Unit 9 Acids, Bases, Salts

Properties of Acids Acids (Table K) Dilute aqueous solutions of acids taste sour Lemons (citric acid) Vinegar (acetic acid) Table K lists the acids in decreasing acid strength Common strong acids  HCl (hydrochloric)  HBr (hydrobromic)  H 2 SO 4 (sulfuric)  HNO 3 (nitric) The formulas of acids either begin with an H or end in a -COOH

Properties of Bases Bases (Table L) Dilute aqueous solutions of bases taste bitter Bases have a slippery or soapy feeling Table L lists the bases in decreasing base strength Group 1 and 2 metals + OH (hydroxide ion) are strong bases Examples: Mg(OH) 2, NaOH Most bases end in an OH

Arrhenius Theory Arrhenius Acids Dissociate (break apart) and form the hydrogen ion (H + ) in solution HCl(aq)  H + (aq) + Cl - (aq) The hydrogen ion is also called a proton Not all substances that contain H are acids Example: CH 4 (methane) It is believed that when a hydrogen ion is released into water It will form a covalent bond with water forming the hydronium ion  H + + H 2 O  H 3 O +

Arrhenius Theory Acids Continued Some acids produce more hydrogen ions per molecule than others Monoprotic- 1 hydrogen ion released per molecule (HCl) Diprotic- 2 hydrogen ions released per molecule (H 2 SO 4 ) Triprotic- 3 hydrogen ions released per molecule (H 3 PO 4 )

Arrhenius Theory Bases Bases dissociate in water or react with water to form hydroxide ions (OH - ) NaOH(aq)  Na + (aq) + OH - (aq) NH 3 (aq) + H 2 O(l)  NH 4 + (aq) + OH - (aq) There are certain compounds that contain the hydroxide ion, but are not bases Alcohols- these are covalent compounds that do not produce ions when they dissolve  Examples: CH 3 OH

Salts Ionic substances composed of a positive ion other than hydrogen A negative ion other than hydroxide Examples: NaCl, KBr, CuCl 2

Electrolytes Electrolyte A substance that when dissolved in water produces ions Acids, bases, and salts are all examples of electrolytes because Each type of substance produces ions in solution that can than conduct an electric current The more ions produced, the greater the conductivity Stronger acids and bases make more ions than weaker ones A nonelectrolyte will not dissociate into ions in solution Will not conduct electricity Example: alcohols, sugars

Acid/Base Reactions Reactions of acids with metals Any metal above H 2 on table J will react with an acid to make a salt and hydrogen gas Mg(s) + HCl(aq)  MgCl 2 (aq) + H 2 (g) These reactions are single replacement reactions

Acid/Base Reactions Neutralization Reactions Strong Acid + Strong Base  Salt + Water HCl(aq) + NaOH(aq)  NaCl(aq) + H 2 O(l) This is a double replacement reaction The H + to OH - ratio must be 1:1 The net ionic equation for all neutralization reactions is  H + + OH -  H 2 O All neutralization reactions are exothermic

Titrations Titration is the process of adding measured volumes of an acid or base of known concentration (using a buret) to an acid or base of unknown concentration (in a beaker with indicator) until neutralization occurs Titrations are performed to determine the concentration of the unknown solution Standard solution – solution of known concentration End point- when indicator changes to a color that represents neutralization

Titration Apparatus

Titrations Titration Equation (table T) M A V A = M B V B Ex #1: What is the concentration of an HCl solution if 50.0 mL of 0.25M KOH are needed to completely neutralize 20.0 mL of the HCl solution? HCl is the acid and KOH is the base The ratio of H to OH is 1:1, so we will simply plug in and solve M A = x V A = 20.0 mL M B = 0.25M V B = 50.0 mL (x)(20.0) = (.25)(50.0) x = 0.625M or 0.63M

Titrations Example #2: what is the concentration of a sulfuric acid solution if 125 mL of a 1.0M KOH solution are needed to neutralize 25 mL of sulfuric acid? Sulfuric acid: H 2 SO 4 Base: KOH The ratio of H and OH are 2:1 When the ratio is not 1:1, we must adjust the equation (2)M A V A = M B V B (2)(x)(25) = (1.0)(125) X = 2.5M

pH Scale When pH = 7 solution is neutral (water, sugars) H + = OH - When pH > 7 solution is basic (alkaline) H + < OH - When pH OH - pH is a measurement of the hydrogen (or hydronium) ions in a solution

pH Scale Since this is a logarithmic scale (powers of 10), each 1 point change in pH is a 10x decrease or increase in hydrogen ions When a solution changes from a pH of 6 to a pH of 4, what happens to the concentration of hydrogen ions? Answer: pH is decreasing so it is becoming more acidic (increase in hydrogen ions, decrease in hydroxide ions)

pH Scale To determine the exact amount a substance increases or decreases by 1. Take the pH values and subtract to get a + value hat # becomes the exponent of a power of So in our previous example pH difference is 6-4 = = 100 x more acidic (100 x more hydrogen ions) You Try When a solution changes from a pH of 3 to 6, what happens to the concentration of hydrogen ions? Answer they decrease by a factor of 10 3 = 1000 H + < OH -

Acid/Base Indicators Table M Any indicator is a substance that changes color when it gains or loses protons (hydrogen) They are used to test the acidity or alkalinity of a substance

Acid/Base Indicators Most often used indicators are phenolphthalein and litmus paper 1. Phenolphthalein 1. turns pink in a base 2. Colorless in an acid 2. Litmus Paper 1. Red litmus – stays red in an acid, turns blue in base 2. Blue litmus – stays blue in base, turns red in acid

Acid/Base Indicators To study any indicator use table M On table M 1. the color on the left corresponds to the # on the left 2. The color on the right corresponds to the # on the right 3. Example: A solution has a pH of 6.0 What color would each indicator be in this solution? Methyl orange = Bromothymol blue = Phenolphthalein = Litmus paper = Bromcresol green = Thymol blue =

Alternative Theory Acids are proton (hydrogen) donors Lose hydrogen Bases are proton acceptors Gain hydrogen