HCl(aq)  H+(aq)+ Cl-(aq) Definitions Arrhenius - In aqueous solution… Acids increase hydrogen ion concentration [H+] HCl(aq)  H+(aq)+ Cl-(aq) H Cl O – + HCl(aq) + H2O(l) H3O+(aq)+ Cl-(aq) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

KOH(aq)  K+(aq)+ OH-(aq) Definitions Arrhenius - In aqueous solution… Bases increase hydroxide ion concentration [OH-] KOH(aq)  K+(aq)+ OH-(aq) (aq) (aq) (aq) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor conjugate acid conjugate base base acid

Conjugate Acids and Bases: From the Latin word conjugare, meaning “to join together.” Reactions between acids and bases always yield their conjugate bases and acids.

More examples of Conjugate Pairs

Strong and Weak Acids/Bases The strength of an acid (or base) is determined by the amount of IONIZATION. HNO3, HCl, H2SO4 and HClO4 are among the only known strong acids.

Strong and Weak Acids/Bases Weak acids are much less than 100% ionized in water. One of the best known is acetic acid = CH3CO2H

“Strong” = complete dissociation in solution “Weak” = incomplete dissociation in solution

Strong and Weak Acids/Bases Strong Base = 100% dissociated in water. NaOH (aq) → Na+ (aq) + OH- (aq) CaO Other common strong bases include KOH and Ca(OH)2 CaO (lime) + H2O --> Ca(OH)2 (slaked lime)

Strong and Weak Acids/Bases Weak base = less than 100% ionized in water One of the best known weak bases is ammonia NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)

Extra (nontestable) information… Factors Affecting Acid Strength The more polar the H-X bond and/or the weaker the H-X bond, the more acidic the compound. Acidity increases from left to right across a row and from top to bottom down a group.

Yet another reason water is special It’s called “amphoteric” because H2O can function as both an ACID and a BASE. Equilibrium constant for water = Kw Kw = [H+] [OH-] = 1.00 x 10-14 at 25 oC

Acid – Base Concentrations pH = 3 pH = 11 10-1 H+ pH = 7 OH- 10-7 concentration (moles/L)‏ H+ OH- OH- H+ 10-14 [H+] > [OH-] [H+] = [OH-] [H+] < [OH-] acidic solution neutral solution basic solution Timberlake, Chemistry 7th Edition, page 332

pH of Common Substance More basic More acidic 14 1 x 10-14 1 x 10-0 0 pH [H1+] [OH1-] pOH 14 1 x 10-14 1 x 10-0 0 13 1 x 10-13 1 x 10-1 1 12 1 x 10-12 1 x 10-2 2 11 1 x 10-11 1 x 10-3 3 10 1 x 10-10 1 x 10-4 4 9 1 x 10-9 1 x 10-5 5 8 1 x 10-8 1 x 10-6 6 6 1 x 10-6 1 x 10-8 8 5 1 x 10-5 1 x 10-9 9 4 1 x 10-4 1 x 10-10 10 3 1 x 10-3 1 x 10-11 11 2 1 x 10-2 1 x 10-12 12 1 1 x 10-1 1 x 10-13 13 0 1 x 100 1 x 10-14 14 NaOH, 0.1 M Household bleach Household ammonia Lime water Milk of magnesia Borax Baking soda Egg white, seawater Human blood, tears Milk Saliva Rain Black coffee Banana Tomatoes Wine Cola, vinegar Lemon juice Gastric juice More basic 7 1 x 10-7 1 x 10-7 7 More acidic

(also uses a logarithmic scale) Richter Scale – Earthquakes 6 7 5 . Area Radius Richter Scale Diameter pH [H+] 1000000 564 7 1128 mm 1 0.1 M 100000 178.4 6 356.8 2 0.01 M 10000 56.4 5 112.8 3 0.001 M 1000 17.8 4 35.6 4 0.0001 M 100 5.64 3 11.28 5 0.00001 M 10 1.784 2 1.128 6 0.000001 M 1 0.564 1 ------ 7 0.0000001 M 4 3 2 1

pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14 Kw = [H+] [OH-] = 1 x10-14 The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution • pH and H+ concentration are related as follows: pH = –log10[H+] or [H+] = 10–pH • pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00 • pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7 • pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7 • The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows: pH = –log10[OH–] or [OH–] = 10–pOH • A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00 • The sum of the pH and the pOH for a neutral solution at 25ºC is 7.00 + 7.00 = 14.00 pKw = –log Kw = –log([H3O+] [OH–]) = (–log[H3O+]) + (–log[OH–]) = pH + pOH • At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature pOH pOH = -log[OH-] [OH-] [OH-] = 10-pOH

