ACIDS AND BASES

Acids and Bases Two important classification of compounds - Acids and Bases Properties of ACIDS Taste Sour/Tart Stings and burns the skin Reacts with bases Turns blue litmus paper red Reacts with metals to form H2 gas Neutralizes Bases Donates H+ Conduct electricity. Can be strong or weak electrolytes in aqueous solution Properties of BASES Taste Bitter Feels slippery on skin Reacts with acids Turns red litmus blue Doesn’t react with metals Neutralizes Acids Accepts H+ Chapters 19 and 9.4

Acids Affect Indicators Blue litmus paper turns red in contact with an acid. Chapters 19 and 9.4

Acids React with Active Metals Acids react with active metals to form salts and hydrogen gas: HCl(aq) + Mg(s) → MgCl2(aq) + H2(g) Chapters 19 and 9.4

Sulfuric Acid – H2SO4 Highest volume production of any chemical in the U.S. Used in the production of paper Used in production of fertilizers Used in petroleum refining Chapters 19 and 9.4

Nitric Acid – HNO3 Used in the production of fertilizers Used in the production of explosives Nitric acid is a volatile acid – its reactive components evaporate easily Stains proteins (including skin!) Chapters 19 and 9.4

Hydrochloric Acid - HCl Used in the “pickling” of steel Used to purify magnesium from sea water Part of gastric juice, it aids in the digestion of proteins Sold commercially as “Muriatic acid” Chapters 19 and 9.4

Phosphoric Acid – H3PO4 A flavoring agent in sodas Used in the manufacture of detergents Used in the manufacture of fertilizers Not a common laboratory reagent Chapters 19 and 9.4

Acetic Acid – HC2H3O2 Used in the manufacture of plastics Used in making pharmaceuticals Acetic acid is the acid present in household vinegar Chapters 19 and 9.4

Bases Affect Indicators Red litmus paper turns blue in contact with a base. Phenolphthalein turns purple in a base. Chapters 19 and 9.4

Sodium hydroxide (lye), NaOH Examples of Bases Sodium hydroxide (lye), NaOH Potassium hydroxide, KOH Magnesium hydroxide, Mg(OH)2 Calcium hydroxide (lime), Ca(OH)2 Chapters 19 and 9.4

Bases Neutralize Acids Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl. 2 HCl + Mg(OH)2 MgCl2 + 2 H2O Chapters 19 and 9.4

Acid-Base Theories Arrhenius (1883) NaOH HCl H2SO4 CO32- NH3 H2O Brønsted–Lowry (1923) BF3 AlI3 Lewis (1923-38)

Arrhenius Definition of A & B An acid is a substance that dissociates in water to produce hydrogen ions (H+) Example: Hydrochloric Acid (HCl) A base is a substance that dissociates in water to produce hydroxide ions (OH-) Example: Sodium Hydroxide (NaOH)

What happens when placed in water??? Chapters 19 and 9.4

Svante Arrhenius (1859-1927) Chapters 19 and 9.4

Brønsted-Lowry Definition An acid is any substance that can donate H+ ions. A base is any substance that can accept H+ ions. Expansion on Arrhenius definition of A&B Defines A&B as not necessarily in water solution Does not contain OH- ; covers bases such as ammonia NH3 Chapters 19 and 9.4

Johannes Bronsted Thomas Lowry (1879-1947) (1874-1936) Chapters 19 and 9.4

Brønsted-Lowry Definition H+ is really only just a proton (no electrons or neutrons), so definition is often in terms of protons Brønsted-Lowry Base is a proton acceptor Brønsted-Lowry Acid is a proton donor Monoprotic Acids can only donate 1 H+: HCl Diprotic Acids can donate 2 H+: H2SO4 Triprotic Acids can donate 3 H+: H3PO4 Polyprotic acids = diprotic and triprotic acids

Chapters 19 and 9.4

Chapters 19 and 9.4

Conjugate Acid-Base Pairs Acid-base reactions with water proceed in both directions: Example: NH3 (g) + H2O (l) NH4+ (aq) + OH- (aq) Acid loses H+ = conjugate base Example: H2O (acid) loses its H+, turning it into OH- (conjugate base) Base gains H+ = conjugate acid Example: NH3 (base) gains an H+, turning it into NH4+ (conjugate acid) Chapters 19 and 9.4

Acids and bases come in pairs A “conjugate base” is the remainder of the original acid, after it donates it’s hydrogen ion A “conjugate acid” is the particle formed when the original base gains a hydrogen ion Indicators are weak acids or bases that have a different color from their original acid and base Chapters 19 and 9.4

The Hydronium Ion Water can pick up a H+ ion to form a hydronium ion: H+ + H2O ® H3O+ H3O+ = hydronium ion With acids, water is a Brønsted-Lowry base (accepts protons Example: HCl (g) + H2O(l) ® H3O+ (aq) + Cl- (aq) With bases, water is a Brønsted-Lowry acid (donates protons) Example: NH3 (g) + H2O (l) ® NH4+ (aq) + OH- (aq) Compound that can act as either a proton donor or acceptor = amphoteric Chapters 19 and 9.4

Lewis Acids and Bases Gilbert Lewis (1875-1946) Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction Lewis Acid - electron pair acceptor Lewis Base - electron pair donor Most general of all 3 definitions; acids don’t even need hydrogen!

