1. Aqueous Reactions and Solution Stoichiometry Pg 105
Chapter 4
Aqueous Reactions and Solution Stoichiometry
Pg 105

2. Aqueous Solutions
-Aqueous Solutions are solutions that have water as the dissolving medium.
-Many reactions contain substances that have been dissolved in water, making them aqueous solutions.
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3. 3 Main Major Chemical Reaction Types Involving Aqueous Solutions
Precipitation Reactions
Acid-Base Reactions
Redox Reactions

4. 4.1 General Properties of Aqueous Solutions
Solutions are homogeneous mixtures
Usually has more solvent than solute.
Solute is the substance being dissolved in the solvent

5. Electrolytic Properties
Pure water is a bad conductor
The presence of Ions in water makes it into a good conductor
Aqueous solution that conduct electricity such as NaCl(aq) or other ionic compounds are electrolyte.
Solutions that do not form ions like sucrose and other molecular compounds are nonelectrolytes.

6. Ionic Compounds in Water
Ionic compounds dissolve in water dissociating into component ions (ex. NaCl -> Na+&Cl-)
The polar nature of water makes it a very effective solvent
The polarity helps prevent anions and cations from rejoining.

7. Molecular Compounds in Water
Structure usually remains unchanged, they usually do not form ions
Acids and a few other compounds like ammonia react with water forming ions making an electrolyte.
Ex. HCl make H+ and Cl- ions

8. Strong & Weak Electrolytes
Strong Electrolytes = Most ionic compounds and a few molecular compounds.
Weak Electrolytes = Molecular compounds that produce few ions when dissolved
If the chemical reaction goes both ways, breaking into ions, and recombining, than the substance is a weak electrolyte.

9. HCl(aq) --> H+ + Cl- One arrow means strong electrolyte
HC2H3O29(aq) <--> H+ + C2H3O2-
Double arrow means weak electrolyte

10. 4.2 Precipitation Reactions
Precipitation Reaction = Reactions that result in the formation of an insoluble product.
Precipitate = Insoluble solid formed by a reaction in a solution

11. Solubility Guidelines for Ionic Compounds
Solubility = Amount of substance that can be dissolved in given amount of solvent
If less than .01 mol dissolves in a liter, substance is insoluble. In these substances intermolecular attraction is stronger than the waters polarity.
Table 4.1 pg 111 (Solubility Guidelines for Common Ionic Compounds in Water)
All ionic compounds with 1A elements or ammonia ions are soluble in water.

12. Is Sodium Carbonate Soluble (Na2CO3)
Yes. Carbonate is usually insoluble, but when paired with a 1A element, Sodium, the compound becomes soluble.

13. Exchange (Metathesis) Reactions
Exchange or Metathesis Reaction = AX+BY --> AY+BX
Precipitation and Acid Base Reactions conform to this pattern

14. What precipitate forms when BaCl2 and K2SO4 are mixed?
BaSO4, SO42- is soluble but Ba2+ is not

15. Ionic Equations
Molecular Equation = complete chemical formulas of reactants and products
Complete Ionic Equation = All Soluble strong Electrolytes are shown as ions
Spectator ions = ions that are present in the same form on both product and reactant side. These are dropped out to form a Net Ionic Equation.

16. Steps to Write a Net Ionic Equation
Write a balanced Molecular Equation
Rewrite to show ions that are formed during dissociation or ionization, only the strong electrolytes are written in ionic form
Cancel spectator ions on both sides

17. Write the net ionic equation for the mixing of CaCl2 and Na2CO3
CaCl2(aq) + Na2CO3(aq) --> CaCO3(S) + 2NaCl(aq)
Ca2++ 2Cl- + 2Na+ + CO32--->CaCO3(s) + 2Na+ + 2Cl-
Ca2+(aq) + CO32-(aq)-->CaCO3(s)

18. Acid-Base Reactions Acids and Bases are common Electrolytes
Are some of the most common compounds we encounter

19. Acids
Substances that ionize in aqueous solutions upping H+ concentration
Protic refers to amount of H+ ions ionizing. Monoprotic = 1, Diprotic = 2.
Diprotic Acid ionization occurs in two steps, One hydrogen is separated at a time.

