Chem 40S Unit 5 Notes

Chemistry 40S Unit 4 – Acids & Bases

Acids & Bases


Acids & Bases
Properties of Acidic Solutions
• Taste sour. For example, lemons (citric acid) and vinegar
(acetic acid).
• Burn when touching skin.
• Turn blue litmus red.
• Neutralize basic solutions.
• React with carbonates to produce carbon dioxide gas.
For example, when you add vinegar to baking soda
(sodium bicarbonate), fizzing occurs. This fizzing is the
production of carbon dioxide.
• Corrosive to metals. Many acids react with active metals
to produce hydrogen gas.
• Another property of acids is that they are electrolytes


Acids & Bases
Properties of Basic Solutions
• Taste bitter.
• Feel slippery.
• Turn red litmus blue.
• Neutralize acidic solutions.
• Bases are electrolytes


Acids & Bases

Arrhenius Definition
vs
Bronstead-Lowry Definition


Arrhenius Definition


Acids & Bases
Arrhenius Definition of Acids & Bases
• Arrhenius noticed that acidic and basic solutions
were electrolytes. He determined that acids and
bases must ionize or dissociate in water.
• According to Arrhenius, an acid is defined as a
substance which releases hydrogen ions in water.

Example: hydrochloric acid:
HCl(aq) → H+(aq) + Cl¯(aq)


Acids & Bases
Arrhenius Definition of Acids & Bases
• According to Arrhenius, a base is defined as a
substance which releases hydroxide ions in water.

Example 3, sodium hydroxide:

NaOH(s) → Na+(aq) + OH¯(aq)


Acids & Bases
Problems with Arrhenius Definition of
Acids & Bases
• The Arrhenius definition says that acids and bases
can only occur in water solutions – not true.
• There are many substances which are acidic or
basic but do not have a hydrogen ion or a
hydroxide ion.
• The Arrhenius Theory is not a complete loss,
however, since it was important in establishing
the concept of dissociation.


Bronstead-Lowry Definition


Acids & Bases
Brønsted-Lowry Definition of Acids & Bases
• In 1923 a Danish chemist named Johannes
Brønsted (1879-1947) and an English chemist
Thomas Lowry (1874-1936) independently
developed a more general definition of acids and
bases within months of each other.
• According to the Brønsted-Lowry definition, an
acid is a proton or H+ ion donor. A base is
defined as a proton acceptor.


Acids & Bases
• In this example, hydrogen chloride
(hydrochloric acid) reacts with water by
donating a proton. Water acts as the base,
accepting the proton. The result is the H3O+
ion called the hydronium ion.


Acids & Bases
• In this example, ammonia's properties as a base are
better explained by the Brønsted-Lowry definition.
Ammonia accepts a proton from water, making
ammonia a base and water the acid. The result is the
ammonium ion and the hydroxide ion.


Acids & Bases
• In the Brønsted-Lowry definition of acids and bases,
substances like water can act as BOTH an acid and a
base. These types of substances are called
amphoteric (the root "amph" is similar to the root of
amphibian which means "having two lives", on land
and in water).


Acids & Bases
Conjugate Acids and Bases
• The general form of a Brønsted-Lowry acid-base
reaction is
Acid + Base → Conjugate Acid + Conjugate Base
• The conjugate acid is what remains after a base
has accepted a proton and the conjugate base is
what remains after the acid has donated its proton.


Acids & Bases
• Let's take another look at the reaction of
ammonia with water.
NH3(g) + H2O(l) → NH4+(aq) + OH¯(aq)

Base Acid

• The reverse reaction would be
NH4+(aq) + OH¯(aq) → NH3(g) + H2O(l)

Acid Base →


Acids & Bases
• Notice an acid results form the ammonia
accepting a proton from water. The ammonium
ion can donate a proton to the hydroxide ion. The
hydroxide ion accepts the proton making it a base.
• In the first reaction, ammonia is the base and the
ammonium ion is its conjugate acid. Water is the
acid in the first reaction and the hydroxide ion is
its conjugate base.
• NH3(g) and NH4+(aq) are called a conjugate acid-
base pair, as are H2O(l) and OH¯(aq).


