1. Unit 11: Reactivity Acid-Base, Precipitation, Oxidation-Reduction(Redox) These are the 3 categories of reactions

2. Properties of Acids Tart or _Sour___________ React with certain_metals_____ to produce hydrogen gas React with bases to produce water and a salt Turns blue litmus to the color __RED_____ Electrolytes = __conduct_______ electricity Acids – vinegar, carbonated drinks, citrus fruit

3. Properties of Bases __Bitter_______ taste Feel _slippery______ React with acids to form water and a salt Turns red litmus to the color _blue_______ Electrolytes = __conduct___ electricity Bases – antacid tablets, household cleaning agents

4. Ions in Solution ___Acidic_______ – contain more H + than OH - _Basic_________ – contain more OH - than H + _Neutral____ – contain equal amounts of H + and OH -

5. Naming Acids - reminder I ate something icky. I bite something delicious. (binary acids) HX “X” ends in -ide Hydro _____ic acid “X” ends in -ate _______ic acid “X” ends in -ite _______ous acid HCl: hydrochloric acid HF: hydrofluoric acid H 2 SO 3 : sulfurous acid HNO 2 : nitrous acid H 2 SO 4 : sulfuric acid HNO 3 : nitric acid

6. Naming Bases - reminder A base is a compound that produces hydroxide ions, (OH - ), when dissolved in water. Bases are named the same way any other ionic compound is named. NaOH – sodium hydroxide Ca(OH) 2 – calcium hydroxide

7. Arrhenius Acids and Bases Svante Arrhenius (Swedish chemist) (1859-1927) 1890’s – formulated first useful theory of acids and bases. Acid: contain H + ; the dissolved compound pr0duces a hydrated hydrogen ion (H 3 O + (aq), hydronium) and an anion. Base(alkali): contains at least one OH - ion that dissociates when dissolved in water. Dissolved compound produces hydroxide ion and a cation. Reaction produces H2O and a dissolved salt. This process is known as Neutralization.

8. Bronsted-Lowry 1920’s – definitions of acids and bases expanded independently by – Danish chemist Johannes Bronsted and English chemist Thomas Martin Lowry More broad definition of acids/bases Acids- hydrogen ion, or proton, donors Bases- hydrogen ion, or proton, acceptors

9. Bronsted-Lowry (cont.) A conjugate acid and conjugate base are formed. Base gains hydrogen ion; particle formed is conjugate acid. The particle that remains when an acid has donated a hydrogen ion is the conjugate base. Example: when HCl gas dissolves and ionizes in water, HCl is the proton donor (acid), water is the proton acceptor (base), the hydronium ion is the conjugate acid, and the chloride ion is the conjugate base. HCl (g) + H 2 O (l) H 3 O + (aq) + Cl - (aq) acid base conjugate acid conjugate base

10. Neutralization Reactions acid + base  a salt + water HCl + NaOH  NaCl + H 2 O Neutralization reactions are a type of double replacement reactions.

11. Water Water can act as an acid or a base H2O as an acid donates one H + to become OH - H2O as a base accepts one H + to become H 3 O + This is called the Hydronium ion (H 3 O + )

12. Neutral Solution [H + ] = hydrogen ion concentration [OH - ] = hydroxide ion concentration Neutral Solution—Any aqueous solution in which the concentration of H +, [H + ], and the concentration of OH -, [OH - ], are equal. In pure water at 25 o C, [H + ] = [OH - ] = 1.0 x 10 -7 M. In any aqueous solution, [H + ] and [OH - ] are interdependent. That is, When [H + ] increases, [OH - ] must decrease, and vice versa.

