Acids and Bases Chapter 16 Acids and Bases John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten

Acids and Bases HW CHAPTER 16 –ACID-BASE EQUILIBRIUM Bronsted Lowry 15,16,17,19,23,25 Kw 29 pH scale 35, 37 Strong acids and bases 41, 43 (a to c) 45, 47 WEAK ACIDS 53, 55, 57, 61-a, 65 73, 75, 81, 85, to 110 red problems ONLY

Acids and Bases Electrolytes Substances that dissociate into ions when dissolved in water. A nonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so.

Acids and Bases Electrolytes and Nonelectrolytes Soluble ionic compounds tend to be electrolytes.

Acids and Bases Electrolytes and Nonelectrolytes Molecular compounds tend to be nonelectrolytes, except for acids and bases.

Acids and Bases Electrolytes A strong electrolyte dissociates completely when dissolved in water. A weak electrolyte only dissociates partially when dissolved in water.

Acids and Bases Strong Electrolytes Are… Strong acids

Acids and Bases Strong Electrolytes Are… Strong acids Strong bases

Acids and Bases Some Definitions Arrhenius  Acid:Substance that, when dissolved in water, increases the concentration of hydrogen ions.  Base:Substance that, when dissolved in water, increases the concentration of hydroxide ions.

Acids and Bases Some Definitions Brønsted–Lowry  Acid:Proton donor  Base:Proton acceptor

Acids and Bases A Brønsted–Lowry acid… …must have a removable (acidic) proton. A Brønsted–Lowry base… …must have a pair of nonbonding electrons.

Acids and Bases If it can be either…...it is amphiprotic. HCO 3 − HSO 4 − H2OH2O

Acids and Bases What Happens When an Acid Dissolves in Water? Water acts as a Brønsted–Lowry base and abstracts a proton (H + ) from the acid. As a result, the conjugate base of the acid and a hydronium ion are formed.

Acids and Bases 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.

Acids and Bases Conjugate Acid-Base Pairs Whatever is left of the acid after the proton is donated is called its conjugate base. Similarly, whatever remains of the base after it accepts a proton is called a conjugate acid. Consider –After HA (acid) loses its proton it is converted into A - (base). Therefore HA and A - are a conjugate acid-base pair. –After H 2 O (base) gains a proton it is converted into H 3 O + (acid). Therefore, H 2 O and H 3 O + are a conjugate acid-base pair. Conjugate acid-base pairs differ by only one proton.

Acids and Bases Example – Identify the acid and the base in each equation, and identify each acid-base pair. 1.HNO 3 + NH 3  NO NH CH 3 COOH + OH -  H 2 O + CH 3 COO - Identify the acid and base for the reverse reaction in each example.

Acids and Bases Acid and Base Strength Strong acids are completely dissociated in water.  Their conjugate bases are quite weak. Weak acids only dissociate partially in water.  Their conjugate bases are weak bases.The weaker the acid the stronger its conjugate base.

Acids and Bases Acid and Base Strength Substances with negligible acidity do not dissociate in water.  Their conjugate bases are exceedingly strong.

Acids and Bases Relative Strengths of Acids and Bases The stronger the acid, the weaker the conjugate base. H + is the strongest acid that can exist in equilibrium in aqueous solution. OH - is the strongest base that can exist in equilibrium in aqueous solution.

Acids and Bases Acid and Base Strength In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base. HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl − (aq) H 2 O is a much stronger base than Cl −, so the equilibrium lies so far to the right K is not measured (K>>1).