Acids & Bases
• So, for this reversible reaction:

• In general for the acid HA, and for the base


Acids & Bases
Hydrolysis
• The Brønsted-Lowry definition can explain
why a solution of sodium hydrogen
carbonate (baking soda) is basic and can
react with either acidic or basic solutions.
• In water, NaHCO3 produces two ions:

NaHCO3(s) → Na+(aq) + HCO3¯(aq)


Acids & Bases
Advantages of The Brønsted-Lowry Theory
• The Arrhenius Theory of acids and bases was
limited to aqueous solutions. The Brønsted-Lowry
Theory expands the definition of an acid and base
to a proton donor or acceptor. This means the acid
or base can be in any state.
• The Brønsted-Lowry Theory is able to explain
why substances such as the hydrogen carbonate
ion can act as an acid and a base.


Acids & Bases
Advantages of The Brønsted-Lowry Theory

• All Arrhenius acids are Brønsted-Lowry acids,
but NOT all Brønsted-Lowry acids are
Arrhenius acids.

• All Arrhenius bases are Brønsted-Lowry bases,
but NOT all Brønsted-Lowry bases are
Arrhenius bases.


Acids & Bases

Your task is to
write an
analogy
describing
what a
Bronstead-
Lowry acid or
Base is.


Strong & Weak Acids & Bases


Strong Acids
• Strong acids are substances that easily
donate protons. This means they must easily
dissociate or ionize in water. The bond
holding the proton to the rest of the
molecule is not very strong.
• Strong acids are usually strong
electrolytes.
• Strong acids are considered to
COMPLETELY dissociate in water.


Strong Acids
• For example,
HCl(g) + H2O(l) → H3O+(aq) + Cl¯(aq)
• The following are examples of strong acids.
You must memorize the names and
chemical formulas of these acids!
Chemical Formula Name of Acid
HClO4 perchloric acid
HCl hydrochloric acid
H2SO4 sulfuric acid
HNO3 nitric acid
HBr hydrobromic acid


Weak Acids
• Weak acids are poor or weak electrolytes. They
do not conduct an electric current in water very
well.
• Weak acids are poor electrolytes because they
ionize incompletely.
• Weak acids have a high affinity for their
proton(s). The hydrogens are bound quite
strongly to the rest of the molecule, so it takes a
very strong base to “rip” the proton from a weak
acid.


Weak Acids
• An example of a weak acid is acetic acid. Only a
small number of aqueous acetic acid molecules
react with water to form an acetate ion and a
hydronium ion. The rest exist as acetic acid
molecules.
• This establishes an equilibrium between the ions
and the molecules of acetic acid. We show this
using the double arrows:
HC2H3O2(aq) + H2O(l) ↔ H3O+(aq) + C2H3O2¯(aq)


Weak Acids
• Here are more examples of weak acids.
Their names and formulas should be
memorized!
Chemical Formula Name of Acid
H2CO3 carbonic acid
HNO2 nitrous acid
HF hydrofluoric acid


Strong Bases
• Strong bases are substances have a very high
affinity for protons, H+ ions. Strong bases are
strong electrolytes, that is, they completely
dissociate when dissolved in water.
• The most common strong bases are sodium
hydroxide, NaOH, and potassium hydroxide,
KOH.
NaOH(s) → Na+(aq) + OH¯(aq)
KOH(s) → K+(aq) + OH¯(aq)


Strong Bases
• Other strong bases are
Chemical Formula Name of Base
Mg(OH)2 Magnesium hydroxide
Ca(OH)2 Calcium hydroxide
CaO Calcium oxide (lime)


Weak Bases
• Just as with weak acids, weak bases are
poor proton acceptors. These do not have a
large affinity for protons.
• For example, ammonia is a weak base. Its
dissociation should be memorized.
NH3(g) + H2O(l) ↔ NH4+(aq) + OH¯(aq)


Weak Bases
• Other examples of molecular compounds
which exist as bases are
Chemical Formula Name of Base
CH3NH2 methylamine
C6H5NH2 aniline
C5H5N pyridine


Weak Bases
• The most common weak bases are the
conjugate bases of strong acids.
• For example, the carbonate ion:
CO32¯(aq) + H2O(l) ↔
HCO3¯(aq) + OH¯(aq)


Acid Strength Table


Using the Acid Strength Chart
Example 1
• Arrange the following in decreasing
strength as an acid and decreasing strength
as bases. Note some may be used in both
groups.
HNO2, OH¯, HCO3¯, HPO42¯, HTe¯, C2H3O2¯


Acid-Base Reactions
• Recall, an acid donates a proton and a base
accepts a proton.

• The stronger acid of the two tends to donate its
proton, so the reaction favours the direction
AWAY from the strongest acid. For example, if
Acid1 was the stronger acid, the forward reaction
would be favoured, away from the stronger acid.


