Unit 9 Chemical Equilibrium & Acid-Base Chemistry

Reversible Reactions A chemical reaction in which the products can react to re-form the reactants is called a reversible reaction. A reversible reaction is written using double arrows to show that the reaction is proceeding in both directions. Example:

Dynamic Equilibrium A reversible reaction reaches dynamic equilibrium when the rate of its forward reaction equals the rate of its reverse reaction and the concentrations of its products and reactants remain unchanged. At equilibrium, both reactions continue, but there is no net change in the composition of the system. Visual Concept

Equilibrium  Equal At equilibrium, the rates of the forward and reverse reactions are equal. But the concentrations aren’t necessarily equal. Some reactions reach equilibrium only after almost all reactants are consumed (products are favored.) Others reach equilibrium when only a small percentage of reactants are consumed (reactants are favored.)

Le Châtelier’s Principle Le Châtelier’s Principle: When a system at equilibrium is disturbed, the system shifts in a direction that minimizes the disturbance. A shift in equilibrium will result from a change to any of the following: Concentration Volume/Pressure Temperature

Change in Concentration If you increase the concentration by adding more of a reactant or product, the system will shift to produce less of that substance. If you lower the concentration by removing some of a reactant or product, the system will shift to produce more of that substance.

Change in Volume/Pressure When you increase the pressure (usually by decreasing volume of the container), the system shifts so the least number of gas molecules are formed (less collisions = lower pressure.) When you decrease the pressure, the system will shift so the greatest number of gas molecules are formed.

Change in Temperature For every reversible reaction, one direction is endothermic and the other is exothermic. If the temperature is increased, the endothermic reaction will be favored (because it takes in some of the excess heat.) If the temperature is decreased, the exothermic reaction will be favored (produces heat.)

Le Châtelier’s Principle Sample Problem 2 SO2(g) + O2(g) Û 2 SO3(g) DH° = -198 kJ How will the reversible reaction above shift in response to each of the following stresses? adding more O2 to the container condensing and removing SO3 compressing the gases cooling the container doubling the volume of the container warming the mixture Shift right Shift right Shift right Shift right Shift left Shift left

The Law of Mass Action [C]c [D]d K = [A]a [B]b The relationship between the chemical equation and the concentrations of reactants and products is called the Law of Mass Action. for the general equation aA + bB  cC + dD, the Law of Mass Action is: Lowercase letters represent coefficients. Always products over reactants. Pure solids and pure liquids are not included. [C]c [D]d K = [A]a [B]b

The Equilibrium Constant The equilibrium constant (K) reflects how the concentrations of the reactants and products compare at equilibrium. It can also be a ratio of pressures (in atmospheres) if a reaction involves gaseous reactants and/or products. K is unitless.

The Value of K K > 1 more product molecules present than reactant molecules (the position of equilibrium favors products.) K < 1 more reactant molecules present than product molecules (the position of equilibrium favors reactants.) K = 1 reactant and product particles are present in exact equal concentrations at equilibrium.

Equilibrium Constant Sample Problem Equilibrium concentrations of [H2] = 0.033 M, [I2] = 0.53 M and [HI] = 0.934 M were observed at 445oC for the reaction: H2(g) + I2(g) Û 2 HI(g) Write an equilibrium expression for the above reaction. 2. Calculate the value of Kc for this reaction at 445oC. [HI]2 K = [H2][I2] [HI]2 [0.934]2 Kc = = = 49.9 [H2][I2] [0.033] [0.53]

The Reaction Quotient [C]c [D]d Q = [A]a [B]b When a reaction is not at equilibrium, how do you know in which direction it will proceed? the answer is to compare the equilibrium constant to a ratio of current concentrations called the reaction quotient (Q). for the general equation aA + bB  cC + dD: The non-equilibrium concentrations (or pressures) are used. [C]c [D]d Q = [A]a [B]b

Q vs. K We calculate Q in order to compare it with K. Q < K means the reaction will proceed in the forward direction ([products] increase and [reactants] decrease.) Q > K means the reaction will proceed in the reverse direction ([products] decrease and [reactants] increase.) Q = K means the reaction is at equilibrium ([products] and [reactants] will not change.)

