Chapter 4: Chemical Reactions

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