Acid-Base Reactions
Example 2
• Complete the reaction below, indicating the
acids and bases. which direction is favoured
and why?
• HCO3¯ + PO43¯ ↔


The Equilibrium Law


Acid Dissociation Constant
• Strong acids dissociate completely, and
therefore do not establish an equilibrium.
However, weak bases do establish equilibrium.
This means, for weak acids we can make an
equilibrium law.
• In general, for the weak acid HA

HA + H2O ↔ H3O+ + A¯


• The equilibrium law would be

• but water is a liquid and its concentration does not
change, so we remove it and replace Kc with Ka,
the acid dissociation constant or ionization
constant.

• The acid dissociation constant reflects the
equilibrium that exists for an acid in solution.


• The larger the Ka, the more product, so the greater the
dissociation. The larger the Ka, the stronger the acid.
• For example, hydrochloric acid completely dissociates
according to the equation
HCl(g) + H2O(l) ↔ H3O+(aq) + Cl¯(aq)
• For a 1.0 mol/L solution, H3O+ and Cl¯
will both be 1.0 mol/L and there will
be no HCl.

• The smaller the Ka, the less product, so the weaker the acid.
For example, the weak acid, acetic acid has a Ka equal to 1.8 x
10-5.


Base Dissociation Constant
• Just as with acids, the base dissociation constant, Kb,
reflects the strength of a base. The higher the value of
Kb, the stronger the base.
• For example, the weak base ammonia ionizes
according to the equation

NH3(g) + H2O(l) ↔ NH4+(aq) + OH¯(aq)

• The equilibrium law is


• The Kb for ammonia at 25°C is 1.8 x 10-5. In general,
the Kb for the weak base, B, whose dissociation is like
that of ammonia is

B + H2O(l) ↔ BH+(aq) + OH¯(aq)


• In general, for the weak base BOH, the equilibrium
law would be
BOH(aq) ↔ B+(aq) + OH¯(aq)

• The strong base, sodium hydroxide, dissociates
completely according to the equation
NaOH(s) ↔ Na+(aq) + OH¯(aq)
• Since no NaOH is present in solution, the Kb is very
large.


Calculating the Dissociation Constant
Example 1
• A 0.10 mol/L solution of acetic acid is only partly
ionized. If at equilibrium, the hydronium ion
concentration is 1.3 x 10-3 mol/L, what is the acid
dissociation constant, Ka?


Calculating the Concentration of
Dissociated Species
• If we know the acid or base dissociation constant, we
can calculate all the species in the solution. This
requires the use of the “ICE” table.

Example 2
• HA is a weak acid with a Ka of 7.3 x 10-8. What is the
concentration of all species (HA, H3O+ and A¯) if the
initial concentration of HA is 0.50 mol/L?


Percent Dissociation
• Acids and bases can be described in terms of strong
or weak, concentration and degree of dissociation.
The acid and base dissociation constants represent the
acid's or base's degree of dissociation. Another way to
describe the amount of dissociation is by percent
dissociation.
• Percent dissociation is calculated in a similar manner
as calculating your percentage on a test or
assignment.


• For percent dissociation, we use the hydronium ion,
or hydroxide ion concentrations for reasons which
will become more clear later in this module. In
general, for the acid HA

• For the acid HA,

• Or, for the base BOH,


• Where [HA] and [BOH] are the initial concentrations
of the acid and base, respectively.

Example 3
• Calculate the percent dissociation of a 0.100 mol/L
solution of formic acid (HCH2O2) if the hydronium
ion concentration is 4.21 x 10-3 mol/L.


Calculating Dissociation Constant
Given Percent Dissociation
Example 4
• Calculate the Kb of the hydrogen phosphate ion
(HPO42¯) if a 0.25 mol/L solution of sodium
hydrogen phosphate is dissociated is 0.080%.


Ionization of Water
• Recall, that water is amphoteric. That is, it
can act as both an acid and a base.
HA + H2O(l) ↔ H3O+(aq) + A¯(aq)
or
B + H2O(l) ↔ BH+(aq) + OH¯(aq)
• Very sensitive instruments have shown that
pure water actually dissociates into ions, or
ionizes, slightly. We call this process self-
ionization or autoionization.


Ionization of Water
• The equation for self-ionization is written as:
H2O(l) + H2O(l) ↔ H3O+(aq) + OH¯(aq)
or
H2O(l) ↔ H+(aq) + OH¯(aq)

• This indicates there is an equilibrium
established between hydronium and hydroxide
ions


Ion Product of Water
• If an equilibrium is established between
hydronium ions, hydroxide ions and water
molecules, an equilibrium law can be written:

• Since water is a liquid, the product of Ka and
water results in the ion product for water, Kw.


