1. Types of Chemical Reactions And Solution Stoichiometry
Chapter 4

2. Section 4.1: Water, The Common Solvent
Hydration of an ionic compound will occur when the partial positive end of a water becomes attracted to the anions in the compound; likewise for the partial negative center of the water and the cations.
Solubility depends on the strength of the intermolecular attractions between the ions and water, as well as the intramolecular attractions of the cations and anions of the compound.
Water has polar, covalent bonds. The oxygen atom is more electronegative, making electrons more attracted to it than to the hydrogen. This will create a dipole on the molecule, with a more positive end and a more negative end.
When an ionic compound dissolves in water, the ions dissociate completely, as shown in the equation.
NH4NO3(s)  NH4+(aq) + NO3-(aq)

3. What can dissolve in H2O? WHY? Insoluble Soluble Fats Alcohols
ex: bacon grease
Oils
ex: cooking oil
Non-Polar Substances
ex: turpentine
Soluble
Alcohols
ex: C2H5OH
Sugars
ex: C6H12O6
Ionic compounds
ex: NaCl, KOH, LiBr
Draw the Lewis structures for alcohol, glucose, and water. Observe how the highly EN O and less EN H are present in each of these structures (hydrogen bonding).
Like Dissolves Like speaks only about polarity of a molecule. Polar things are attracted to other polar molecules; non polar molecules are attracted to other non-polar molecules. It has nothing to do with shape.
WHY?
Because of intermolecular forces: the OH group on the sugars and alcohols is particularly attractive to a water molecule.
Generally speaking: “Like Dissolves Like”

4. Section 4.2: Strong and Weak Electrolytes
Solute + Solvent = Solution
Strong electrolytes conduct electricity
Weak electrolytes barely conduct electricity
Conductivity depends upon ionization
Solute is what is being dissolved (not necessarily solid).
Solvent is doing the dissolving (typically more of this present).
Solution is homogeneous mixture.
More ions present, more conductivity because the ions are the substances capable of carrying a charge and complete the circuit.

5. All of these dissociate completely in water. Weak Electrolytes
Strong Electrolytes
Soluble salts
Strong acids
Strong bases
All of these dissociate completely in water.
Weak Electrolytes
Weak acids
Weak bases
All of these partially dissociate in water
HCl  H+ + Cl-
NaOH  Na+ + OH-
HC2H3O H+ + C2H3O2
Weak electrolytes are represented at equilibrium because the ions combine to make the molecule again. In the case of ammonia, the ionization occurs very slowly since it is not favored to occur.
Non-Electrolytes are completely molecular substances in water (not even a little dissociation); Non polar substances.

6. Section 4.3: Composition of Solutions
Concentration is measured in molarity, molality, and many others.
Concentration DOES NOT directly express the number of ions present in a solution.
M= moles solute
liters solution
MgCl2  Mg Cl-
1.0 M M M

7. Sample Problems
Calculate the number of moles of Cl- ions in 1.75 L of 1 x 10-3 M ZnCl2.
A chemist needs 1.0 L of 0.20 M K2Cr2O7 solution. How much solid K2Cr2O7 must be weighed out to make this solution?
DUH! This is chem 1 material.

8. Standard Solution: a solution whose concentration is accurately known.
Example: M HCl; M NaOH
Creating dilutions
Chemical analysis of a compound
Theoretical Calculations
What would you do to prepare a standard solution? In your answer, include specific pieces of glassware, techniques, or equipment you should use.
ANSWER
NOW

9. moles before dilution = moles after dilution
Dilutions
Dilution is the process used to make the solution less concentrated.
moles before dilution = moles after dilution
Because M =mol/L,
V1(M1) = V2(M2)
Lab Technique: Use a pipet to deliver the correct amount of original solution to a volumetric flask. Add some water, swirl. Fill to line, invert.

10. You have a large quantity of 1. 5 M NaOH solution available
You have a large quantity of 1.5 M NaOH solution available. Dilute this to 100.0mL of a 0.05 M solution. Submit your calculations and store your final product for use in our first lab. DO
NOW

11. Section 4.4: Types of Chemical Reactions
There are more than just these few types, but in this chapter we will cover…
Precipitation
Acid-base
Oxidation-Reduction

12. Section 4.5: Precipitation Reactions
Precipitation Reactions (double displacement)
Forms a solid precipitate from aqueous reactants.
Color of precipitate can help in identification
Solubility rules help BUNCHES
MORE…

13. Solubility RULES
All compounds containing alkali metal cations and the ammonium ion are soluble.
All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are soluble.
All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+, and Hg2+.
All sulfates are soluble except those containing Hg2+, Pb2+, Sr2+, Ca2+, and Ba2+.
All hydroxides are only slightly soluble, except those containing an alkali metal, Ca2+, Ba2+,and Sr2+. NaOH and KOH are the most soluble hydroxides.
All compounds containing PO43-, S2-, CO32-, and SO32- are only slightly soluble except for those containing alkali metals or the ammonium ion.

