Acid Base Equilibria n Acids and bases are found in many common substances and are important in life processes. n Group Work: Make a list of some common acids and bases. How do we know which is which? n There are several models for what constitutes an acid or a base -- three models to be discussed.

14.1 Acids and Bases: A Brief Review n Acid: Base: tastes sourtastes bitter stings skinfeels slippery corrosive to metals releases CO 2 from carbonates turns litmus redturns litmus blue turns phenolphthalein colorlessturns ph. pink n React together to form a salt with loss of the characteristic acid/base properties

Arrhenius Definition Acids produce hydrogen ions (H + ) in aqueous solution. Acids produce hydrogen ions (H + ) in aqueous solution. n Bases produce hydroxide ions (OH - ) when dissolved in water. n Limits to aqueous solutions. n Only one kind of base. n NH 3 ammonia could not be an Arrhenius base.

Arrhenius Definition Not realistic: H + has a radius of cm, which gives a very concentrated charge, so it associates with H 2 O as H(H 2 O) 4 +, which we usually simplify to H 3 O + or H + (aq) Not realistic: H + has a radius of cm, which gives a very concentrated charge, so it associates with H 2 O as H(H 2 O) 4 +, which we usually simplify to H 3 O + or H + (aq) OH – is also associated with H 2 O as OH(H 2 O) 3 – which we usually write as OH – (aq). OH – is also associated with H 2 O as OH(H 2 O) 3 – which we usually write as OH – (aq).

Bronsted-Lowry Definitions n And acid is an proton (H + ) donor and a base is a proton acceptor. n Acids and bases always come in pairs. n HCl is an acid.. n When it dissolves in water it gives its proton to water. n HCl(g) + H 2 O(l) H 3 O + + Cl - n Water is a base makes hydronium ion

Acid/Base Pairs n General equation n HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) n Acid + Base Conjugate acid + Conjugate base n This is an equilibrium. n Competition for H + between H 2 O and A - n The stronger base controls direction. n If H 2 O is a stronger base it takes the H + n Equilibrium moves to right.

Acid/Base Pairs – Group Work n Write a balanced equation showing how the following substances behave as acids in water and identify the conjugate acid-base pairs. HNO 3 HCO 3 - H 3 PO 4 H 2 PO 4 - HNO 3 (aq) + H 2 O(l) ⇌ H 3 O + (aq) + NO 3 - (aq) HNO 3 (aq) + H 2 O(l) ⇌ H 3 O + (aq) + NO 3 - (aq) HCO 3 - (aq) + H 2 O(l) ⇌ H 3 O + (aq) + CO (aq) HCO 3 - (aq) + H 2 O(l) ⇌ H 3 O + (aq) + CO (aq) H 3 PO 4 (aq) + H 2 O(l) ⇌ H 3 O + (aq) + H 2 PO 4 - (aq) H 3 PO 4 (aq) + H 2 O(l) ⇌ H 3 O + (aq) + H 2 PO 4 - (aq) H 2 PO 4 - (aq) + H 2 O(l) ⇌ H 3 O + (aq) + HPO (aq) H 2 PO 4 - (aq) + H 2 O(l) ⇌ H 3 O + (aq) + HPO (aq) n acid 1 base 2 acid 2 base 1

AcidBase StrongestHClO 4 ClO 4 - Weakest acidsH 2 SO 4 HSO 4 - bases acidsH 2 SO 4 HSO 4 - bases HII - HBrBr - HClCl - HNO 3 NO 3 - H 3 O + H 2 O HSO 4 - SO H 2 SO 3 HSO 3 - H 3 PO 4 H 2 PO 4 - HNO 2 NO 2 - HFF - CH 3 CO 2 HCH 3 CO 2 - H 2 CO 3 HCO 3 - H 2 SHS - NH 4 + NH 3 HCNCN - HCO 3 - CO HS - S 2 - H 2 OOH - WeakestNH 3 NH 2 - Strongest acidsOH - O 2 - bases back

Leveling Effect All acids above H 3 O + in the table are strong acids, which dissociate completely in aqueous solution. All acids above H 3 O + in the table are strong acids, which dissociate completely in aqueous solution. All bases below OH - in the table are strong bases, which dissociate completely in aqueous solution. All bases below OH - in the table are strong bases, which dissociate completely in aqueous solution. n The table can be used to make predictions, based on the principle that the stronger acid reacts with the stronger base to form a weaker acid and a weaker base.