Lewis Acids and Bases Lewis acid = accepts a pair of electrons during a reaction Lewis base = donates a pair of electrons during a reaction Covers acids and bases not covered by Brønsted-Lowry definition Table 19.4, Pg. 592, Text Type Acid Base Arrhenius H+ producer OH- producer Brønsted-Lowry H+ (proton) donor H+ acceptor Lewis Electron-pair acceptor Electron-pair donor Chapters 19 and 9.4

The pH Scale pH is based on the concentration of the hydronium ion in a solution Concentrations range from 100 M (strong) to 10-14 M (weak) of [H3O+] pH ranges from 0 to 14 If [H30+] concentration = 10-2 M, pH = 2 If [H30+] concentration = 10-10 M, pH = 10 Water is 10-7 M, pH = 7 pH of 0-6 = Acidic pH of 7 = Neutral pH of 8-14 = Basic Chapters 19 and 9.4

pH …..defined as the negative base-10 logarithm of the hydronium ion concentration. pH = −log [H3O+] For pure H2O: [1.0  10−7 ] = 7.0 Problem: Calculate pH when [H3O+] = 2.3 x 10-3 M 10x log log - 2.3 * 10^- 3 = 2.64 base 10 log Problem: Calculate [H3O+] when pH = 2.3 ? antilog 10x log 10^ -2.3 = 5.0 * 10-3

pH and pOH Calculations

Chapters 19 and 9.4

pH These are the pH values for several common substances. pH + pOH = 14

pH and pOH equilibrium in pure Water [H3O+] [OH−] = Kw [1.0  10−7 ] [1.0  10−7 ] = 1.0  10−14 −log [H3O+] + −log [OH−] = −log Kw pH + pOH = pKw Because in pure water [H3O+] = [OH−], Kw = [1.0  10−7 ] [1.0  10−7 ] = 1.0  10−14 + 7 = 14

pH and pOH equilibrium in Water to which Acids & Bases are Added Add base OH- H2O(l) + H2O(l) H3O+(aq) + OH−(aq) [H3O+] [OH−] H2O Kw = [1.0  10−7 ] [1.0  10−7 ] = 1.0  10−14 Kw = [1.0  10−8 ] [1.0  10−6 ] = 1.0  10−14 pH + pOH = pKw 8 + 6 = 14

Measuring pH Why measure pH? Solutions we use – swimming pools soil conditions for plants medical diagnosis soaps and shampoos, etc. Sometimes we can use indicators, other times we might need a pH meter Chapters 19 and 9.4

Acid-Base Strength An acid or a base is considered strong if they completely dissociate into ions (H+ and OH-) in water Strong Acids HCl and H2SO4 Strong Bases Hydroxides, e.g. NaOH Conjugate acid-base pairs have an inverse relationship (works for both acids and bases) The stronger the acid, the weaker the conjugate base The weaker the acid, the stronger the conjugate base Chapters 19 and 9.4

Strong Acids and Bases to know

Measuring pH with wide-range paper 1. Moisten indicator strip with a few drops of solution, by using a stirring rod. 2.Compare the color to the chart on the vial – read the pH value. Chapters 19 and 9.4

Acid-Base Indicators Although useful, there are limitations to indicators: usually given for a certain temperature (25 oC), thus may change at different temperatures what if the solution already has color, like paint? the ability of the human eye to distinguish colors is limited Chapters 19 and 9.4

Some of the many pH Indicators and their ranges Chapters 19 and 9.4

Red Cabbage Juice as an indicator Red cabbage juice mixed with baking soda (left) and with vinegar (right). On the top, a drop of unmixed juice. Chapters 19 and 9.4

Acid-Base Indicators A pH meter may give more definitive results some are large, others portable works by measuring the voltage between two electrodes; typically accurate to within 0.01 pH unit of the true pH needs to be calibrated Chapters 19 and 9.4

Acids Neutralize Bases HCl + NaOH → NaCl + H2O -Neutralization reactions ALWAYS produce a salt and water. -Of course, it takes the right proportion of acid and base to produce a neutral salt Chapters 19 and 9.4

Acid-Base Properties of Salts Neutralization reaction = reaction of an acid and a base Acid + base react – form a salt and water Type of salt depends on reactants Acid Base Salt Strong Neutral Weak Acidic Basic Neutral, basic, or acidic

Chapters 19 and 9.4

Chapters 19 and 9.4

Buffers Buffers are solutions in which the pH remains relatively constant, even when small amounts of acid or base are added made from a pair of chemicals: a weak acid and one of it’s salts; or a weak base and one of it’s salts Chapters 19 and 9.4

Chapters 19 and 9.4

Acid rain Chapters 19 and 9.4

Causes of emissions Chapters 19 and 9.4

pH readings nationwide Chapters 19 and 9.4

Acid rain effects limestone and Marble Chapters 19 and 9.4

Effects of Acid Rain on Marble (calcium carbonate) George Washington: BEFORE George Washington: AFTER Chapters 19 and 9.4

Naming Acids and Bases Naming Acids: Three Rules Naming Bases Name of anion ends in –ide Acid name begins with hydro- Stem of anion has suffix –ic Name of anion ends in –ite Stem of anion has suffix –ous Name of anion ends in –ate All three end with the word “acid” Naming Bases Named just like ionic compounds – cation + anion Chapters 19 and 9.4

Chapters 19 and 9.4