20. Bases Substances that accept H+ ions, or increases OH- concentration.
Does not need to have an OH- ion, if accepts H+ like NH3 (ammonia is a weak electrolyte)

21. Strong and Weak Acids and Bases
Strong Acids and Bases are strong electrolytes that completely ionize in solutions
Weak Acids and Bases are electrolytes that partly ionize in solutions
Table 4.2 pg 115 (Common Strong Acids and Bases)

22. Identifying Strong and Weak Electrolytes
Is the compound ionic, yes -> probably strong electrolyte
Not ionic, is it an acid
Yes, is an acid, if strong, is a strong electrolyte if weak, is a weak electrolyte.
Not an acid, is it NH3 or another molecular base, yes -> weak base, no -> probably nonelectrolyte

23. Classify HNO3 as a strong, weak, or non Electrolyte
HNO3 is a strong acid making it a strong electrolyte.
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24. Neutralization Reactions and Salts
When acid and base react together, it is a neutralization reaction.
These reactions form a salt and a water.
Salt = any ionic compound whose cation comes form a base and anion comes from an acid.

25. Write the net ionic equation for HC2H3O2 and Ba(OH)2
HC2H3O2(aq)+OH-(aq)--> H2O(l) + C2H3O2-(aq)

26. Acid-Base Reaction with Gas Formation
The sulfide ion and carbonate ion react with acids to form gases
2HCl(aq) + Na2S(aq) --> H2S(g)+2NaCl(aq)
HCl(aq)+NaHCO3(aq)-->NaCl(aq) + H2O(l)+CO2(g)

27. 4.4 Oxidation - Reduction Reactions

28. Oxidation and Reduction
Metals undergoing erosion are losing electrons and forming cations
Loss of electrons is known as oxidization
The gain of electrons by a substance is called reduction

29. Oxidation Numbers
Oxidization number is the actual charge of the of the atom.
In elemental form the Oxidization number is 0
Oxidization of monatomic ions equals the charge
Nonmetals are usually negative, Oxygen usually is -2, Hydrogen is +1, Fluorine is -1.
Sum of oxidation numbers in neutral compound is 0 or equal to the charge in a polyatomic ion.

30. Determine the oxidation stat of sulfur in H2S
-2, Hydrogen is always +1 2H = +2 so S = -2 so that sum of oxidation numbers = 0

31. Oxidation of Metals by Acids and Salts
Displacement Reactions = ion is solution is displaced or replaced through the oxidation of an element.
A+BX-->AX+B

32. Write the net ionic equation for the reaction of aluminum and hydrobromic acid
2Al(S) + 6H+(aq)+6Br--->2Al3+(aq)+6Br-(aq)+3H2(g)
2Al(s)+6H+(aq)-->2Al3++3H2(g)

33. The Activity Series
Table 4.5 pg 124 Activity Series of Metals in Aqueous Solution
Is a table of metals arranged in order of decreasing ease of oxidation.
Alkali and Alkaline Earth Metals are at the top.
Metals on the list can be oxidized by any of the ions below them.

34. Can Pb(NO3)2 can oxidize Zn, Cu, or Fe
Zn and Fe
Refer to table 4.5

35. 4.5 Concentrations Of Solutions
Concentration = Amount of solute dissolved in a given quantity of solvent or solution.

36. Molarity
Molarity (M) = (moles of solute)/(Volume of Solution in Liters)

37. Calculate the molarity of a solution made by dissolving 23
Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form 125ml of solution
23.4g*(1mol Na2SO4/142g Na2SO4) = .165 mols
125ml*(1L/1000ml)=.125
.165mols Na2SO4/.125L = 1.32M

38. Dilution Adding water to lower the concentration is called dilution
Mi*Vi=Mf*Vf
i = initial f = final M = Molarity V= Volume

39. How many mL of 3.0 M H2SO4 are required to make 450 mL of .10 M H2SO4 ?
Vi = (MfVf)/Mi
((.10M)(450mL))/3.0M = 15 mL

40. 4.6 Stoichiometry and Chemical Analysis

41. Titrations
Second solution of known concentration is called the standard solution
Combining the standard solution with a solution of unknown concentration to get a chemical reaction is called titration
Equivalence point is where equivalent quantities have been brought together indicators change the color helping us to find this point.
If molar ratio is 1 to 1 you may use the dilution equation.
If not, convert standard solution to mols, then use molar ratio to give you the mols of the unknown, then convert to grams.

42. How many grams of chloride ion are in the sample of the water if 20
How many grams of chloride ion are in the sample of the water if 20.2 mL of .1 M Ag is required to react with all the chloride in the sample?
(20.2 mL solution) * (1L/1000mL solution) * (.1mol Ag+/L solution) = 2.02 * 10-3
1mol Ag+ : 1mol Cl-
(2.02*10-3 mol Cl-) * (35.5g Cl-/1 mol Cl-) = 7.17 * 10-2 g Cl-