Predicting Acid-Base Reactions HCl + HSO 3 - ⇌ H 2 SO 3 + Cl - HCl + HSO 3 - ⇌ H 2 SO 3 + Cl - n stronger stronger weaker weaker acid base We must also consider H 2 O as a possible acid or base. Thus, HNO 3 will transfer its proton to H 2 O, not to Cl - because H 2 O is a stronger base than Cl -. We must also consider H 2 O as a possible acid or base. Thus, HNO 3 will transfer its proton to H 2 O, not to Cl - because H 2 O is a stronger base than Cl -.

Group Work n Write an equation showing the position of equilibrium for the following mixtures. Remember that H 2 O can also be either an acid or a base. HSO 4 - and F - HS - and HCO 3 - HSO F - ⇌ SO HF HSO F - ⇌ SO HF HS - + HCO 3 - ⇌  H 2 S + CO HS - + HCO 3 - ⇌  H 2 S + CO 3 2 -

14.2 Acid dissociation constant K a n The equilibrium constant for the general equation. n HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) n K a = [H 3 O + ][A - ] [HA] n H 3 O + is often written H + ignoring the water in equation (it is implied).

Acid dissociation constant K a n HA(aq) H + (aq) + A - (aq) n K a = [H + ][A - ] [HA] n We can write the expression for any acid. n Strong acids dissociate completely. n Equilibrium far to right. n Conjugate base must be weak.

Back to Pairs n Strong acids n K a is large n [H + ] is equal to [HA] n A - is a weaker base than water n Weak acids n K a is small n [H + ] <<< [HA] n A - is a stronger base than water

Types of Acids n Polyprotic Acids- more than 1 acidic hydrogen (diprotic, triprotic). n Oxyacids - Proton is attached to the oxygen of an ion. n Organic acids contain the Carboxyl group -COOH with the H attached to O n Generally very weak.

Amphoteric n Behave as both an acid and a base. n Water autoionizes n 2H 2 O(l) H 3 O + (aq) + OH - (aq) n K W = [H 3 O + ][OH - ]=[H + ][OH - ] n At 25ºC K W = 1.0 x n In EVERY aqueous solution. n Neutral solution [H + ] = [OH - ]= 1.0 x10 -7 n Acidic solution [H + ] > [OH - ] n Basic solution [H + ] < [OH - ]

14.3 pH Scale n pH= -log[H + ] n Used because [H + ] is usually very small n As pH decreases, [H + ] increases exponentially n Sig figs only the digits after the decimal place of a pH are significant n [H + ] = 1.0 x pH= sig figs n pOH= -log[OH - ] n pKa = -log K

Measuring pH n litmus or pH paper n color changes of indicators n voltage generated by electrodes (pH meter)

pH Indicators

Relationships n K W = [H + ][OH - ] n -log K W = -log([H + ][OH - ]) n -log K W = -log[H + ]+ -log[OH - ] n pK W = pH + pOH So… n K W = 1.0 x n = pH + pOH n [H + ], [OH - ], pH and pOH –Given any one of these we can find the other three through equilibrium relationships (K w )

BasicAcidicNeutral [H + ] pH Basic [OH - ] pOH

14.5 pH and K a of Acid Solutions n Always write down the major ions in solution. n Remember these are equilibria. n Remember the chemistry. n Don’t try to memorize; there is no one way to do this. Apply good chemistry!

