Chapter 17
• Acids are substances that increase the concentration of hydrogen
ions in solution.
• Bases are substances that increases the concentration of hydroxide
ions in solution.
• Acids and bases can be either strong or weak electrolytes.
• The proton produced by acids combines with water to form
hydronium ion, H3O+.
• A strong acid/base completely ionizes in solution. (NaOH, HCl)
• Weak electrolytes ionize only to a small extent. (a large amount of
acid remains in solution – HC2H3O2)
• Equilibrium constants for weak acids are called ionization constants.
• Acids or bases that ionize extensively, with K>1, are referred to as
strong acids or bases.
• Acids or bases that do not ionize extensively, with K<1, are referred to
as weak acids or bases.
• Bronsted and Lowry said that an acid is a proton donor, such as nitric
acid or ammonium.
• A Bronsted base is a proton acceptor such as ammonia.
• Monoprotic acids donate one proton (HF, HCl, HNO3) while
polyprotic acids are capable of donating two or more (H2SO4).
• Polyprotic bases may accept more than one proton, such as sulfate or
phosphate ion.
• Molecules that can act as either a Bronsted acid or base are called
amphiprotic, such as water!
• Write a balanced equation for the reaction that occurs
when H3PO4, phosphoric acid, donates a proton to water to
form the dihydrogen phosphate ion. Is the dihydrogen
phosphate ion an acid, a base, or amphiprotic?
• Write a balanced equation for the reaction that occurs
when the cyanide ion, CN-, accepts a proton from water to
form HCN. Is CN- a Bronsted acid or base?
• A pair of compounds or ions that differ by the presence of one H+ ion
is called a conjugate acid-base pair.
• Every reaction between a Bronsted acid and Bronsted base involves
H+ transfer and has two conjugate acid-base pairs.
• In the following reaction, identify the acid on the left and its
conjugate base on with right. Similarly, identify the base on the left
and its conjugate acid on the right.
HNO3 + NH3  NH4+ + NO3-
• Two water molecules interact with each other to produce a
hydronium ion and a hydroxide ion.
• This property is called autoionization. Water will still conduct a
small amount of electricity since it contains low concentration of
hydronium and hydroxide.
• Kw = [H3O+][OH-] = 1.0 x 10-14 at 25 oC
• In pure water the two ion concentrations are equal and the water is
said to be neutral.
• Adding acid increases the hydronium concentration, adding base
increases the hydroxide. Both additions would disturb the
equilibrium.
• A solution of the strong acid HCl has [HCl] = 4.7 x 10-3 M. What are
the concentrations of H3O+ and OH- in this solution? (remember
that HCl is 100% ionized in water)
• pH = -log[H3O+]
• pOH = -log[OH-]
• pKw = 14.00 = pH + pOH
• Approximate pH of solutions can be determined using an acid-base
indicator.
• Substances that change color in a known pH range are Bronsted
acids or bases for which the acid and its conjugate base have
different colors.
• Modern pH meters are preferable!
• What is the pH of a 0.0012 M NaOH solution?
• The pH of a diet soda is 4.32 at 25oC. What are the hydronium and
hydroxide ion concentrations in the soda?
• If the pH of a solution of the strong base Sr(OH)2 is 10.46, what is the
concentration of the Sr(OH)2 in mol/L?
• After reading sections 17.1-17.3, you should be able to do the
following…
• P. 801 (1-14)
• The lower the pH the stronger the acid.
• For a strong base/acid, the [OH-]/[H3O+] concentration is equal to
the original base/acid concentration.
• For a weak base/acid, the ion concentration is much less than the
original acid concentration.
• For the general acid HA, we can write
Ka = [H3O+][A-]/[HA]
where the Ka is the equilibrium constant for an acid in water.
• K is less than 1 for weak acids. Ka increases as acid strength increases.
• For a weak base B in water…
Kb = [BH+][OH-]/[B]
Ionization Constants for Some Acids and their Conjugate Bases are on
page 770.
The weaker the acid, the stronger its conjugate base.
• Which is the stronger acid, H2SO4 or H2SO3?
• Is benzoic acid, C6H5CO2H, stronger or weaker than acetic
acid?
• Which has the stronger conjugate base, acetic acid or boric
acid?
• Which is the stronger base, ammonia or the acetate ion?
• Which has the stronger conjugate acid, ammonia or the
acetate ion?
• Anions that are conjugate bases of strong acids are such
weak bases that they have no effect on solution pH.
• There are numerous basic anions; all are the conjugate bases
of weak acids.
• Acidic anions arise from polyprotic acids, and are
amphiprotic as well.
• Alkali metal and alkaline earth cations have no measurable
effect on solution pH.
• All metal cations are hydrated in water. Only when the
cation is a +2 or +3 does the ion act as an acid.
• For each of the following salts in water, predict whether the pH will
be greater than, less than, or equal to 7.
• KBr
• NH4NO3
• AlCl3
• Na2HPO4
• The negative log of the Ka value (pKa) becomes smaller as the acid
strength increases.
pKa = -logKa
• What is the pKa value for benzoic acid, C6H5CO2H?
• Is chloroacetic acid (ClCH2CO2H), pKa = 2.87, a stronger or weaker
acid than benzoic acid?
• What is the pKa for the conjugate acid of ammonia? Is this acid
stronger or weaker than acetic acid?
• As Ka decreases, Kb increases. The product of the two is equal to the
autoionization constant for water.
• Kw = KaKb
• Ka for lactic acid, CH3CHOHCO2H, is 1.4 x 10-4. What is Kb for the
conjugate base of this acid, CH3CHOHCO2-? Where does this base
fit in Table 17.3?