Q, K, and the Direction of Reaction

Reaction Quotient Sample Problem For the reaction below, which direction will it proceed if PI2 = 0.114 atm, PCl2 = 0.102 atm & PICl = 0.355 atm? I2(g) + Cl2(g)  2 ICl(g) Kp = 81.9 First calculate Q: Then, compare it with K: (ICl)2 (0.355)2 Q = = = 10.8 (I2)(Cl2) (0.114) (0.102) Q < K Reaction will proceed to the right 10.8 81.9

Properties of Acids Taste sour. React with metals to release H2 gas. React with bases to produce salts and water. Change the color of acid-base indicators. Conduct electric current.

Properties of Bases Taste bitter. Feel slippery. React with acids to produce salts and water. Change the color of acid-base indicators. Conduct electric current.

Arrhenius Acids and Bases An Arrhenius acid produces hydrogen ions, H+, in aqueous solution. An Arrhenius base produces hydroxide ions, OH−, in aqueous solution. A strong acid (or base) ionizes completely. A weak acid (or base) releases only a few ions.

Arrhenius Theory HCl ionizes in water, producing H+ and Cl– ions NaOH dissociates in water, producing Na+ and OH– ions

Hydronium Ion The H+ ions (protons) produced by the acid are so reactive they cannot exist in water. instead, they react with a water molecule to form a hydronium ion, H3O+. H+ + H2O  H3O+ Chemists use H+ and H3O+ interchangeably.

Arrhenius Acid-Base Reactions The H+ from the acid combines with the OH- from the base to make a molecule of H2O. It is often helpful to think of H2O as H-OH. The cation from the base combines with the anion from the acid to make a salt. acid + base → salt + water HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

Brønsted-Lowry Acids and Bases In a Brønsted-Lowry Acid-Base reaction, an H+ ion (proton) is transferred. Does not have to take place in aqueous solution. Broader definition than Arrhenius. A Brønsted-Lowry acid is a molecule or ion that is a proton donor. A Brønsted-Lowry base is a molecule or ion that is a proton acceptor.

Conjugate Pairs Brønsted-Lowry theory allows for reversible reactions. The original base has an extra H+ after the reaction. It will act as an acid in the reverse process. The original acid has a lone pair of electrons after the reaction. It will act as a base in the reverse process. each reactant and the product it becomes is called a conjugate pair.

Brønsted-Lowry Acid-Base Reactions Sample Problem Identify the Brønsted-Lowry Acids and Bases and Their Conjugates in the Reactions below: a. b. H2SO4 + H2O  HSO4– + H3O+ acid base conjugate base conjugate acid HCO3– + H2O  H2CO3 + HO– base acid conjugate acid conjugate base

Amphoteric Compounds An amphoteric substance is one that can react as either an acid or a base. Example: water Water can act as an acid. Water can act as a base. base acid acid base Visual Concept

Polyprotic Acids Molecules with more than one ionizable H are called polyprotic acids. 1 H = monoprotic, 2 H = diprotic, 3 H = triprotic (Ex: HCl = monoprotic, H2SO4 = diprotic, H3PO4 = triprotic) Polyprotic acids ionize in steps (each ionizable H removed sequentially.) Removing the first H automatically makes removing the second H harder. (Ex: H2SO4 is a stronger acid than HSO4)

Lewis Acids and Bases A Lewis acid is an atom, ion, or molecule that accepts an electron pair to form a covalent bond. A Lewis base is an atom, ion, or molecule that donates an electron pair to form a covalent bond. The Lewis definition is the broadest of the three definitions.

Comparing the Three Definitions Visual Concept

Strong and Weak Acids and Bases Strong acid – fully dissociates in water (almost every molecule breaks up to form H+ ions. Weak acid – partially dissociates in water. Strong base – fully dissociates in water (almost every molecule breaks up to form OH- ions. Weak base – partially dissociates in water. A Strong Acid A Weak Acid

Ionization Constants for Weak Acids and Bases The acid ionization constant (Ka) is the equilibrium constant for the ionization reaction of a weak acid: A base ionization constant (Kb) can also be created for the ionization reaction of a weak base: HA(aq) + H2O(l)  H3O+(aq) + A-(aq) [H3O+] [A-] Ka = [HA] B(aq) + H2O(l)  BH+(aq) + OH-(aq) [BH+] [OH-] Kb = [B]