Strong Acid: HNO 3 HNO 3 is completely dissociated into the ions, H 3 O + and NO 3 - (work sample 14.7) HNO 3 is completely dissociated into the ions, H 3 O + and NO 3 - (work sample 14.7)

Strong Acids n HBr, HI, HCl, HNO 3, H 2 SO 4, HClO 4 n ALWAYS WRITE THE MAJOR SPECIES n Completely dissociated n [H + ] = [HA] i n [OH - ] is going to be small because of equilibrium K w = = [H + ][OH - ]

Changes in pH with Dilution n Group Work: pH for factors of 10 dilution? n What is the pH of 1.0 M HCl? n pH = 0.00 n What is the pH of 0.10 M HCl (a 1:10 dilution)? n pH = 1.00 n What is the pH of M HCl? n pH = 2.00

Changes in pH with Dilution What is the pH of 1.0 x M HCl? What is the pH of 1.0 x M HCl? n pH = 3.00 What is the pH of 1.0 x M HCl? What is the pH of 1.0 x M HCl? n pH = 4.00 What is the pH of 1.0 x M HCl? What is the pH of 1.0 x M HCl? n pH = 5.00 What is the pH of 1.0 x M HCl? What is the pH of 1.0 x M HCl? n pH = 5.996

Changes in pH with Dilution What is the pH of 1.0 x M HCl? What is the pH of 1.0 x M HCl? n pH = What is the pH of 1.0 x M HCl? What is the pH of 1.0 x M HCl? n pH = n Why does the pH stop changing at a value of about 7? n Water has a pH of 7 due to autodissociation, so it is never possible to get a pH higher than 7 by addition of water. –If [HA] < water contributes the H +

14.5 pH of Weak Acids n Except for the strong acids, most acids do not ionize completely. These acids are called weak acids. HF(aq) + H 2 O(l) ⇌ H 3 O + (aq) + F - (aq) HF(aq) + H 2 O(l) ⇌ H 3 O + (aq) + F - (aq)

14.5 pH of Weak Acids n K a will be small. n ALWAYS WRITE THE MAJOR SPECIES. n It will be an equilibrium problem from the start. n Determine whether most of the H + will come from the acid or the water. n Compare K a or K w n Rest is just like chapter 13.

Example n Calculate the pH of 2.0 M acetic acid HC 2 H 3 O 2 with a Ka 1.8 x10 -5 n Calculate pOH, [OH - ], [H + ]

A mixture of Weak Acids n The process is the same. n Determine the major species. n The stronger acid will predominate. n Bigger K a if concentrations are comparable n Calculate the pH of a mixture 1.20 M HF (Ka = 7.2 x ) and 3.4 M HOC 6 H 5 (Ka = 1.6 x )

Some Weak Acids

Percent dissociation n = amount dissociated x 100 initial concentration n For a weak acid percent dissociation increases as acid becomes more dilute. n Calculate the % dissociation of 1.00 M and M Acetic acid (Ka = 1.8 x n As [HA] 0 decreases [H + ] decreases but % dissociation increases. –Le Chatelier principle with dilution (pg 642)

The other way n What is the Ka of a weak acid that is 8.1 % dissociated as M solution? n Compare to Sample Exercise 14.11

pH and Ka n One way to find the value of K a for a weak acid is using concentration & pH data. n The pH of M HNO 2 is What is K a of HNO 2 ? [H 3 O + ] = = M [H 3 O + ] = = M HNO 2 + H 2 O  H 3 O + + NO 2 - HNO 2 + H 2 O  H 3 O + + NO 2 - I C E K a = [H 3 O + ][NO 2 - ]/[HNO 2 ] K a = [H 3 O + ][NO 2 - ]/[HNO 2 ] K a = (0.0149) 2 /0.485 = 4.58 x K a = (0.0149) 2 /0.485 = 4.58 x

14.6 Bases n The OH - is a strong base. n Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved. n The hydroxides of alkaline earths Ca(OH) 2 etc. are strong dibasic bases, but they don’t dissolve well in water. n Used as antacids because [OH - ] can’t build up.