Acids and Bases Acid and Base Strength Acetate is a stronger base than H 2 O, so the equilibrium favors the left side (K<1). HC 2 H 3 O 2 (aq) + H 2 OH 3 O + (aq) + C 2 H 3 O 2 − (aq)

Acids and Bases Autoionization of Water As we have seen, water is amphoteric. In pure water, a few molecules act as bases and a few act as acids. This is referred to as autoionization. H 2 O (l) + H 2 O (l) H 3 O + (aq) + OH − (aq)

Acids and Bases Ion-Product Constant The equilibrium expression for this process is K c = [H 3 O + ] [OH − ] This special equilibrium constant is referred to as the ion-product constant for water, K w. At 25°C, K w = 1.0  10 −14

Acids and Bases In most solutions [H + (aq)] is quite small. We define In neutral water at 25  C, pH = pOH = In acidic solutions, [H + ] > 1.0  10 -7, so pH < In basic solutions, [H + ] The higher the pH, the lower the pOH, the more basic the solution. The pH Scale

Acids and Bases pH Therefore, in pure water, pH = −log (1.0  10 −7 ) = 7.00 An acid has a higher [H 3 O + ] than pure water, so its pH is <7 A base has a lower [H 3 O + ] than pure water, so its pH is >7.

Acids and Bases pH These are the pH values for several common substances.

Acids and Bases Other “p” Scales The “p” in pH tells us to take the negative log of the quantity (in this case, hydrogen ions). Some similar examples are  pOH −log [OH − ]  pK w −log K w

Acids and Bases Watch This! Because [H 3 O + ] [OH − ] = K w = 1.0  10 −14, we know that −log [H 3 O + ] + −log [OH − ] = −log K w = or, in other words, pH + pOH = pK w = 14.00

Acids and Bases February 28 Measuring pH: a/b indicators pH meter Section 16.5 Strong a/b – key points Section 16.6 Weak acids  Ka  Problems a) Calculating Ka and % Ionization from measured pH and initial concentration

Acids and Bases HW WEAK ACIDS 53, 55, 57, 61-a, 65

Acids and Bases Daily Quiz 1. Write the formation of nitrous acid from its anhydride 2. Calculate the pH of an aqueous solution of LiOH that has a pH of 12.5

Acids and Bases How Do We Measure pH? For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.

Acids and Bases Most pH and pOH values fall between 0 and 14. There are no theoretical limits on the values of pH or pOH. (e.g. pH of 2.0 M HCl is ) Examples: 1.Consider a solution with [H + ] = 6.2 x Calculate the pH, pOH, and [OH - ] Is this solution acidic or basic? 2.Consider a solution with pOH = Calculate the pH, [H + ], and [OH - ] Is this solution acidic or basic?

Acids and Bases Measuring pH Most accurate method to measure pH is to use a pH meter. However, certain dyes change color as pH changes. These are indicators. Indicators are less precise than pH meters. Many indicators do not have a sharp color change as a function of pH.

Acids and Bases How Do We Measure pH? For less accurate measurements, one can use  Litmus paper “Red” paper turns blue above ~pH = 8 “Blue” paper turns red below ~pH = 5  An indicator

Acids and Bases Strong Acids You will recall that the seven strong acids are HCl, HBr, HI, HNO 3, H 2 SO 4, HClO 3, and HClO 4. These are, by definition, strong electrolytes and exist totally as ions in aqueous solution. For the monoprotic strong acids, [H 3 O + ] = [acid].

Acids and Bases Strong Bases Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca 2+, Sr 2+, and Ba 2+ ).

Acids and Bases Strong bases are strong electrolytes and dissociate completely in solution. The pOH of a strong base is given by the initial molarity of the hydroxide ion. Be careful of stoichiometry. In order for a hydroxide to be a base, it must be soluble. Bases do not have to contain the OH - ion: O 2- (aq) + H 2 O(l)  2OH - (aq) H - (aq) + H 2 O(l)  H 2 (g) + OH - (aq) N 3- (aq) + 3H 2 O(l)  NH 3 (aq) + 3OH - (aq)

Acids and Bases Strong basic solutions Ionic metal oxides Na2O and CaO are used in industry to produce strong basic solutions. Find the pH of a solution formed by dissolving 0.01 mol of Na2O in enough water to produce a liter of solution. 12.3