(less acidic than vinegar, What is the pH of a solution whose pOH is 11.09? Is this solution acidic or basic? pH + pOH = 14 pH + 11.09 = 14 Solution is acidic  (less acidic than vinegar, more than orange juice) pH = 14 – 11.09 = 2.91

pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14 Kw = [H+] [OH-] = 1 x10-14 The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution • pH and H+ concentration are related as follows: pH = –log10[H+] or [H+] = 10–pH • pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00 • pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7 • pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7 • The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows: pH = –log10[OH–] or [OH–] = 10–pOH • A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00 • The sum of the pH and the pOH for a neutral solution at 25ºC is 7.00 + 7.00 = 14.00 pKw = –log Kw = –log([H3O+] [OH–]) = (–log[H3O+]) + (–log[OH–]) = pH + pOH • At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature pOH pOH = -log[OH-] [OH-] [OH-] = 10-pOH

pH = 0.620 Solution is acidic  pH = - log [H+] = -log [0.240M H+] Calculate the pH of a 0.240 M nitric acid solution. Is this solution acidic or basic? pH = - log [H+] = -log [0.240M H+] pH = 0.620 Solution is acidic  (less acidic than battery acid, more than stomach fluids)

pH of Common Substances vinegar 2.8 water (pure)‏ 7.0 soil 5.5 gastric juice 1.6 carbonated beverage 3.0 drinking water 7.2 bread 5.5 1.0 M NaOH (lye)‏ 14.0 orange 3.5 potato 5.8 blood 7.4 1.0 M HCl milk of magnesia 10.5 apple juice 3.8 urine 6.0 detergents 8.0 - 9.0 bile 8.0 lemon juice 2.2 tomato 4.2 milk 6.4 ammonia 11.0 seawater 8.5 coffee 5.0 bleach 12.0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 acidic neutral basic [H+] = [OH-] Timberlake, Chemistry 7th Edition, page 335

pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14 Kw = [H+] [OH-] = 1 x10-14 The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution • pH and H+ concentration are related as follows: pH = –log10[H+] or [H+] = 10–pH • pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00 • pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7 • pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7 • The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows: pH = –log10[OH–] or [OH–] = 10–pOH • A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00 • The sum of the pH and the pOH for a neutral solution at 25ºC is 7.00 + 7.00 = 14.00 pKw = –log Kw = –log([H3O+] [OH–]) = (–log[H3O+]) + (–log[OH–]) = pH + pOH • At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature pOH pOH = -log[OH-] [OH-] [OH-] = 10-pOH

What is the hydroxide ion concentration of a solution with a hydronium ion concentration of 3.3 x 10-10M? Is this solution acidic or basic? Kw = [H+] [OH-] = 1.0 x10-14 1.0 x10-14 = [3.3 x 10-10] [OH-] [OH-] = 3.0 x 10-5 M Solution is basic  (less basic than ammonia, more than sea water)

pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14 Kw = [H+] [OH-] = 1 x10-14 The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution • pH and H+ concentration are related as follows: pH = –log10[H+] or [H+] = 10–pH • pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00 • pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7 • pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7 • The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows: pH = –log10[OH–] or [OH–] = 10–pOH • A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00 • The sum of the pH and the pOH for a neutral solution at 25ºC is 7.00 + 7.00 = 14.00 pKw = –log Kw = –log([H3O+] [OH–]) = (–log[H3O+]) + (–log[OH–]) = pH + pOH • At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature pOH pOH = -log[OH-] [OH-] [OH-] = 10-pOH

Calculate the [H+] of a solution with a pH of 4.60 pH = - log [H+] 4.60 = -log [H+] [H+] = 2.51 x 10-5 M 2nd log -4.6 10- 4.60 You can check your answer by working backwards. pH = - log [2.51x10-5 M] pH = 4.60

pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14 Kw = [H+] [OH-] = 1 x10-14 The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution • pH and H+ concentration are related as follows: pH = –log10[H+] or [H+] = 10–pH • pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00 • pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7 • pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7 • The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows: pH = –log10[OH–] or [OH–] = 10–pOH • A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00 • The sum of the pH and the pOH for a neutral solution at 25ºC is 7.00 + 7.00 = 14.00 pKw = –log Kw = –log([H3O+] [OH–]) = (–log[H3O+]) + (–log[OH–]) = pH + pOH • At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature pOH pOH = -log[OH-] [OH-] [OH-] = 10-pOH

