1. Chapter 4: Chemical Reactions
Chemistry 1061: Principles of Chemistry I
Andy Aspaas, Instructor

2. Chemical reactions Ions in aqueous solution
Molecular and ionic equations
Types of reactions
Precipitation reactions
Acid-base reactions
Oxidation-reduction reactions
Solutions
Concentration and dilutions
Quantitative analysis
Gravimetric and volumetric analyses

3. Ions in aqueous solutions
Ionic theory of solutions: Arrhenius, 1884
When dissolved in water, the individual ions of ionic substances completely separate and enable the solution to conduct electricity
Pure water is a poor conductor of electricity
Electrolyte: substance that dissolves in water to give an electrically conducting solution
Generally, ionic solids that dissolve in water are electrolytes
A few molecular electrolytes, Ex. HCl (g)
Nonelectrolytes: dissolve in water, poorly conducting solution, usually neutral molecular substances

4. Strong and weak electrolytes
The extent to which a solution conducts electricity indicates the “strength” of the dissolved electrolyte
Strong electrolytes: exist in solution almost entirely as ions
Ex. NaCl
Weak electrolytes: dissolve in water to give only a small percentage of dissociated ions
Ex. NH3

5. Solubility rules
Solubility: ability of a substance to dissolve completely in water
Ex. Sugar, NaCl, ethyl alcohol are soluble
Ex. Calcium carbonate, benzene are insoluble
Soluble ionic compounds are strong electrolytes
8 solubility rules can determine whether an ionic compound is soluble or not

6. Solubility rules
Li+, Na+, K+, NH4+

7. Molecular and ionic equations
Molecular equation: chemical equation in which reactants and products are written as if they were molecular substances, even if they exist as ions in solution
Explicit in the actual compounds added to a solution, and the products obtained
Complete ionic equation: all strong electrolytes are written as their dissociated ions (aq)
Insoluble compounds are written as a solid compound, not ions

8. Net ionic equations
Spectator ion: ion in an ionic equation that does not take part in the reaction
Appears in ionic form on both sides of a reaction
Net ionic equation: equation in which all spectator ions have been canceled
Several different reactions can have the same net ionic equation

9. Precipitation reactions
Precipitate: insoluble compound formed during a chemical reaction in solution
Predicting precipitation reactions:
Exchange reaction most common, each compound “trades partners” to form products
Write molecular equation
Use solubility rules to determine phase lables for each product and reactant; (aq) if soluble, (s) if insoluble
If all components of reaction are soluble, no reaction occurs
If a product is insoluble, it forms as a precipitate
A net ionic equation shows the reaction at the ionic level

10. Acid-base reactions
Acids: vinegar (acetic acid), lemon juice (citric acid), Coca-Cola (phosphoric acid and carbonic acid), battery acid (sulfuric acid)
Bases: Drano (sodium hydroxide), ammonia, Milk of Magnesia (magnesium hydroxide)
Brønsted-Lowry acid: molecule or ion that donates a proton to another species in a proton-transfer reaction
Brønsted-Lowry base: molecule or ion that accepts a proton in a proton transfer reaction

11. Strong acids and strong bases
Strong acids and bases ionize completely in water
Strong acids: HClO4, H2SO4, HI, HBr, HCl, HNO3
Strong bases: LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
Weak acids and bases only partly ionize in water

12. Neutralization reactions
Reaction between acid and base to produce a salt and possibly water
Salt: ionic compound formed in neutralization reaction
Start by writing molecular equation
Acid anion and base cation form the salt
Water is usually a product
Net ionic equation: write any strong acid or base as its dissociated ions

13. Acid-base reactions with gas formation
Carbonates (CO32-) form H2O and CO2 when reacted with acids
Sufites (SO32-) form H2O and SO2 when reacted with acids
Sulfides (S2-) form H2S when reacted with acids

14. Oxidation-reduction reactions
Oxidation-reduction reactions (redox) involve transfer of electrons
Oxidation number: actual charge of an atom if it exists as a monatomic ion, or a hypothetical charge assigned by a few rules
Elemental atoms always have ox. # 0
Oxygen is usually -2
Hydrogen is usually +1
Halogens usually -1 (unless bonded to another halogen or oxygen)
Sum of ox. #’s of atoms in a compound is 0, sum of ox #’s in a polyatomic ion is the charge on the ion

15. Describing oxidation-reduction reactions
If a species loses electrons, it is oxidized
If a species gains electrions, it is reduced
LEO, GER
Use oxidation numbers to determine this
Oxidizing agent: species that oxidizes another species, and is itself reduced
Reducing agent: species that reduces another species, and is itself oxidized

16. Combustion reaction
Reaction in which a substance reacts with oxygen, usually accompanied by release of heat and production of a flame
Organic compounds combust to form CO2 and H2O
Metals combust to form metal oxides

17. Molar concentration
Molarity: measure of concentration
= (moles of solute / liters of solution)
Unit: mol/L
Diluting solutions: MiVi = MfVf

18. Gravimetric analysis
Determination of amount of a species by precipitating that species out as an insoluble compound, and weighing the product
Mass precipitated product  moles product  moles unknown species  mass unknown species

19. Volumetric analysis
Titration: method for determining amount of one substance by adding a precise volume of another substance until the two substances completely react
Colored pH indicator often used to detect endpoint
Volume added solution  moles added solution  moles unkn. solution  molarity or grams unkn. solution