Acids and Bases Examples: Calculate the pH and pOH of each: M NaOH M HCl M KOH M HNO x M RbOH M HClO 4

Acids and Bases Examples: Calculate the pH and pOH of each: pHpOH M NaOH M HCl M KOH M HNO x M RbOH M HClO

Acids and Bases Weak acids are only partially ionized (dissociated) in solution. There is a mixture of ions and unionized acid in solution. Therefore, weak acids are in equilibrium: Weak Acids

Acids and Bases Dissociation Constants For a generalized acid dissociation, the equilibrium expression would be This equilibrium constant is called the acid-dissociation constant, K a. [H 3 O + ] [A − ] [HA] K c = HA (aq) + H 2 O (l) A − (aq) + H 3 O + (aq)

Acids and Bases Dissociation Constants The larger the Ka the stronger the acid (i.e. the more ions are present at equilibrium relative to unionized molecules). If Ka >> 1, then the acid is completely ionized and the acid is a strong acid

Acids and Bases Percent Ionization Percent ionization is another method to assess acid strength. For the reaction

Acids and Bases Calculating K a from pH Weak acids are simply equilibrium calculations. The pH is used to calculate the equilibrium concentration of H +. –Write the balanced chemical equation clearly showing the equilibrium. –Write the equilibrium expression. –Use an ICE Chart / find the [H+] from the pH –Find the value for K a.

Acids and Bases Calculating K a from the pH The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is Calculate K a for formic acid at this temperature. We know that [H 3 O + ] [COO − ] [HCOOH] K a =

Acids and Bases Calculating K a from the pH The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is Calculate K a for formic acid at this temperature. To calculate K a, we need the equilibrium concentrations of all three things. We can find [H 3 O + ], which is the same as [HCOO − ], from the pH.

Acids and Bases Calculating K a from the pH pH = −log [H 3 O + ] 2.38 = −log [H 3 O + ] −2.38 = log [H 3 O + ] 10 −2.38 = 10 log [H 3 O + ] = [H 3 O + ] 4.2  10 −3 = [H 3 O + ] = [HCOO − ]

Acids and Bases Calculating K a from pH Now we can set up a table… [HCOOH], M[H 3 O + ], M[HCOO − ], M Initially Change −4.2    10 −3 At Equilibrium 0.10 − 4.2  10 −3 = =  10 −3

Acids and Bases Calculating K a from pH [4.2  10 −3 ] [0.10] K a = = 1.8  10 −4

Acids and Bases Calculating Percent Ionization Percent Ionization =  100 In this example [H 3 O + ] eq = 4.2  10 −3 M [HCOOH] initial = 0.10 M [H 3 O + ] eq [HA] initial

Acids and Bases Calculating Percent Ionization Percent Ionization =   10 − = 4.2%

Acids and Bases Examples: Calculate K a. 1.A 0.25 M solution of a hypothetical acid, HA, has a pH of x A 0.10 M solution of acetic acid (CH 3 COOH) has a pH of x 10 -5

Acids and Bases  Using Ka to calculate pH from initial concentration  5% Rule. When to make approximations and when to use quadratic equation.  Percent ionization  Polyprotic acids

Acids and Bases Using K a to Calculate pH Percent ionization relates the equilibrium H + concentration, [H + ] eqm, to the initial HA concentration, [HA] 0. The higher percent ionization, the stronger the acid. Percent ionization of a weak acid decreases as the molarity of the solution increases. For acetic acid, 0.05 M solution is 2.0 % ionized whereas a 0.15 M solution is 1.0 % ionized. Use ICE Charts to determine [H + ] -log [H + ] = pH Substitute into the equilibrium constant expression and solve. Remember to turn x into pH if necessary.

