Arrhenius, Bronsted-Lowry, and Lewis Acids/Bases
Neutralization reactions. Neutralization reactions are reactions between acids and bases and, often, they take
the form:
Acid + Base  Salt + Water
The three theories of identifying a compound as being an acid or base are:
1. Arrhenius Acids/Bases
2. Bronsted – Lowry Acids/Bases
3. Lewis Acids/Bases
Arrhenius Acids/Bases
Arrhenius’ theory of acids/bases focuses on what molecules contain. An Arrhenius acid contains a H+. This
can be compounds like HCl, HNO3, HI, etc. All of these compounds contain a hydrogen that we would
characterize as being acidic. An Arrhenius base contains an OH-. This can be compounds like NaOH, KOH,
CsOH, etc. All of these have a hydroxide group that makes them basic.
Bronsted – Lowry Acids/Bases
When Arrhenius’ definition covers many acids/bases, acid/base reactions are not about containing H+ or OH-.
Bronsted and Lowey modified Arrhenius’ theory with their own additions. A Bronsted-Lowry acid is any
compound that gives H+ during the reaction. A Bronsted-Lowry base is any compound that accepts H+
during the reaction. Therefore, the focus is not on what a molecule contains, but the movement of the acidic
hydrogen such as the following reaction:
focuses on movement of H+
Lewis Acids/Bases
A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor. This is the broadest
definition because it also includes all acids/bases that are in the Arrhenius definition.
Lewis Acid
Lewis Base
Conjugate Acid/Base Pairs
As with any reaction, there is a forward and reverse reaction for acid/base reactions. Consider the following
example:
HX + H2O
H3O+ + X+
In the forward reaction HX is the acid because it donates H and H2O is the base because it accepts H+. In the
reverse reaction, H3O+ is now the acid and X- is now the base. The pair of HX and X- are known as an
conjugate acid-base pair.
Every acid has a conjugate base that is formed when a proton (H+) is removed. Every base has a conjugate
acid that forms when a proton is added
Question 1:
In the following reactions, identify the Acid, Base, Conjugate acid, conjugate base
a) H2SO4 + NaOH  NaHSO4 + H2O
b) NH3 + CH3COOH  NH4+ + CH3COOc) H3PO4 + NH3  H2PO4- + NH4+
d) AlCl3 + Cl-  AlCl4e) H2SO4 + Ca(OH)2  H2O + Ca(HSO4)2
Strong vs. Weak Acids/Bases
Acids and bases can be classified as strong or
weak based on how much it disassociates. Recall
that dissociate means to break apart. Acids/Bases
that completely dissociate are called strong
acids/bases. When we examine a solution of
strong acid, we only see the products, we do not
see any reactants.
Acids/Bases that only partially dissociate are
called weak acids/bases. These are in
equilibrium with one another, so we see both
reactants and products present throughout the entire solution.
There are only 7 strong acids and 8 strong bases. You must memorize these. Everything that is not classified
as “strong” on the list below is considered to be a weak acid and in equilibrium with the reactants.
7 Strong Acids
1. HBr
2. HCl
3. HI
4. HNO3
5. H2SO4
6. HClO3
7. HClO4
7 Strong Bases
1. LiOH
2. NaOH
3. KOH
4. RbOH
5. CsOH
6. Ca(OH)2
7. Sr(OH)2
8. Ba(OH)2
Question2: Compare and contrast a strong acid and a weak acid.
pH, pOH, and Calculating pH/pOH
pH scale is used to convey the concentration of H+. The pH scale ranges from 0-14.
pH Between 0-7 = Acidic
pH Between 7-14 = Basic or Alkaline
pH = 7
pH + pOH = 14
pH = - log([H+])
pOH =
- log([OH-])
Recall, for strong acids and bases, the solutions are completely dissociated. This means that if I have a solution
of HCl, it is completely broken into H+ and Cl-. This means that the concentration of a strong acid/base is
the [H+] or [OH-].
For bases, we only know [OH-]. We do not know the [H+], so we cannot calculate the pH in one step. We need
to first calculate pOH and then convert to pH using the equation:
pH = 14 – pOH
Question 3: Identify the following solutions as acidic, basic, or neutral:
a. pH = 6.7 ____________
b. pH = 8.23 ___________
c. pH = 5.23 ___________
Question4: You have a 0.231 M solution of HBr. What is the pH of the solution?
Question5:You have a 0.512 M solution of HNO3. What is the pH of the solution?
Question6: You have a 0.7801 M solution of KOH. What is the pH of the solution?
Question7: You make a solution by dissolving 4.2 moles of NaOH in 7 L of water. What is the
pH of this solution?
Amphoteric - can behave either as an acid or as a base. ex. - water
Autoionization - involves the transfer of a proton from one water molecule to another to
produce a hydroxide ion and a hydronium ion.