Acid Ionization Constant Sample Problem Calculate the Ka of a 0.100 M solution of acetic acid with a measured [H3O+] of 1.34 x 10-3 M. HC2H3O2(aq) + H2O(l)  H3O+(aq) + C2H3O2-(aq) For every H3O+ produced, there is also a C2H3O2- produced, so the concentrations must be the same. The equilibrium concentration of original acid is the original concentration decreased by the amount ionized (0.100M – 1.34 x 10-3 M = 0.0987 M) [H3O+][C2H3O2-] [1.34 x 10-3M] [1.34 x 10-3M] Ka= = = 1.82 x 10-5 [HC2H3O2] [0.0987 M]

Ka & Kb and Strength of Acids/Bases The strength of an acid or base is measured by the size of its equilibrium constant when it reacts with H2O.

Six Strong Acids and Bases Because these acids and bases are known to dissociate (ionize) to essentially 100% completion, it is meaningless to connect them to equilibrium: 6 Strong Acids 6 Strong Bases HCl – Hydrochloric Acid Ca(OH)2 – Calcium Hydroxide HBr – Hydrobromic Acid Sr(OH)2 – Strontium Hydroxide HI – Hydroiodic Acid Ba(OH)2 – Barium Hydroxide H2SO4 – Sulfuric Acid LiOH – Lithium Hydroxide HNO3 – Nitric Acid NaOH – Sodium Hydroxide HClO4 – Perchloric Acid KOH – Potassium Hydroxide

Strengths of Conjugate Pairs the stronger an acid is at donating H, the weaker the conjugate base is at accepting H. (i.e. strong acids have weak conjugate bases, and weak acids have strong conjugate bases.) Increasing Basicity Increasing Acidity

Predicting Acid Strength The strength of an acid depends on its tendency to ionize (let go of its hydrogen.) For binary acids, the strength of the acid depends on two factors: The stronger the bond, the weaker the acid. The more polar the bond, the stronger the acid.

Periodic Trends are a BEAR All of these increase in the direction toward their letter: B=Basicity E=Electronegativity, ionization Energy, & Electron Affinity A=Acidity R=Radius B E A R

Predicting Acid Strength (continued) The strength of Oxyacids of the form H-O-Y, where Y is any atom (besides H) bonded to O, depends on two factors: The more electronegative the element Y, the stronger the acid. The greater the number of oxygen atoms bonded to Y, the stronger the acid.

Predicting Acid Strength Sample Problem Predict the relative strengths of the following acids: HCl, HBr, and HI HClO, HBrO, and HIO HNO3 and HNO2 Predict the relative strengths of the following bases: Cl-, Br- and I- H2PO3- and H2PO4- LiOH and Mg(OH)2 HCl < HBr < HI HIO < HBrO < HClO HNO2 < HNO3 I- < Br- < Cl- H2PO4- < H2PO3- LiOH < Mg(OH)2

Autoionization of Water Water is actually an extremely weak electrolyte. About 1 out of every 10 million water molecules form ions through a process called autoionization. H2O + H2O Û H3O+ + OH– In pure water at 25°C, [H3O+] = [OH–] = 10-7M.

Ion Product of Water Kw = [H3O+] [OH-] = 1 x 10-14 The product of the H3O+ and OH– concentrations is always the same number, called the ion product of water (Kw). at 25°C If you measure one of the concentrations, you can calculate the other. As [H3O+] increases the [OH–] must decrease so the product stays constant. Kw = [H3O+] [OH-] = 1 x 10-14

Acidic and Basic Solutions Neutral solutions have equal [H3O+] and [OH–] [H3O+] = [OH–] = 1 x 10-7 Acidic solutions have a larger [H3O+] than [OH–] [H3O+] > 1 x 10-7; [OH–] < 1 x 10-7 Basic solutions have a larger [OH–] than [H3O+] [H3O+] < 1 x 10-7; [OH–] > 1 x 10-7