14. Practice Predicting Potassium nitrate and barium chloride
Sodium sulfate and lead (II) nitrate
Potassium hydroxide and iron (III) nitrate

15. ALL REACTIONS SHOULD BE WRITTEN IN NET IONIC FORM

16. Section 4.7: Stoichiometry of Precipitation Reactions
Stoichiometry in a precipitation reaction is performed just like stoichiometry for a molecular reaction.
You need to know which ion comes from which molecular formula.

17. Sample problem
Calculate the mass of solid NaCl needed to add to 1.5 L of 0.1 M silver nitrate solution to precipitate all Ag+ ions in the form of AgCl.
Net Ionic Eq: Ag+ + Cl-  AgCl

18. General Format Write the Net Ionic Equation
Calculate the moles present
Identify the Limiting Reactant*
Use Mole Ratio(s)
Fancy-fy your answer (put in correct units)

19. Try Me!
What mass of precipitate will be produced when 50.0 mL of 0.200M aluminum nitrate is added to mL of M potassium hydroxide?

20. Section 4.8: Acid-Base Reactions
Acids yield H+
Bases yield OH -
Definitions of acid and base vary.
Arrhenius and Bronsted/Lowry are common theories.
Acid-Base rxns are called NEUTRALIZATIONS
Bases are proton acceptors
Acids
are proton
donors

21. Strong Acid-Strong Base (HCl) (NaOH)
Both dissociate completely
H+ + OH-  H2O
Na+ and Cl- are spectators.
Weak Acid - Strong Base
(HC2H3O2) (KOH)
Acetic acid will not dissociate
KOH will completely
HC2H3O2 + OH-  H2O + C2H3O2-
K+ is a spectator.

22. Stoichiometry sample
What volume of M HCl is needed to neutralize 25 mL of 0.35 M NaOH?
H+ + OH-  H2O

23. Titrations…define me! Volumetric analysis Titration Titrant Analyte
Equivalence point
Indicator
Endpoint

24. To complete a successful titration…
The reaction between the titrant and the analyte should be known (you should know WHAT substances you have)
The equivalence point should be marked accurately (you should use the right indicator)
Volume of the titrant needed to reach the equivalence point should be recorded accurately (you should use a buret!)

25. Effective Indicator Ranges

26. Titration Try Me Calc 1
A 50.0 mL sample of a sodium hydroxide solution is to be standardized M of KHP (potassium hydrogen phthalate, KHC8H4O4) is used as the titrant. KHP has one acidic hydrogen mL of the KHP solution is used to titrate the sodium hydroxide solution to the endpoint. What is the resulting concentration of the analyte?

27. Titration Try Me Calc 2
How many milliliters of a M sodium hydroxide solution are needed to neutralize 20.0 mL of a M sulfuric acid solution?

28. Norton Tutorial
Go to the website
Find the tutorial on Acid/Base ionization.
Complete the tutorial question form.

29. Section 4.9: Redox Reactions
What is it??
-A reaction that occurs in conjunction with a transfer of electrons.
We assign oxidation states to individual atoms in a reaction to observe the change in electrons.
Oxidation states
are written with
the +/- sign
before the quantity.
Ion charges are
written with the
+/- sign behind
the quantity.

30. Assigning Oxidation States
The Oxidation State of…
Quantity of Oxid. State
Examples
An atom in element form
Zero
Na(s), O2(g)
A monatomic ion
Equal to the charge on the ion
Na+, Cl-
Fluorine in a compound
-1 , always
HF, PF3
Oxygen in a compound
-2, except in peroxide where it is -1
H2O, CO2, H2O2
Hydrogen in a compound
+1, always
H2O, HCl, NH3

31. Oxidation= an increase in the oxidation state
Reduction = a decrease in the oxidation state
oxidation
2Na(s) + Cl2(g)  2NaCl(s)
reduction

32. The metal is oxidized and the other substance is reduced.
Metal Atom
Oxidized Substance:
Loss of electrons
Oxidation state increases
Gets Smaller
Called the Reducing Agent
Other Atom e-
Other Ion
Metal
Ion
Reduced Substance:
Gain of electrons
Oxidation state decreases
Gets Bigger
Called the Oxidizing Agent
The metal is oxidized and the other substance is reduced.

33. Section 4.10: Balancing Redox
How To, in Acid:
Write the ½ reactions
Balance the non-H and non-O atoms
Balance O by adding H2O where needed
Balance H by adding H+ where needed
Balance charge using e-
Multiply by coefficients until both e- are equal for each ½ reaction
Add the ½ reactions together (cancel stuff)

34. Redox Sample Problem Balance: MnO4- + Fe2+  Fe3+ + Mn2+ Check+2 3+ +3
Check
Charges! +7 +2 8+ 1- 2+
x5!
x5!

35. Redox Try Me Problem
As2O3(s) + NO3-  H3AsO4 + NO(g)

36. How To, in Base:
Repeat steps from old procedure
Cancel out H+ by adding OH- ions
Re-write H+ and OH- as water
Add ½ reactions together (and cancel stuff)

37. Redox Try Me Problem 2
Balance, in base:
Ag(s) + CN- + O2  Ag(CN)2-