• The equilibrium law for water becomes:
• KW = [H3O+][OH¯]At 25°C, the concentration
of the hydronium and hydroxide ions are equal
at 1.0 x 10-7 mol/L.
• Therefore, at 25°C, the value of KW is
constant at 1.0 x 10-14.


• The equilibrium law for water becomes:

KW = [H3O+][OH¯]

• At 25°C, the concentration of the hydronium
and hydroxide ions are equal at 1.0 x 10-7 mol/
L.
• Therefore, at 25°C, the value of KW is
constant at 1.0 x 10-14.


Ionization of Water
• If the ionization of water occurs by the
equation
H2O(l) + H2O(l) ↔ H3O+(aq) + OH¯(aq)
• we can predict the effect of dissolving an acid
or base on hydronium and hydroxide ion
concentrations by using Le Chatelier's
Principle.
• Adding Acid ?
• Adding Base ?


Ionization of Water


Ionization of Water
Adding Acid
• When an acid is dissolved in water, the acid produces
a large amount of H3O+ ions. If the H3O+ ion
concentration increases, the equilibrium will shift to
the left to use up some of the added hydronium and
maintain Kw at 1.0 x 10-14.
• Since equilibrium shifts left, the hydroxide ion
concentration is reduced. Therefore, adding a strong
acid to water increases the hydronium ion
concentration and reduces the hydroxide ion
concentration.


Ionization of Water
Adding Base
• When a base is dissolved in water, the hydroxide ion
concentration increases. According to Le Chatelier's
Principle, the equilibrium shifts left to use up some of
the added hydroxide and maintain KW at 1.0 x 10-14.
• Since equilibrium shifts left, the hydronium ion
concentration is reduced. Therefore, adding a strong
base to water increases the hydroxide ion
concentration and reduces the hydronium ion
concentration.


Calculating Hydroxide Ion Concentration
• Example 1
If 2.5 moles of hydrochloric acid is dissolved
in 5.0 L of water, what is the concentration of
the hydroxide ions? Assume the volume
remains unchanged.
Calculating Hydronium Ion Concentration
• Example 2
O.40 g of NaOH is dissolved in water to make
a solution with a volume of 1.0 L. What is the
hydronium ion concentration in this solution?


The pH Scale


Defining pH
• In 1909, a Danish chemist, named Soren Sorensen
(1868 - 1939), developed a simplified system for
referring to the degree of acidity of a solution. He
used the pH or the potenz (power) of hydrogen.
Therefore, pH describes the concentration of the
hydronium or hydrogen ions in solution.
• pH is defined as the negative logarithm of the
hydronium ion concentration, or
pH = -log[H3O+] or pH = -log[H+]
(We will use the H3O+ion and H+ ion interchangeably.)


The pH Scale
• Values for pH in most solutions range from
0.0 to 14.0. Pure water is considered to be
neutral, or a pH of 7.0. The lower the pH, the
more acidic the solution. The higher the pH
the more alkaline or basic the solution.
– pH < 7 acidic
– pH = 7 neutral
– pH > 7 basic (alkaline)


The pH Scale
• Below is a pH scale with the pH values of some
common solutions.


pH Calculations
Example 1
• Calculate the pH of an HCl solution whose
concentration is 5.0 x 10-6 mol/L.


pH Calculations
Example 2
• The pH of a solution is 3.25. Calculate the
hydronium ion concentration in the solution.


Defining pOH
• Recall, the ion product for water:
KW = [H3O+][OH¯]
• If we take the negative log of each term, we get
-log(KW) = pH + pOH
• According to the rules of logs, when multiplying
terms is equivalent to adding their logs.
• If we calculate the negative log kW,
-log(KW) = -log(1.0 x 10-14) = 14.00
• So,
pH + pOH = 14.00


pOH Calculations
Example 3
• The pH of a solution is 10.30, what is the
hydroxide ion concentration?


pOH Calculations
Example 4
• What is the pH of a 5.0 x 10-5 mol/L Mg(OH)2
solution?


Calculating pH, Given Ka
• We can now calculate the pH of a weak
acid(or base) solution, given the percent
dissociation or Ka (Kb)and the acid
concentration.


Example 5
• Calculate the pH of a 0.10 mol/L hydrogen
sulfide solution. (Ka=1.0 x 10-7)


Example 5 - continued

Chem 40S Unit 5 Notes

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