Acidic, Basic or Neutral? Sample Problem Calculate [OH] at 25°C when [H3O+] = 1.5 x 10-9 M, and determine if the solution is acidic, basic, or neutral. First calculate [OH-]: Then, compare [H3O+] with [OH-]: Kw = [H3O+][OH-] Kw 1.0 x 10-14 [OH-] = = = 6.7 x 10-6 M 1.5 x 10-9 [H3O+] < [H3O+] [OH-] The solution is basic 1.5 x 10-9 M 6.7 x 10-6 M

The pH Scale The acidity/basicity of a solution is often expressed as pH. pH is defined as the negative of the common logarithm of the hydronium ion concentration. pH = −log [H3O+] pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral.

pH of Some Common Substances

The pH Scale Sample Problem Calculate the pH at 25°C when the [OH] = 1.3 x 10-2 M, and determine if the solution is acidic, basic, or neutral First calculate [H3O+]: Then, calculate pH: Kw = [H3O+][OH-] Kw 1.0 x 10-14 [H3O+] = = = 7.7 x 10-13 M 1.3 x 10-2 [OH-] pH = −log [H3O+] = -log(7.7 x 10-13 M) pH = 12.1 The solution is basic

Finding the Ionization Constant from pH Sample Problem A 0.100 M weak acid (HA) solution has a pH of 4.25. Find Ka for the acid. First calculate [H3O+]: Then, calculate Ka: pH = -log [H3O+] 10-pH 10-4.25 = 5.6 x 10-5 M = = [H3O+] HA(aq) + H2O(l)  H3O+(aq) + A-(aq) [H3O+] [A-] (5.6 x 10-5) (5.6 x 10-5) Ka = = = 3.1 x 10-8 [HA] (0.100 - 5.6 x 10-5)

The pOH Scale Another way of expressing the acidity/basicity of a solution is pOH. pOH is defined as the negative of the common logarithm of the hydroxide ion concentration. pOH = −log [OH-] pOH < 7 is basic; pOH > 7 is acidic, pOH = 7 is neutral

Relationship Between pH and pOH The sum of the pH and pOH of a solution is 14. pH + pOH = 14.0

Neutralization Reactions A neutralization reaction is a double displacement reaction in which an acid and a base in an aqueous solution react to produce a salt and water. A salt is an ionic compound made up of a cation from a base and an anion from an acid.

Acid Base Properties of Salts Salts of strong acids and strong bases are neutral. Ex: HCl(aq) + NaOH (aq)  NaCl (aq) + H2O (l) Salts of strong acids and weak bases are acidic. Ex: NH3(aq) + HCl(aq)  NH4Cl (aq) Salts of strong bases and weak acids are basic Ex: 2NaOH(aq)+ H2CO3 (aq)  Na2CO3 (aq) + 2H2O(l) Salts of weak acids and weak bases can be acidic, basic or neutral depending on the relative strength of acids and bases.

Buffers Buffers are solutions that resist changes in pH when an acid or base is added. Buffers contain both a weak acid and its conjugate base (or a weak base and its conjugate acid.) The weak acid neutralizes added base. The conjugate base neutralizes added acid.

Buffering Effectiveness A good buffer should be able to neutralize moderate amounts of added acid or base. However, there is a limit to how much can be added before the pH changes significantly. The buffering capacity is the amount of acid or base a buffer can neutralize. The buffering range is the pH range over which the buffer can be effective. The effectiveness of a buffer depends on: The relative amounts of acid and base. The absolute concentrations of acid and base.

Buffers in Human Blood Many of the chemical reactions that occur in the body are pH-dependent. Ideally, the pH of the blood should be maintained at a slightly basic 7.4. pH below 6.8 or above 7.8 can be fatal. Fortunately, we have buffers in the blood to protect against large changes in pH.

Acid-Base Indicators An acid-base indicator is a chemical dye that changes colors at definite pH values. There are a variety of indicators that change color at different pH levels. A properly selected indicator can be used to visually "indicate" the approximate pH of a sample.

Litmus Paper A common indicator is found on litmus paper. It is red below pH 4.5 and blue above pH 8.2.

Phenolphthalein Phenolphthalein is an organic compound often used as an acid-base indicator. Phenolphthalein is colorless in acidic solutions, but turns pink when the pH is greater than 8.3.