Acids and Bases Calculating pH from K a Calculate the pH of a 0.30 M solution of acetic acid, HC 2 H 3 O 2, at 25°C. HC 2 H 3 O 2 (aq) + H 2 O (l) H 3 O + (aq) + C 2 H 3 O 2 − (aq) K a for acetic acid at 25°C is 1.8  10 −5.

Acids and Bases Calculating pH from K a The equilibrium constant expression is [H 3 O + ] [C 2 H 3 O 2 − ] [HC 2 H 3 O 2 ] K a =

Acids and Bases Calculating pH from K a We next set up a table… [C 2 H 3 O 2 ], M[H 3 O + ], M[C 2 H 3 O 2 − ], M Initially Change−x−x+x+x+x+x At Equilibrium 0.30 − x  0.30 xx We are assuming that x will be very small compared to 0.30 and can, therefore, be ignored.

Acids and Bases Calculating pH from K a Now, (x) 2 (0.30) 1.8  10 −5 = (1.8  10 −5 ) (0.30) = x  10 −6 = x  10 −3 = x

Acids and Bases Calculating pH from K a pH = −log [H 3 O + ] pH = −log (2.3  10 −3 ) pH = 2.64

Acids and Bases 5% RULE As a general rule if x is greater than 5% than the initial value we can not considered negligible and a quadratic equation has to be solved. If x is less than 5% it is ok to considered negligible. SOLVE THE PROBLEM CONSIDERING X NEGLIGIBLE AND CHECK!

Acids and Bases Examples – Calculate pH, pOH and percent dissociation for each of the following: M HF (K a = 7.2 x ) pH = 1.80, pOH = 12.20, 4.5% ionized M HF pH = 1.63, pOH = 12.37, 3.1% ionized M NH 4 + (from NH 4 Cl) (K a = 5.6 x ) pH = 4.73, pOH = 9.27,.0030% ionized M HNO 2 (K a = 4.0 x ) pH = 1.68, pOH = 12.32, 1.9% ionized M HNO 3 pH = -0.78, pOH = 14.78, 100% ionized

Acids and Bases Percent Ionization Percent ionization is another method to assess acid strength. For the reaction

Acids and Bases

Acids and Bases Polyprotic Acids Have more than one acidic proton. If the difference between the K a for the first dissociation and subsequent K a values is 10 3 or more, the pH generally depends only on the first dissociation.

Acids and Bases Dilute Sulfuric Acid Solutions First Dissociation is Strong (large K a ) Second Dissociation is weak (K a2 = 1.2 x )  Make ICE chart for second dissociation – initial [H + ] is equal to initial [HSO 4 - ]  Determine x (using quadratic if necessary)  Solve for [H + ] eq Example – Calculate the pH of M H 2 SO 4

Acids and Bases Polyprotic Acids

Acids and Bases Weak bases Relationship between Ka and Kb Hydrolysis of salts

Acids and Bases Weak Bases Bases react with water to produce hydroxide ion.

Acids and Bases Weak Bases The equilibrium constant expression for this reaction is [HB] [OH − ] [B − ] K b = where K b is the base-dissociation constant.

Acids and Bases K b can be used to find [OH − ] and, through it, pH.

Acids and Bases Types of Weak Bases Bases generally have lone pairs or negative charges in order to attack protons. Most neutral weak bases contain nitrogen. Amines are related to ammonia and have one or more N- H bonds replaced with N-C bonds (e.g., CH 3 NH 2 is methylamine). Anions of weak acids are also weak bases. Example: ClO - is the conjugate base of HClO (weak acid):

Acids and Bases Calculating pH and pOH of Weak Base Solutions Use ICE Chart to determine [OH - ] eq Use [OH - ] to find pOH, and thus pH Examples – Calculate pOH and pH of each of the following solutions: M NH 3 (K b = 1.8 x ) pH = 11.22pOH = M HONH 2 (K b = 1.1 x ) pH = 9.81pOH = 4.19