Why do chemists use titrations? Quantitative analysis — used to determine the amounts or concentrations of substances present in a sample by using a combination of chemical reactions and stoichiometric calculations

Titration standard solution unknown solution Definition Analytical method in which a standard solution is used to determine the concentration of an unknown solution. Quantitative analysis — used to determine the amounts or concentrations of substances present in a sample by using a combination of chemical reactions and stoichiometric calculations Titration – A method in which a measured volume of a solution of known concentration, called the titrant, is added to a measured volume of a solution containing a compound whose concentration is to be determined (the unknown) – Reaction must be fast, complete, and specific (only the compound of interest should react with the titrant) – Equivalence point — point at which exactly enough reactant has been added for the reaction to go to completion Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Buret stopcock Erlenmeyer flask

Titration Vocabulary Titrant The substance added to the analyte in a titration (a standard solution) Analyte The substance being analyzed Equivalence point The point in a titration at which the quantity of titrant is exactly sufficient for stoichiometric reaction with the analyte.

Analyte Acid-Base Titration Titrant If the concentration of the titrant is known, then the unknown concentration of the analyte can be determined. Acid-Base Titration Titrant In acid-base titrations, a buret is used to deliver measured volumes of an acid or base solution of known titration (the titrant) to a flask that contains a solution of a base or an acid, respectively, of unknown concentration (the unknown). If the concentration of the titrant is known, then the concentration of the unknown can be determined. Plotting the pH changes that occur during an acid-base titration against the amount of acid or base added produces a titration curve; the shape of the curve provides important information about what is occurring in solution during the titration. Before addition of any strong base, the initial [H3O+] equals the concentration of the strong acid. Addition of strong base before the equivalence point, the point at which the number of moles of base (or acid) added equals the number of moles of acid (or base) originally present in the solution, decreases the [H3O+] because added base neutralizes some of the H3O+ present. Addition of strong base at the equivalence point neutralizes all the acid initially present and pH = 7.00; the solution contains water and a salt derived from a strong base and a strong acid. Addition of a strong base after the equivalence causes an excess of OH– and produces a rapid increase in pH. A pH titration curve shows a sharp increase in pH in the region near the equivalence point and produces an S-shaped curve; the shape depends only on the concentration of the acid and base, not on their identity. For the titration of a monoprotic strong acid with a monobasic strong base, the volume of base needed to reach the equivalence point can be calculated from the following relationship: moles of base = moles of acid (volume)b (molarity)b = (volume)a (molarity)a VbMb = VaMa Analyte

Buret Reading

Acidic, basic, or neutral??

The “perfect pink” for a titration with phenolphthalein Phenolphthalein is clear in acidic solutions but pink in basic solutions!

point at which exactly enough reactant has been added for the solution to be neutralized and no more Indicator - changes color to indicate pH change pink pH Example… phenolphthalein is colorless in acid and pink in basic solution Endpoint = 7 colorless Volume base added

Equivalence point (endpoint) Point at which equal amounts of H+ and OH- have been added. Determined by… indicator color change dramatic change in pH Most common acids and bases are not intensely colored – Rely on an indicator Endpoint — point at which a color change is observed, which is close to the equivalence point in an acid-base titration Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

moles H+ = moles OH- Titration At endpoint... Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

MA = 0.66M HCl Titration Acid (H+) Base (OH-) M = ? M = 1.3M 25.5 mL of 1.3M KOH are required to neutralize 50.0 mL of HCl. Find the molarity of HCl. MA = 0.66M HCl Acid (H+) M = ? V = 50.0 mL Base (OH-) M = 1.3M V = 25.5 mL A reaction in which an acid and a base react in stoichiometric amounts to produce water and a salt Strengths of the acid and base determine whether the reaction goes to completion 1. Reactions that go to completion a. Reaction of any strong acid with any strong base b. Reaction of a strong acid with a weak base c. Reaction of weak acid with a weak base 2. Reaction that does not go to completion is a reaction of a weak acid or a weak base with water Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Typical Example... pH Titration Data NaOH added (mL) pH Titration of an Acid With an Base 0.00 1.00 10.00 1.37 20.00 1.95 22.00 2.19 24.00 2.70 25.00 7.00 26.00 11.30 28.00 11.75 30.00 11.96 40.00 12.36 50.00 12.52 14.0 phenolphthalein - pink 12.0 10.0 8.0 pH equivalence point 6.0 4.0 phenolphthalein - colorless Solution of NaOH Solution of NaOH 2.0 Na+ OH- 0.0 0.0 10.0 20.0 30.0 40.0 Volume of 0.100 M NaOH added (mL) H+ Cl- Solution of HCl 25 mL