13. Ion-product Constant for Water (K w ) The product of the concentrations of the hydrogen ions and hydroxide ions in water is called the ion- product constant for water (K w ). K w = [H + ] x [OH - ] = 1.0 x 10 -14 M K w = (1.0 x 10 -7 M) x (1.0 x 10 -7 M) = 1.0 x 10 -14 M If you know one, you can calculate the other. pH scale is derived from the value of K w. Brackets, [ ], denote “concentration of”

14. The pH concept pH is the French abbreviation for “pouvier d’hydrogen.” In English, this translates to the “power of hydrogen.” The pH of a solution is the negative logarithm of the hydrogen-ion concentration. Pure water, pH = 7, is said to be neutral. The formula is: pH = -log[H + ] pH = -log (1 x 10 -7 ) = 7 Likewise…………. pOH = -log[OH - ] Since K w = 1 x 10 -14, for any solution:pH + pOH = 14 pOH = 14 – pH pH = 14 - pOH

15. pH (cont.) Each unit of pH represents a power of 10. Example: Baking soda has pH of 8.4 Milk has pH of 6.4 This difference is 2 pH units. Since each unit represents a power of 10, the [H + ] in milk is 10 2 (or 100) times the [H + ] in baking soda.

16. The pH Scale

17. Strength and Dissociation Strong & weak acids are determined by how much they dissolve into ions in water. Strong acids – completely ionize higher conductivity lower pH than same concentration of weak acid examples: HCl and H2SO4 Weak acids - only partially ionize lower conductivity examples: carbonic acid, acetic acid

18. Strength and Dissociation (cont.) Strong base - higher conductivity higher pH than same concentration of a weak base example: NaOH Weak base - lower conductivity examples: ammonia & sodium bicarbonate

19. Acid-Base Indicators An _indicator__ is an acid or a base that undergoes dissociation in a known pH range. An indicator is useful for measure of pH because its acid form and base form have different colors in solution. Ex: litmus paper, pH paper, phenolphthalein

20. Acid-Base Indicator Table (See pg. 590)

21. pH Meters A __pH Meter__ is used to make rapid, accurate pH measurements. If a pH meter is connected to a computer or chart recorder, it can be used to make a continuous recording of pH changes. Most pH meters are accurate to within 0.01 pH unit of the true pH.

22. Acid Strength The _strength______ of an acid or a base tells you the degree of ionization Strong acids & bases break down into __all__________ ions. (Completely ionized, or dissociated) Ex: Strong acid – hydrochloric acid, sulfuric acid Strong base – Sodium hydroxide Weak acids & bases break down into just a few ions (Partially ionized, or dissociated) Ex: Weak acid – acetic acid, carbonic acid Weak base – Sodium bicarbonate (baking soda)

23. Titration Lab technique used to experimentally determine the molarity of a solution. A titration is a carefully monitored addition of one solution into another solution. Often done by the use of a burette, which allows for accurate measurement needed to neutralize the other solution. An indicator or pH meter is needed to identify at what point the neutralization process is complete. (ex: Phenolphthalein - colorless to pink around pH=9)

24. Oxidation-Reduction Also known as redox reactions Very common in nature – rusting and other types of corrosion are typical examples A chemical reaction in which one reactant loses electrons and another gains electrons. Example: Fe Fe +2 + 2e - (Iron loses e - s.) O 2 + 4e - 2O -2 (Oxygen gains e - s.) Put together: 2Fe + O 2 2FeO Note: In order to be balanced, have total of 4 electrons lost by iron and 4 gained by oxygen.

25. Redox (cont.) The reactant being oxidized (losing electrons) is called the reducing agent. The reactant that is reduced (gaining electrons) is called the oxidizing agent. Typically, redox reactions happen in water or moist surroundings. Note: ALL batteries produce electricity via redox reactions. Regular batteries eventually consume the reactants. Rechargable batteries reverse the redox reaction by adding electricity (electrons), so they last longer. Remember to use Activity Series of Metals for info, too!

26. Precipitation Precipitation reaction may occur when two solutions of IONIC substances are mixed. If one or both pairs of ions in the combined solutions form ionic compounds with very little to no solubility in water (unable to dissolve or very little), then a new, insoluble ionic solid will precipitate out. Note: this is NOT the same as the physical change of crystallization seen in some solutions due to temperature changes.