Bromothymol Blue Bromothymol Blue (BTB) is a useful indicator for substances that have a relatively neutral pH (near 7). BTB is yellow in acids, green in neutral solutions, and blue in bases.

Acid-Base Titration Titration is a method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration. The analyte is a measured volume of an acid or base of unknown concentration. The standard solution (titrant) is an acid or base solution whose concentration is known. Standard solution (titrant) analyte

Titration Procedure In a titration procedure, a measured volume of an analyte is placed in a beaker or flask, and initial pH recorded. The standard solution (titrant) is filled in a burette. A couple of drops of an acid-base indicator are added to the flask. The standard solution is slowly added to the unknown solution in the flask. As the two solutions are mixed the acid and the base are neutralized.

Titration Procedure (cont’d) As the base is added to the acid, H+ reacts with OH– to form water. But there is still excess acid present so the color does not change. Once enough base has been added to neutralize all the acid, the indicator changes color. The difficulty is determining when there has been just enough titrant added to complete the reaction…without going over!

The Equivalence Point End point - The point at which an indicator changes color. Equivalence point – The point at which the moles of acid added equals the moles of base that you started with (should be the same as the end point.) An abrupt change in pH occurs at the equivalence point.

Acid-Base Titration Sample Problem The titration of 10.00 mL of HCl solution of unknown concentration requires 12.54 mL of 0.100 M NaOH solution to reach the end point. What is the concentration of the unknown HCl solution? HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) First, write the neutralization reaction 12.54 mL NaOH 1 0.100 1 L NaOH mol NaOH mol HCl 1.25 x 10-3 mol HCl = 1 1000 mL NaOH 1 mol NaOH L NaOH moles 1.25 x 10-3 mol HCl = = 0.125 M HCl M = L 0.010000 L HCl

Titration of Strong Acid with Strong Base Since the salt produced is neutral, the solution at the equivalence point has a pH of 7. the pH starts off low and increases as you add more of the base. The pH doesn't change very much until you get close to the equivalence point. Then it surges upwards very steeply

Titration of Strong Base with Strong Acid This curve is very similar to the previous one for the titration of a strong acid with strong base. The main difference is that the curve starts basic and then turns acidic after the equivalence point (rather than vice-versa.)

Titration of Weak Acid with Strong Base The salt is basic, so equivalence point is at a pH > 7. Before the equivalence point, the solution acts as a buffer. The start of the graph shows a relatively rapid rise in pH but this slows down due to the buffering effect.

Titration of Weak Base with Strong Acid Salt formed is acidic, hence, equivalence point comes at a pH < 7. This curve is very similar to the titration of a weak acid with a strong base. The main differences are that the curve starts basic and has an acidic equivalence point.

Titration of a Polyprotic Acid A polyprotic acid titration will have more than one equivalence point. The first equivalence point represents the titration of the first proton, while the second equivalence point represents the titration of the second proton.

Titration Curves Sample Problem Two acidic solutions were titrated with a strong base. Which curve represents a weak acid and which represents a strong acid? Strong Acid Weak Acid

Acid Rain Rain is naturally somewhat acidic (pH ~5.6) due to atmospheric CO2. Carbon dioxide combines with rainwater to form carbonic acid: CO2 + H2O → H2CO3. Rain water with a pH < 5.6 is called acid rain. Acid rain is linked to damage in ecosystems and structures.

What Causes Acid Rain? Nonmetal oxides such as SO2 and NO2 are acidic: 2 SO2 + O2 + 2 H2O  2 H2SO4 4 NO2 + O2 + 2 H2O  4 HNO3 Processes that produce nonmetal oxide gases increase the acidity of the rain natural – volcanoes, bacterial action. man-made – combustion of fuel

Sources of SO2 from Utilities Weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced.

Damage from Acid Rain Acids react with metals, so acid rain damages bridges, cars, and other metallic structures. Acid reacts with carbonates, so acid rain damages buildings and other structures made of limestone or cement. Acidified lakes affect aquatic life. Acid dissolves and leaches minerals from soil, which weakens and kills trees.

Acid Rain Legislation The 1990 Clean Air Act was passed to reduce acid rain. It requires industries to minimize pollutant gas emissions. As a result, the acidity of rain in the northeast has stabilized and is beginning to be reduced.