Bases without OH - n Bases are proton acceptors. n NH 3 + H 2 O NH OH - n It is the lone pair on nitrogen that accepts the proton. n Many weak bases contain N n B (aq) + H 2 O(l) BH + (aq) + OH - (aq) n K b = [ BH + ][ OH - ] [ B ]

Strength of Bases n Hydroxides are strong. n Others are weak. n Smaller K b = weaker base. n Consider a solution of 4.0 M pyridine (Kb = 1.7 x ) –What are the major species present in solution? –Determine the [OH-] and pH of the solution. N:N:

14.7 Polyprotic acids n Always dissociate stepwise. n Denoted Ka 1, Ka 2, Ka 3, etc… (Table 14.4) n H 2 CO 3 + H 2 O H + + HCO 3 - Ka 1 = 4.3 x n HCO H 2 O H + + CO 3 -2 Ka 2 = 4.3 x n Conjugate base in 1 st step is acid in 2 nd.

Polyprotic Acid Observations n The first H + comes of much easier than the second. n K a for the first step is much bigger than K a for the second.

Calculate the Concentration… n …of all the ions in a solution of 5.00 M sulfurous acid. n Ka 1 = 1.5 x n Ka 2 = 1.0 x n Find the pH of the solution n See Sample Exercise for another example.

Calculate the Concentration n Of all the ions in a solution of 1.00 M Arsenic acid H 3 AsO 4 n Ka 1 = 5.0 x n Ka 2 = 8.0 x n Ka 3 = 6.0 x n In calculations we can normally ignore the second dissociation.

Sulfuric acid, an interesting case n Calculate the pH and concentration of all species in a 1.5 M solution of H 2 SO 4 –Pg 651, Table 14.4 contains Ka’s for polyprotics n Calculate the concentrations in a 1.5 x M solution of H 2 SO 4 n See Sample Exc & 17 n Summary pg. 655 –In first step H 2 SO 4 is a strong acid. –2 nd step, it is a weak acid: Ka 2 = 1.2 x 10 -2

14.8 Acid/Base Properties of Salts n Salts are ionic compounds. –Recall Model: Ionic compounds dissociate into ions in solution. The ions move around independently. n These cations and anions can (but do not have to!) act as acids or bases depending on how they react with water (referred to as a hydrolysis reaction) n Ch 14, #99 discuss on WB’s

Salts that make Neutral Solutions n Salts containing the cation from a strong base and the anion from a strong acid are neutral. n Hydrolysis is not observed with ions derived from strong acids or bases: Cations of group I and II (except Be 2 + ) Anions: Cl -, Br -, I -, NO 3 -, ClO 4 - n for example NaCl, KNO 3 n There is no equilibrium established from these ions.

Basic Salts n If the anion of a salt is the conjugate base of a weak acid the solution will be basic. n Consider an aqueous solution of NaF n The aqueous species are Na +, F -, and H 2 O n F - + H 2 O HF + OH - n The equilibrium constant for hydrolysis is just a K a or K b, depending on the type of hydrolysis.

Basic Salts n K b =[HF][OH - ] [F - ] n but Ka = [H + ][F - ] [HF] n K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF]

Basic Salts n K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF]

Basic Salts n K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF]

Basic Salts n K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] n K a x K b =[OH - ] [H + ]

Basic Salts n K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] n K a x K b =[OH - ] [H + ] n K a x K b = K w

pH of Basic Salts n The anion of a weak acid is a weak base. n The CN - ion competes with OH - for the H + n Relative base strength: OH - > CN - > H 2 O n Calculate the pH of a solution of 1.00 M NaCN. n Ka of HCN is 6.2 x

Acidic salts n If the cation of a salt is the conjugate acid of a weak base the solution will be acidic. n The same development as bases leads to n K a x K b = K W n Calculate the pH of a solution of 0.40 M NH 4 Cl (the K b of NH x ). n Other acidic salts are those of highly charged metal ions.