Acids and Bases pH of Basic Solutions What is the pH of a 0.15 M solution of NH 3 ? [NH 4 + ] [OH − ] [NH 3 ] K b = = 1.8  10 −5 NH 3 (aq) + H 2 O (l) NH 4 + (aq) + OH − (aq)

Acids and Bases pH of Basic Solutions Tabulate the data. [NH 3 ], M[NH 4 + ], M[OH − ], M Initially At Equilibrium x  0.15 xx

Acids and Bases pH of Basic Solutions (1.8  10 −5 ) (0.15) = x  10 −6 = x  10 −3 = x 2 (x) 2 (0.15) 1.8  10 −5 =

Acids and Bases pH of Basic Solutions Therefore, [OH − ] = 1.6  10 −3 M pOH = −log (1.6  10 −3 ) pOH = 2.80 pH = − 2.80 pH = 11.20

Acids and Bases Relationship between strength of acid and conjugate base: When two reactions are added to give a third, the equilibrium constant for the third reaction is the product of the equilibrium constants for the first two: Reaction 1 + reaction 2 = reaction 3 has Relationship Between K a and K b

Acids and Bases For a conjugate acid-base pair Therefore, the larger the K a, the smaller the K b. That is, the stronger the acid, the weaker the conjugate base. Taking negative logarithms:

Acids and Bases

Acids and Bases Nearly all salts are strong electrolytes. Salts exist entirely of ions in solution. Acid-base properties of salts are a consequence of the reaction of their ions in solution. The reaction in which ions produce H + or OH - in water is called hydrolysis. Acid-Base Properties of Salt Solutions

Acids and Bases HYDROLYSIS OF SALTS STRONG PARENT ACID-PARENT BASE  NEUTRAL SALT STRONG PARENT ACID – WEAK PARENT BASE  ACIDIC SALT WEAK PARENT ACID – STRONG PARENT BASE  BASIC SALT WEAK AND WEAK DEPENDS ON WHICH IS GREATER Ka OR Kb

Acids and Bases An Anion’s Ability to React with Water Anions, X -, can be considered conjugate bases from acids, HX. For X - comes from a strong acid, then it is neutral. If X - comes from a weak acid, then The pH of the solution can be calculated using equilibrium! Anions from strong acids are neutral. Anions from weak acids are basic.

Acids and Bases Reactions of Cations with Water Cations with acidic protons (like NH 4 + ) will lower the pH of a solution. Most metal cations that are hydrated in solution also lower the pH of the solution.

Acids and Bases Reactions of Cations with Water Attraction between nonbonding electrons on oxygen and the metal causes a shift of the electron density in water. This makes the O-H bond more polar and the water more acidic. Greater charge and smaller size make a cation more acidic.

Acids and Bases Hydrolysis of Metal Ions Metal ions are positively charged and attract water molecules (via the lone pairs on O). The higher the charge, the smaller the metal ion and the stronger the M-OH 2 interaction. Hydrated metal ions act as acids: The pH increases as the size of the ion increases (e.g. Ca 2+ vs. Zn 2+ ) and as the charge increases (Na + vs. Ca 2+ and Zn 2+ vs. Al 3+ ).

Acids and Bases Hydrolysis of Metal Ions

Acids and Bases A Cation’s Ability to React with Water Polyatomic cations with ionizable protons can be considered conjugate acids of weak bases. Some metal ions react in solution to lower pH. Combined Effect of Cation and Anion in Solution An anion from a strong acid has no acid-base properties. An anion that is the conjugate base of a weak acid will cause an increase in pH (basic ion).

Acids and Bases A cation that is the conjugate acid of a weak base will cause a decrease in the pH of the solution (acidic ion). Metal ions will cause a decrease in pH except for the alkali metals and alkaline earth metals. When a solution contains both cations and anions from weak acids and bases, use K a and K b to determine the final pH of the solution. Use ICE charts to calculate pH and pOH (as for weak acid and weak base solutions)

Acids and Bases Effect of Cations and Anions 1.An anion that is the conjugate base of a strong acid will not affect the pH. 2.An anion that is the conjugate base of a weak acid will increase the pH. 3.A cation that is the conjugate acid of a weak base will decrease the pH.

