15 Acids and Bases Contents 15-1 The Bronsted-Lowry Definitions 15-2 The Ion Product of Water, Kw 15-3 The pH and Other “p” Scales 15-4 Concentrations of Hydrogen Ions in Aqueous Solutions of Acids 15-7 The Common Ion Effect 15-8 Buffer Solutions 15-9 How Indicators Work

Acids and Bases Arrhenius Definitions － Acids produce hydrogen ions in aqueous solution, and bases produce hydroxide ions in aqueous solution. Brønsted-Lowry Definitions － An acid is a proton (H + ) donor, and a base is a proton acceptor. Lewis Acid-Base Definitions － An acid as an electron pair acceptor and a base as an electron donor.

strong acid formula hydrochloric HCl hydrobromic HBr hydroiodic HI Nitric HNO 3 chloric HClO 3 perchloric HClO 4 sulfuric H 2 SO 4 Sulfuric acid is the only strong acid that is diprotic, meaning it has two protons to donate.

Acid H + donor Base H + acceptor 15-1 The Bronsted-Lowry Definitions Terms to know: proton - H + ion hydronium ion - H 3 O + ; results from water reacting with H + H 2 O + H + --> H 3 O + conjugate base - whatever is left from an acid after a proton has been donated conjugate acid - whatever has been formed when a proton has been accepted by a base

Acids are hydrogen ion (H +) donors Bases are hydrogen ion (H + ) acceptors HCl + H 2 O H 3 O + + Cl - donor, acid acceptor, base + + HCl/Cl - ; H 3 O + / H 2 O :Conjugate Acid-Base Pair An acid-base reaction consists of the transfer of a proton for an acid to a base.

In acidic solutions, the protons (H + ) that are released into solution will not remain alone due to their large positive charge density and small size. They are attracted to the negatively charged electrons on the oxygen atoms in water, and form hydronium ions. H + (aq) + H 2 O (l) = H 3 O + (l) [H + ] = [H 3 O + ] O H..

H atom and H + ion atomic structure? Use symbols H + and H 3 O + to represent the same thing

Conjugate Acid-Base Pair Conjugate acid-base pair consists of two substances related to each other by the donating and accepting of a single proton. Acids and bases that are related by loss or gain of H +. NH 3 +HCl NH Cl - NH 4 + /NH 3, HCl/ Cl -

A species like water that can react either as an acid or as a base is said to be amphoteric. HF(aq) + H 2 O(l) H 3 O + (aq) + F - (aq) Acid 1 Base 2 Acid 2 Base 1 NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq) Base 1 Acid 2 Acid 1 Base 2

Summary 1.Acids and Bases can be molecules or ions 2.Not just defined in aqueous solutions 3.Conjugate pairs differ by H + 4.Some substances are both, AMPHOTERIC depending on the reaction.

Autoionization: A reaction in which a substance reacts with itself to form ions. Water is the common amphoteric substance. HF(aq) + H 2 O(l) H 3 O + (aq) + F - (aq) H 2 O as base. NH 3 (aq) + H 2 O(l) NH OH - (aq) H 2 O as acid. K w =[H 3 O + ][OH - ]=1× ion-product constant (25 ℃ ) pH=-log[H + ] 15-2 The Ion Product of Water, Kw

Water –The most common amphoteric substance water can act as both an acid and a base water can autoionize: –H 2 O + H 2 O H 3 O + + OH - –one water molecule acts as an acid (H + donor), the other acts as an acid (an H + acceptor) –K w = [H 3 O + ][OH - ] = 1.00 x o C dissociation or ion product constant for water

K w –In any aqueous solution at 25 o C, the product of [H + ] and [OH - ] will be 1.0 x –So if you know the [H + ], you can figure out the [OH - ] and vice versa –If [H + ] = [OH - ], the solution is neutral –If [H + ] > [OH - ], the solution is acidic –If [H + ] < [OH - ], the solution is basic

pH scale –Because the [H + ] in any solution is generally quite small, it is easier to use the pH scale to represent a solution ’ s acidity. –pH comes from the Danish … potenz or strength of the H + ion –The pH of a solution is usually defined as the negative of the base 10 logarithm. –pH = - log[H + ] –pOH = - log [OH - ] 15-3 The pH and Other “p” Scales

Acids and Bases pH is a log scale –when the pH changes by one, the [H + ] concentration changes by a power of 10. A solution with a pH of 3 has 10 times more H + than a solution with a pH of 4, and 100 times more H + than a solution with a pH of 5. –As pH decreases, the [H + ] increases. –Rule for significant figures for logarithms - the number of places after the decimal point is equal to the number of significant figures in the original number pH = - log 1.0 x M (2 significant figures in 1.0 x ) pH = 9.00 ( 2 places after the decimal point for significant figures)

