1. Chapter 4 Reactions in Aqueous Solutions
Many chemical and almost all biological reactions occur in the aqueous medium
Substances (solutes) that dissolve in water (solvent) can be divided into two categories:
Electrolytes
Non-Electrolytes

2. Three Major Types of Reactions
Precipitation Reaction – the product an insoluble substance separates from the solution
Acid/Base Reactions – A proton transfer from an acid to a base
Oxidation/Reduction (Redox) “the bane of the AP Test” – Electrons are transferred from a reducing agent to an oxidizing agent

3. Solution Stoichiometry
Quantitative studies with known concentrations (Molarity) of solutions
Gravimetric Analysis
Titrations

4. General Properties of a Solution
Solution – a homogenous mixture of two or more substances
Solution may be gaseous (air), solid (alloy) or Liquid (salt water)
In this chapter we will deal only with aqueous solutions
Most common Solvent - Water

5. Electrolytes versus Nonelectrolytes
Electrolytes – Ionic compounds that completely or partially dissociate in solution with the ability to pass electric current in solution
Acids/Bases will ionize in solution, therefore electricity can be conducted
Non-Electrolytes – Molecular compounds that do not dissociate in solution, therefore no electric current can be pass

6. Ionic Compounds in Solution
Water is a great solvent for ionic compounds because it is polar, the positive end attracts the Negative Ion and vice versa

7. Acids and Bases as Electrolytes
Some acid/bases competely dissociate in solution
HCl
HNO3
H2SO4
Ba(OH)2
NaOH
While others only partially dissociate
CH3COOH HF
HNO2
HN3

8. Writing Partial Dissociation Equations
Partial dissociation equations are written with a double arrow, indicating a reversible reaction
Write partial dissociation for CH3COOH

9. Precipitation Reactions
A double replacement reaction (metathesis) in which a product is insoluble

10. Solubility Rules In water at 25 Degrees
All common compounds of Group I and ammonium ions are soluble.
All nitrates, acetates, and chlorates are soluble.
All binary compounds of the halogens (other than F) with metals are soluble, except those of Ag, Hg(I), and Pb.
All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury (I). The latter three are slightly soluble.
Except for rule 1, carbonates, hydroxides, oxides, silicates, and phosphates are insoluble.
Sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium, and ammonium.

11. Soluble or Insoluble at 25 Degrees Celsius in Water
PbSO4
BaCO3
Li3PO4
FeS
Ca(OH)2
Co(NO3)3

12. Net Ionic Equations
Write the correctly balanced equation and decide on state of each product
Write free state of all ions and insoluble product
Cancel out spectator ions – anyone not part of the reaction
Check charges and balancing in net ionic

13. Practice Net Ionic Predict, Balance and write net ionic
Lead Nitrate and Potassium Iodide
Barium Chloride and Sodium Sulfate
Potassium Phosphate and Calcium Nitrate
Aluminum Nitrate and Sodium Hydroxide

14. Acid – Base Reactions
Acids react with metal such as Zn, Mg and Fe to produce hydrogen gas
Acids react with carbonates and bicarbonates to produce carbon dioxide gas, water and the salt

15. Bronsted Acid and Bases
Bronsted Acid is a proton donor
Bronsted Base is a proton acceptor
HCl (aq) H+ (aq) + Cl-(aq)
In water the H+ attracts to the water molecule producing the hydronium ion

16. Monoprotic Acids
Each unit of acid yields one hydrogen ion upon ionization

17. Diprotic Acids
Each unit of the acid gives up two hydrogen ions in two separate steps (they strip)

18. Triprotic Acids
Yield three hydrogen ions in three separate steps (they strip)

19. Bronsted Acid is a proton donor Bronsted Base is a proton acceptor
Classify each of the following as an Bronsted acid or Bronsted base, explain your reasoning based on the definition
HBr
SO-24 HI
HCO-3
NO2

20. Neutralization Reaction Acid and Base will form Salt and Water
Write the net ionic for the following
Hydrochloric acid and Sodium Hydroxide
Sulfuric acid and Aluminum Hydroxide

21. Acid – Base Reactions Leading to Formation of a Gas
Certain Salts – Carbonates, bicarbonates, sulfites and sulfides react with acids to form gaseous products

22. Oxidation Numbers
Oxidation Reaction – refers to half reaction that involves loss of electrons
Reduction reaction – refers to a half reaction that involves the gain of electrons
The extent of oxidation in a redox reaction must be equal to the extent of reduction; that is the number of electrons lost by a reducing agent must be equal to the number of electrons gained by an oxidizing agent

23. The half-reactions of a redox reaction or oxidation-reduction reaction

24. Oxidation Number
The number of charges the atom would have in a molecule if electrons are transfer completely

25. The convention is that the cation is written first in a formula, followed by the anion.
For example, in NaH, the H is H-; in HCl, the H is H+.
The oxidation number of a free element is always 0.
The atoms in He and N2, for example, have oxidation numbers of 0.
The oxidation number of a monatomic ion equals the charge of the ion.
For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3.
The usual oxidation number of hydrogen is +1.
The oxidation number of hydrogen is -1 in compounds containing elements that are less electronegative than hydrogen, as in CaH2.
The oxidation number of oxygen in compounds is usually -2.
Exceptions include OF2, since F is more electronegative than O, and BaO2, due to the structure of the peroxide ion, which is [O-O]2-.
The oxidation number of a Group IA element in a compound is +1.
The oxidation number of a Group IIA element in a compound is +2.
The oxidation number of a Group VIIA element in a compound is -1, except when that element is combined with one having a higher electronegativity.
The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl.
The sum of the oxidation numbers of all of the atoms in a neutral compound is 0.
The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion.
For example, the sum of the oxidation numbers for SO42- is -2.

26. Assign oxidation numbers to all the elements in the following compounds
Na2O
HNO2
Cr2O7-2
PF3
MnO4-

27. Arrange the following species in order of increasing oxidation number of the sulfur atoms
H2S
SO2
SO3 S8
H2SO4
S-2
HS-

28. Concentration Molarity = moles of solute/liters of solution
Example: How many grams of potassium dichromate are required to prepare a 125ml solution whose concentration is 1.83M

29. Concentration
In a biochemical assay a chemist needs a to add 4.07g of glucose to a reaction mixture. Calculate the volume in milliliters the volume of a 3.16M glucose she should use

30. Dilution of Solutions
The procedure of making a less concentrated solution from a high concentration solution
MiVi = MfVf

31. Dilution Problem
Describe hou you would prepare 2.50 * 102 ml of a 2.25M H2SO4 solution, starting with a 7.41 M stock solution of H2SO4

32. Dilution Problem #2
How would you prepare a 200ml of a .866M KOH solution, starting with 5.07M stock solution

33. Acid – Base Titrations
In a titration a solution of an accurately known concentration, called the standard is added gradually to another solution of unknown until reaction is neutralized (equivalence point)
Indicators are used to color the reaction when it is complete

34. Titration Problem
In a titration experiment, a student finds that 25.46ml of a NaOH solution is needed to neutralize .6092g of KHP. What is the concentration of the NaOH solution?

35. Titration Problem #2
How many milliliters of a .836M NaOH solution is needed to neutralized 25ml of a .335M of H2SO4?

36. Solution Stoichiometry
When sodium chloride reacts with silver nitrate, silver chloride precipitates. What mass of silver chloride is produced from 150ml 3M of silver nitrate?

37. When Magnesium chloride reacts with silver nitrate, silver chloride precipitates. What mass of silver chloride is produced from 4.5M in 250ml of silver nitrate? What is the name of the other product of the reaction?