Acids and Bases Effect of Cations and Anions 4.Cations of the strong Arrhenius bases will not affect the pH. 5.Other metal ions will cause a decrease in pH. 6.When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the affect on pH depends on the K a and K b values.

Acids and Bases Examples – State whether each will be acidic, basic, or neutral in aqueous solution: 1.NaClO 2.Fe(NO 3 ) 3 3.KBr 4.LiClO 4 5.NaHCO 3 6.CaCl 2 7.NH 4 NO 3 8.NiI 2

Acids and Bases Examples – State whether each will be acidic, basic, or neutral in aqueous solution: 1.NaClObasic 2.Fe(NO 3 ) 3 acidic 3.KBrneutral 4.LiClO 4 neutral 5.NaHCO 3 basic 6.CaCl 2 neutral 7.NH 4 NO 3 acidic 8.NiI 2 acidic

Acids and Bases Examples – Calculate pH and pOH for each solution: M NH 4 Cl M KBr M KC 2 H 3 O M NaHCO 3

Acids and Bases Factors that Affect Acid Strength Consider H-X. For this substance to be an acid we need: H-X bond to be polar with H  + and X  - (if X is a metal then the bond polarity is H  -, X  + and the substance is a base), the H-X bond must be weak enough to be broken, the conjugate base, X -, must be stable. The 3 factors are then 1)Polarity of H—X bond 2)Bond strenght 3)Stability of conjugate base

Acids and Bases Binary Acids H—X strength most important factor The bigger the atom X the weaker the strength of the bond the more acidic the acid. HF is a weak acid because the bond energy is high, HCl is more acidic because the bond strength is weaker. Acid strength increases down a group. In a period the acidity strength increases with the increase in the polarization of the bond. That depends on the electronegativity of the element. In a period the acidity increases across a period

Acids and Bases Base Strength Conversely, base strength decreases across a period and down a group. The electronegativity difference between C and H is so small that the C- H bond is non-polar and CH 4 is neither an acid nor a base.

Acids and Bases Binary Acids

Acids and Bases Oxyacids In oxyacids, in which an OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.

Acids and Bases For a series of oxyacids, acidity increases with the number of oxygens.

Acids and Bases Carboxilic Acids Resonance in the conjugate bases of carboxylic acids stabilizes the base and makes the conjugate acid more acidic.

Acids and Bases The strength of carboxilic acid increase as the number of electronegative atoms in the acid increases. Trifluoroacetic Ka=5 x 10 -1

Acids and Bases Brønsted-Lowry acid is a proton donor. Focusing on electrons: a Brønsted-Lowry acid can be considered as an electron pair acceptor. Lewis acid: electron pair acceptor (the acceptor in a coordinate bond). Lewis base: electron pair donor ( electron rich – the donor in a coordinate bond). Note: Lewis acids and bases do not need to contain protons. Therefore, the Lewis definition is the most general definition of acids and bases. Lewis Acids and Bases

Acids and Bases Lewis acids generally have an incomplete octet (e.g. BF 3 ). Transition metal ions are generally Lewis acids. Lewis acids must have a vacant orbital (into which the electron pairs can be donated). Compounds with  -bonds can act as Lewis acids: H 2 O(l) + CO 2 (g)  H 2 CO 3 (aq)

Acids and Bases Lewis Acids Lewis acids are defined as electron-pair acceptors. Atoms with an empty valence orbital can be Lewis acids.

Acids and Bases Lewis Bases Lewis bases are defined as electron-pair donors. Anything that could be a Brønsted–Lowry base is a Lewis base. Lewis bases can interact with things other than protons, however.