Methods for Measuring the pH of an Aqueous Solution (a) pH paper (b) Electrodes of a pH meter

Strong and Weak Acids Strong:100% dissociation good H + donor equilibrium lies far to right (HNO 3 ) generates weak base (NO 3 - ) Weak:<100% dissociation not-as-good H + donor equilibrium lies far to left (CH 3 COOH) generates strong base (CH 3 COO - )

15-4 Concentrations of Hydrogen Ions in Aqueous Solutions of Acids Strong acids: For dense and moderately dilute solutions (1.0× to 0.01M) of strong acids that have only one ionizable hydrogen, [H + ] = stoichiometric concentration of the strong acid. HCl(aq) H + (aq) + Cl - (aq)

acidbase conjugate acid conjugate base HA(aq) H + （ aq ） +A - （ aq ） Weak acids:

P584 on the textbook: Table 15.3 Ionization Constants of some Common Weak Acids at 25 ℃

K a Increases Strength vs. K a

Calculating the pH of Weak Acid Solutions 1.0M Before 0 0 Equilibrium (1-x)MxMxMxMxM C/K<500

The 5% Rule The validity of an approximation should always be checked. The 5% rule should be used to evaluate which approximation are reasonable. The approximation is acceptable.

The pH of a Mixture of Weak Acids HNO 2 is assumed to be the dominant producer of H +.

x xx

Percent Dissociation xxx

15.7 The Common Ion Effect Common ion effect: The shift in equilibrium position that occurs because of the addition of an ion already involved in the equilibrium reaction.

The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The presence of a common ion suppresses the ionization of a weak acid or a weak base. Consider mixture of CH 3 COONa (strong electrolyte) and CH 3 COOH (weak acid). CH 3 COONa (s) Na + (aq) + CH 3 COO - (aq) CH 3 COOH (aq) H + (aq) + CH 3 COO - (aq) common ion 16.2

A buffer solution is a solution of: 1.A weak acid or a weak base and 2.The salt of the weak acid or weak base Both must be present! A buffer solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base CH 3 COOH (aq) H + (aq) + CH 3 COO - (aq) Consider an equal molar mixture of CH 3 COOH and CH 3 COONa Adding more acid creates a shift left IF enough acetate ions are present

15.8 Buffer Solutions A Buffer Solution (or buffered solution) is one that resists a change in pH when either hydroxide ions or protons are added. A buffered solution may contain a weak acid and its salt or a weak base and its salt. (HF+NaF, NH 3 +NH 4 Cl)

How Do the H + /OH - Ions Work in Buffered Solutions The equilibrium concentration of H + and the pH are determined by the ratio [HA]/[A - ].

The Effect of Added Bases When OH - are added, HA is converted to A -, causing the ratio [HA]/[A - ] to decrease. If the amount of HA and A - originally present are very large compared with the amount OH - added, the change in [HA]/[A - ] ratio is small.

The Effect of Added Acids When protons are added to a buffered solution, the added H + ions react with A - to form the weak acid. If [HA] and [A - ] are large compared with the [H + ] added, only a slight change in the pH occurs.

Unique Properties of Buffer Solution The Effect of Dilution － The pH of a buffer solution remains essentially independent of dilution. The Effect of Added Acids and Bases － A buffer solution resists pH change after addition of small amounts of strong acids or bases Buffer Capacity

Which of the following are buffer systems? (a) KF/HF (b) KCl/HCl, (c) Na 2 CO 3 /NaHCO 3 (a) KF is a weak acid and F - is its conjugate base buffer solution (b) HCl is a strong acid not a buffer solution (c) CO 3 2- is a weak base and HCO 3 - is it conjugate acid buffer solution 16.3

What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? HCOOH (aq) H + (aq) + HCOO - (aq) Initial (M) Change (M) Equilibrium (M) x-x+x+x x x+x x x 16.2 Mixture of weak acid and conjugate base! K a for HCOOH = 1.8 x [H + ] [HCOO - ] K a = [HCOOH] x = X pH = 3.98

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15.9 How Acid-Base Indicators Work Add a few drops of the phenolphthalein indicator to a acidic solution. (pH=1) The ratio shows that the predominant form of the indicator is HIn, resulting in a red solution.

As OH - is added to this solution, [H + ] decreases and the equilibrium shift to right, changing HIn to In -. A color change from red to reddish purple will occur. For most indicators, about 1/10 of the initial form must be converted to the other form before a new color is apparent.

Indicators