1. Chapter 4 Aqueous Reactions and Solution Stoichiometry

2. General Properties of Aqueous Solutions

3. Ionic Compounds in Water

4. Molecular Compounds in Water

5. Strong and Weak Electrolytes Simulation Electrolyte activity

6. Precipitation Reactions

7. Solubility Guidelines

8. Double Displacement (Exchange) Reactions

9. Ionic Equations

10. Acid-Base Reactions Acids and Bases are also common electrolytes, which dissociate in water (soluble/separate) Acids are substances that ionize in aqueous solutions to form hydrogen ions (H + ) thereby increasing the concentration of H + (aq) ions. Because hydrogen atoms consist of only one proton and one electron, is H + simply one proton. Thus acids are often called proton donors. Just like the ionic compounds we have been studying, water also causes the H + (aq) ion to separate from its other half. Molecules of different acids can ionize to form different numbers of H + (aq) ions. Both HCl and HNO 3 are monoprotic acids yeilding only one H + (aq) ion. H 2 SO 4, however, is a diprotic acid that yields two H + (aq) ions. This occurs to two steps however.

11. Acid-Base Reactions cont. Bases are substances that accept (react with) H + (aq) ions. Bases produce hydroxide ions (OH - ) when they dissolve in water. Ionic hydroxide compounds such as NaOH, KOH, and Ca(OH) 2 are among the most common bases, separating just as the ions we have been studying separate. Compounds that do not contain OH - ions can also be bases. For example, ammonia is a common base that when added to water, it accepts an H + from the water and thereby produces an OH - ion.

12. Strong and Weak Acids and Bases Acid and bases that completely dissociate in solution (strong electrolytes) are strong acids and bases. (pg 130 in AP Text) Those that only partially dissociate in solution (weak electrolytes) are weak acid and bases. Not only strong acids are highly reactive though… So strong electrolytes include all soluble ionic compounds and strong acids. Weak electrolytes include weak acids and bases.

13. Neutralization Reactions and Salts When a solution of a strong acid and strong base are mixed a neutralization reaction occurs because the H + the acid combines with the OH - from the base to make water and a salt (ionic compound whose anion comes from an acid and whose cation comes from an base. HCl + NaOH  H 2 O + NaCl Video simulation Also have net ionic equations… Sample Exercise 4.7 pg 133

14. Acid-Base Reactions with Gas Formation Many bases beside OH- react with H+ to for molecular compounds such as the sulfide ion and the carbonate ion. Both of these ions reaction with acids to form gases. See page 134 for net ionic equations. Net Ionic Equation

15. Oxidation-Reduction Reactions Reactions in which electrons are transferred between reactants-redox reactions When an atom, ion or molecule has become more positively charged by losing electrons it is oxidized (oxidation). LEO When an atom, ion or molecule has become more negatively charged by gaining electrons, it is reduced (reduction). GER The oxidation of one substance is always accompanied by the reduction of another substance. As electrons are transferred between them. Calcium Water Demo on each table

16. Oxidation Numbers We need a system to keep track of whether electrons are lost or gained by a substance to determine if a redox reaction has taken place. Oxidation number is the hypothetical charges assigned to the atom before or after a reaction takes place. Oxidation numbers are assigned by the following rules: –For an atom in it elemental form, the oxidation number is always zero. –For any monatomic ion the oxidation number equals the charge on the ion. –The oxidation number of hydrogen is usually +1 when bonded to nonmetals and -1 when bonded to metals. –The oxidation number of fluorine is -1 in all compounds, but the other halogens have positive oxidation numbers when combined with oxygen. –The sum of the oxidation numbers of all atoms in a neutral compound is zero and the sum of the oxidation numbers in a polyatomic ion equals the charge of the ion.

17. Oxidation of Metals by Acids and Salts In this chapter we will look at the oxidation of metals by acids and salts, but in Ch 20 we will look at more complex redox reactions. The reaction of a metal with either an acid or a metal salt is a single displacement reaction. A + BX  AX + B Also has net ionic reactions… Sample Exercise 4.9 pg 140

18. Activity Series of Metals Different metals vary in the ease with which they are oxidized also known as activity. Zn is oxidized by solutions of Cu +2 but Ag is not… meaning that Zn loses electrons more readily than Ag. Elemental metals become ions by losing electrons through oxidation. Therefore, extensive oxidation of metals can lead to the failure of metal machinery parts or structures if the parts are made of metals with high activity (ease of oxidation). A list of metals arranged in order of decreasing ease of oxidation is called an activity series (seen in our previous lab). Alkali and Alkaline Earth Metals are most easily oxidized, elemental form turns into a compound, and therefore are called the active metals. Transition metals are less active and form compounds less readily, called noble metals because of their low reactivity in elemental form, and are therefore used for jewelry and coins. Any metals on the list can be oxidized by the ions of elements below it. Which also means that each metal is reduced by the elements above it.

19. Concentrations of Solutions The behavior of solutions often depends on how much of a solute is in it. Concentration is the term used to explain the amount of solute dissolved in a given quantity of solvent. More solute, greater concentration. Molarity (M) expresses the concentration of a solution as the number of moles of solute in a liter of solution Molarity= moles of solute/volume in liters of solution Sample Exercises 4.11 on pg 144 and practice exercise pg 145. Three quantities needed for the calculation so if we know two of the three and then we can figure out the third. Sample Exercises 4.13 on pg 146

20. Dilutions Solutions that are used routinely in the lab are often purchased in concentrated form. HCl for example is purchased as a 12M solution. Solutions of lower concentrations can them be obtained by adding water, process called dilution. M conc X V conc =M dil X V dil Sample Exercise 4.14 pg 148

21. Solution Stoichiometry If you know the chemical equation and the amount of one reactant consumed in the reaction, you can calculate the quantities of the other reactants and products. Coefficients in a balanced equation give the relative number of moles of reactants and products. So all values given should be converted to moles so that the mole ratio in the equation can be used to do conversion from one reactant to another part of the equation. If you have mass convert into moles before you convert to mass of other substance. If you have molarity, convert into moles as well before you convert to molarity of other substance. See chart on pg 149.

22. Titrations To determine concentration of a particular solute in a solution, chemists often carry out a titration, which involves combining a sample of the solution with a reagent solution of known concentration called a standard solution. Can be conducted using acid-base, precipitation or oxidation-reduction reactions. The point at which stoichiometrically equivalent quantities are brought together is known as the equivalence point for the titration (ratios from balanced equation achieved). Easy to determine the equivalence point for acid and bases (solution becomes neutral and can be detected by indicators such as phenothalein) Practice Exercises on pgs 152-153.