2H2O(l)→ H3O+(aq) + OH-(aq)
Ion-Product Constant Kw (or dissociation constant for water) - always refers to the
autoionization of water. Using the above reaction for the autoionization of water:
2H2O(l)→ H3O+(aq) + OH-(aq)
Write equilibrium constant expression for this reaction???
Kw
or
= [H3O+][OH-]
= [H+]. [OH-]
[H+] = [OH-] = 1.0 x 10-7
Kw = (1.0 x 10-7 M)2 = 1.0 x 10-14 mol2/L2
It is important to recognize the meaning of Kw. The product of [H+] and [OH-] must always
equal 1x 10-14. This lends itself to three possibilities:
A) A neutral solution, where [H+] = [OH-]
B) An acidic solution, where [H+] > [OH-].
C) A basic solution, where [H+] < [OH-].
Calculating the pH of Weak Acid Solution
Step 1 - List the major species in the solution. +
Step 2 - Choose the species that can produce H , and write a balanced equation for
the reaction producing H+.
Step 3 - Using the values of the equilibrium constant for the reactions you+ have
written, decide which equilibrium will dominate in producing H .
Step 4 - Write the equilibrium expression for the dominant equilibrium.
Step 5 USE ICE TABLES using x for change in concentrations
Step 6 Solve for x the "easy" way: that is by assuming
[HA]0 - x  [ HA]0
Step 7 - Using the 5% rule verify whether the approximation is valid.
Step 8 - Calculate [H+] and pH
5% Rule: If x is <5% of the initial concentration, the assumption is probably a good one to
ignore that x from subtraction.
- If x>5% of the initial concentration, then it may be best to solve the quadratic equation or use
successive approximations. In that case, you will use two values, use the value of x smaller than
the number.
Percent Dissociation:
For a given weak acid, the percent dissociation
increases as the acid becomes more
dilute. For solutions of any weak acid HA, [H+] decreases as [HA]0 decreases, but
the percent dissociation increases as [HA]0 decreases.
Bases
Strong Base - one which dissociates completely in water and has large Kb values.
General reaction for a base B and water is
given by:
+ H2O(l)
B(aq)
Acid
Base
BH+(aq)
Conjugate
Acid
+
OH-(aq)
Conjugate
Base
The equilibrium constant for this
equation:
Question8: A student prepared a 0.10 M solution of formic acid (HCHO2) and measured its
pH. The pH at 25oC was found to be 2.38. Calculate the Ka for formic acid at this temperature.
Question9:
Niacin, one of the B vitamins, has the following molecular structure:
A 0.020 M solution of niacin has a pH of 3.26.
What is the acid-dissociation constant, Ka, for niacin?
Question10: A 0.10 M solution of formic acid (HCOOH) contains 4.2 x 10-3 M H+(aq).
Calculate the percentage of the acid that is ionized.
Question11:A 0.020 M solution of niacin has a pH of 3.26. Calculate the percent
ionization of the niacin
Question12:Calculate the pH of a 0.20 M solution of HCN. Refer to the table for Ka
Question13:The Ka for niacin is 1.6 x 10-5. What is the pH of a 0.010 M solution of
niacin?
Question14: Calculate the percentage of HF molecules ionized in
a) a 0.10 M HF solution
b) a 0.010 M HF solution
Question15:
The percent ionization of niacin (Ka = 1.5 x 10-5) in a 0.020 M solution is 2.7%.
Calculate the percentage of niacin molecules ionized in a solution that is
a) 0.010 M
b) a 1.0 x 10-3 M
Question16:
Anions of weak acids are also weak bases. Example: ClO– is the conjugate base of HClO (weak
acid):
ClO–(aq) + H2O(l) → HClO(aq) + OH– (aq)
Kb = 3.3 x 10–7
A solution made by adding solid sodium hypochlorite (NaClO) to enough water to make 2.00 L
of solution has a pH of 10.50. How many moles of NaClO were added to the water?
Hint:
Question17:A solution of NH3 in water has a pH of 11.17. What is the molarity of the solution?
Use the following table for Ka / Kb to answer:
Question18: Calculate
a) the base-dissociation constant, Kb, for the fluoride ion (F-);
b) the acid-dissociation constant, Ka, for the ammonium ion (NH4+).
c) Which of the following anions has the largest base-dissociation constant: NO2-, PO43-, or N3-?
Question19:
The base quinoline has the following structure:
Its conjugate acid is listed in handbooks as having a pKa of 4.90.
What is the base-dissociation constant for quinoline?
Question20:
The hydrogen sulfite ion (HSO3-) is amphoteric.
a) Write an equation for the reaction of HSO3- with water, in which the ion acts as an acid.
b) Write an equation for the reaction of HSO3- with water, in which the ion acts as a base.