• All proton transfer reactions proceed from the stronger acid and base
to the weaker acid and base. (Equilibrium favors the weaker acid
and base.)
• Which is the stronger Bronsted acid, HCO3- or NH4+? Which has
the stronger conjugate base?
• Is a reaction between HCO3- ions and NH3 product- or reactantfavored?
HCO3-(aq) + NH3(aq)  CO32-(aq) + NH4+(aq)
• Write the net ionic equation for the possible reaction between acetic
acid and sodium hydrogen sulfate, NaHSO4. Does the equilibrium
lie to the left or right?
• After reading sections 17.4-17.5, you should be able to do the
following…
• P. 801 (15-17,23-25,30-31,34-35)
• Strong Acid with a Strong Base
• mixing equal quantities will produce a neutral solution
• Weak Acid with a Strong Base
• mixing equal quantities produces a basic solution; pH depends on Kb for anion
produced
• Strong Acid with a Weak Base
• mixing equal quantities produces an acidic solution; pH depends on Ka for the
cation produced
• Weak Acid with a Weak Base
• mixing equal quantities produces a solution in which the pH depends on the Ka
and Kb for the cations and anions produced
• Equal molar quantities of HCl and NaCN are mixed. Is the resulting
solution acidic, basic, or neutral?
• Equal molar quantities of acetic acid and sodium sulfite, Na2SO3, are
mixed. Is the resulting solution acidic, basic, or neutral?
• You can use an ICE table to calculate K from initial concentrations
and measured pH.
• A solution prepared from 0.055 mol of butanoic acid dissolved in
sufficient water to give 1.0L of solution has a pH of 2.72. Determine
Ka for butanoic acid. The acid ionizes according to the balanced
equation
CH3CH2CH2CO2H + H2O  H3O+ + CH3CH2CH2CO2-
• Due to the fact that very little ionization occurs in a weak acid, we
can assume that the acid concentration at equilibrium is basically
the same as the initial acid concentration.
• This is valid whenever [HA]0 is greater than or equal to 100Ka.
• We can use ICE tables to calculate the pH of a solution of a weak
acid or base and using known equilibrium constants.
• What are the equilibrium concentrations of acetic acid, acetate ion,
and H3O+ for a 0.10M solution of acetic acid (Ka = 1.8x10-5)? What is
the pH of the solution?
• What are the equilibrium concentrations of HF, fluoride ion, and
H3O+ when a 0.015M solution of HF is allowed to come to
equilibrium? What is the pH of the solution?
• Sodium hypochlorite, NaOCl, is used as a disinfectant in swimming
pools and water treatment plants. What are the concentrations of
HOCl and OH- and the pH of a 0.015M solution of NaOCl?
• In order to calculate pH in a resulting solution, you must write a
balanced equation and decide whether the products are acid or
basic.
• You must then find initial concentrations and can calculate pH by
solving an equilibrium problem.
• Calculate the pH after mixing 15mL of 0.12M acetic acid with 15mL
of 0.12M NaOH. What are the major species in solution at
equilibrium (besides water) and what are their concentrations?
• After reading 17.6 and 17.7, you should be able to do the following…
• P. 802 (38-39,42-45,50-55,60-62)
• Acids that can donate more than one proton are polyprotic.
• Each step has its own Ka, which becomes progressively smaller due to
the increased energy required to remove a proton.
• The pH of many polyprotic acids depends primarily on the
hydronium ion generated in the first ionization step.
• What is the pH of a 0.10M solution of oxalic acid, H2C2O4? What
are the concentrations of H3O+, HC2O4-, and the oxalate ion, C2O42?
• This concept is based on the sharing of electron pairs between acid
and base.
• A Lewis acid is a substance that can accept a pair of electrons from
another atom to form a new bond.
• A Lewis base is a substance that can donate a pair of electrons to
another atom to form a new bond.
• The product of an acid-base reaction in the Lewis sense is often
called an acid-base adduct. This type of bond is called a coordinate
covalent bond.
• The formation of hydronium and ammonium are both examples.
• All metal cations form hydrated cations in which the metal
ion is surrounded by water molecules, such a [Fe(H2O)6]2+.
• These structures are called complex ions, or coordination
complexes.
• Hydroxide ion is a Lewis base and binds readily to metal
cations to form metal hydroxides. Metal hydroxides are
usually amphoteric.
• Acidic oxides such as carbon dioxide and sulfur dioxide.
• Due to oxygen’s high electronegativity, electrons are polarized away
from the other element which can therefore react with a Lewis base
such as hydroxide ion.
• Describe each of the following as a Lewis acid or a Lewis base. (Draw
the Lewis dot structure. Are there lone pairs on the central atom? If
so, it may be a Lewis base. Does the central atom lack an electron
pair? If so it can behave as a Lewis acid.
• PH3, BCl3, H2S, HS-
• Stronger acids have weak H-X bonds (such as HCl) and weaker acids
have strong H-X bonds (such as HF).
• Oxoacids, such as HNO3, contain an atom bonded to one or more
oxygen atoms, some with hydrogen atoms attached.
• Inductive effect: the attraction of electrons from adjacent bonds by more
electronegative atoms
• Carboxylic acids, hydrated metal cations, and anions all act as
Bronsted bases.
• Organic amines, such as ammonia compounds, act as Bronsted and
Lewis bases.
• Which is the stronger acid, H2SeO3 or H2SeO4?
• Which is the stronger acid, Fe(H2O)62+ or Fe(H2O)63+?
• Which is the stronger acid, HOCl or HOBr?
• After reading sections 17.8-17.10, you should be able to do the
following…
• P. 804 (65-68,69-71,73-75)