
Cr(OH)3+ ClO3-  CrO4-2 + Cl-

54. Things To Look For
Make sure the number of electrons are equal on both sides of the equation
Check whether is acidic or basic
Sometimes an element can be oxidized and reduced
Disproportionation

55. Homework
p. 183 #’s 58a-d,62,64ab,65ab