In both cases identify the conjugate acid-base pairs.
Question21:
When lithium oxide (Li2O) is dissolved in water, the solution turns basic from the reaction of
the oxide ion (O2-) with water.
Write the reaction that occurs, and identify the conjugate acid-base pairs.
Question22:
For each of the following reactions, use the above figure to predict whether the equilibrium lies
predominantly
to the left or to the right:
a) HSO4(aq) + CO32-(aq) →SO42-(aq) + HCO3-(aq)
b) HPO42-(aq) + H2O(l)→ H2PO4(aq) + OH-(aq)
c) NH4+(aq) + OH-(aq) → NH3(aq) + H2O(l)
Question23:
Indicate whether solutions with each of the following ion concentrations is neutral, acidic, or
basic:
a) [H+] = 4 x 10-9 M
b) [OH-] = 1 x 10-7 M
c) [OH-] = 7 x 10-13 M
d) Calculate the concentration of OH- in a 0.15M solution of NH3.
Acid-Base Properties of Salt Solutions
• Acid-base properties of salts are a consequence of the reactions of their ions in solution.
An Anion’s Ability to React with Water
• Anions from weak acids are basic: They will cause an increase in pH.
• Anions from strong acids are neutral. They do not cause a change in pH.
• Anions with ionizable protons (e.g., HSO4– ) are amphiprotic.
• They are capable of acting as an acid or a base.
• If Ka > Kb, the anion tends to decrease the pH.
• If Kb > Ka, the anion tends to increase the pH.
A Cation’s Ability to React with Water
• Polyatomic cations that have 1 or more ionizable protons are conjugate acids of weak bases �
decrease pH.
• Metal cations of Group 1A and heavy alkaline earth metals are cations of strong bases and do
not alter pH.
• Other metal ions can cause a decrease in pH.
Combined Effect of Cation and Anion in Solution
• The pH of a solution may be qualitatively predicted using the following guidelines:
• Salts derived from a strong acid and a strong base are neutral.
• Examples are NaCl and Ca(NO3)2.
• Salts derived from a strong base and a weak acid are basic.
• Examples are NaClO and Ba(C2H3O2)2.
• Salts derived from a weak base and a strong acid are acidic.
• An example is NH4Cl.
• Salts derived from a weak acid and a weak base can be either acidic or basic.
• Equilibrium rules apply!
• We need to compare Ka and Kb for hydrolysis of the anion and the cation.
• For example, consider NH4CN.
• Both ions undergo significant hydrolysis.
• Is the salt solution acidic or basic?
• • The Ka of NH4+
+ is smaller than the Kb of CN–, so the solution should be basic.
Question 24) Determine whether aqueous solutions of each of the following salts will be acidic,
basic, or neutral:
a) Ba(CH3COO)2
b) NH4Cl
c) CH3NH3Br
d) KNO3
e) Al(ClO4)3
Question 25)
In each of the following, circle which salt in each of the following pairs will form the more
acidic (or less basic) 0.010 M solution:
a) NaNO3 or Fe(NO3)3
b) KBr, or KBrO
c) CH3NH3Cl or BaCl2
d) NH4NO2 or NH4NO3
Question26:
Predict whether the salt Na2HPO4 will form an acidic or a basic solution on dissolving in water.
Question27:
Predict whether the dipotassium salt of citric acid (K2HC6H5O7) will form an acidic or basic
solution in water.
Factors That Affect Acid Strength
• The H–X bond must be polar with Hδ+ and Xδ-.
• The strength of the bond: the H–X bond must be weak enough to be broken.
• The stability of the conjugate base, X–: the greater the stability of the conjugate base, the more
acidic the molecule.
Binary Acids
• Group trends:Acid strength increases down a group; base strength decreases down a group.
• Period trends: Acid strength increases and base strength decreases right across a period as the
electronegativity of X increases.
Oxyacids Acids that contain OH groups (and often additional oxygen atoms) bound to the
central atom
• All oxyacids have the general structure Y–O–H, e.g. H2SO4
• The strength of the acid depends on Y and the atoms attached to Y.
• As the electronegativity of Y increases, so does the acidity of the substance.
• For oxyacids with the same number of OH groups and the same number of oxygen atoms:
• Acid strength increases with increasing electronegativity of the central atom, Y.
• Example: HClO > HBrO > HIO
• Acid strength increases as the number of oxygen atoms attached to Y increases.
• Example: HClO4 > HClO3 > HClO2 > HClO
Question28:
Which is the most acidic?
H2SO4, H2SeO3, H2SeO4
Question29:
In each of the following pairs choose the compound that leads to the more acidic (or less basic) solution:
a) HBr, HF;
b) PH3, H2S;
c) HNO2, HNO3;
d) H2SO3, H2SeO3.