Acidic Salts – Metal ions n Ions are often modified when dissolved in solution. Examine the photos of Fe(III) salts and solutions. Why do they have different colors? Fe(NO 3 ) 3. 6H 2 O contains pink Fe(H 2 O) Fe(NO 3 ) 3. 6H 2 O contains pink Fe(H 2 O) Solutions may hydrolyze to give yellow Fe(H 2 O) 5 OH 2 + or even reddish brown Fe(H 2 O) 3 (OH) 3 Solutions may hydrolyze to give yellow Fe(H 2 O) 5 OH 2 + or even reddish brown Fe(H 2 O) 3 (OH) 3 FeCl 3. 6H 2 O contains ions such as yellow Fe(H 2 O) 5 Cl 2 + FeCl 3. 6H 2 O contains ions such as yellow Fe(H 2 O) 5 Cl 2 +

Hydrolysis of Metals n Hydrolysis is more important for more highly charged ions and smaller ions. n Hydrogen from water easier to remove.

Hydrolysis of Metals n Table below gives values of K a for metal ions: IonRadiusKa IonRadiusKa Na + 95 pm3.3 x Na + 95 pm3.3 x Li + 60 pm1.5 x Be pm3.2 x Mg pm3.8 x Ba pm1.5 x Cr pm9.8 x Zr pm6.0 x n Greater values of K a for ions with larger charge and smaller size.

Hydrolysis of Metals n Calculate the pH of a 0.20 M CrCl 3 solution. (From the table, the Ka value for Cr 3+ is 9.8 x )

Anion of weak acid, cation of weak base n If both of the ions have acid/base properties, compare the Ka of the cation to the Kb of the anion n K a > K b acidic n K a < K b basic n K a = K b neutral n Predict whether an aqueous solution of ammonium cyanide will be acidic, basic, or neutral. n Summary Table 14.6, pg 660

14.9 Structure and Acid base Properties n Draw lewis structures (that obey octet rule) for the following oxyacids. HClO 4 K a = Large (~10 7 ) HClO 3 K a = ~1 HClO 2 K a = 1.2 x 10 –2 HClOK a = 3.5 x 10 –8 n Can you explain the relative strengths of these acids using your knowledge of atomic structure and bonding?

Strength of oxyacids ClOH O O O Electron Density

Strength of oxyacids ClOH O O Electron Density

Strength of oxyacids Electron Density ClOHO

Strength of oxyacids Electron Density ClOH

Strength of oxyacids n The more oxygen bonded to the central atom, the more acidic the hydrogen. n The oxygens are electronegative n Pulls electron density away from hydrogen n HClO 4 > HClO 3 > HClO 2 > HClO –Remember that the H is attached to an oxygen atom.

14.9 Structure and Acid base Properties n Draw lewis structures (that obey octet rule) for the following oxyacids. HOClK a = 4 x 10 –8 HOBrK a = 2 x 10 –9 HOIK a = 2 x 10 –11 HOCH3K a = ~ 10 –15 n Can you explain the relative strengths of these acids using your knowledge of atomic structure and bonding?

14.9 Structure and Acid base Properties n Any molecule with an H in it is a potential acid. n The stronger the X-H bond the less acidic (compare bond dissociation energies). n The more polar the X-H bond the stronger the acid (use electronegativities). n The more polar H-O-X bond -stronger acid.

14.10 Acid-Base Properties of Oxides n Non-metal oxides dissolved in water can make acids. n SO 3 (g) + H 2 O(l) H 2 SO 4 (aq) n Ionic oxides dissolve in water to produce bases. n CaO(s) + H 2 O(l) Ca(OH) 2 (aq)

14.11 Lewis Acids and Bases n Most general definition. n Acids are electron pair acceptors. n Bases are electron pair donors. BF F F :N:N H H H

Lewis Acids and Bases n Boron triflouride wants more electrons. BF F F :N:N H H H

Lewis Acids and Bases n Boron triflouride wants more electrons. n BF 3 is Lewis acid NH 3 is a Lewis base. B F F F N H H H

Lewis Acids and Bases Al +3 ( ) H H O Al ( ) 6 H H O