Download Acids and Bases - Hobbs High School

Acids and Bases; pH
Acid Base Reactions
• We learned that an acid base reaction was
a special type of a double displacement
reaction.
• An acid reacts with a base to produce a
salt and water.
What is an acid?
• According to the Arrhenius acid-base
concept, a substance is classified as an acid if
it ionizes to form hydrogen(+) ions in
aqueous solution. For example, hydrochloric
acid reacts with water to form hydrogen ions
which are transferred to a water molecule to
form a hydronium ion (H3O+).
What is an acid?
• Other systems classify substances as
acids if they act as proton donors
(Bronsted-Lowry theory) or as electronpair acceptors (Lewis theory).
• These two classification methods are not
limited to solutions in water, but are used
for other solvents as well.
Common Acids
• Common acids are citric
acid (from certain fruits and
veggies, notably citrus
fruits)
• ascorbic acid (vitamin C, as
from certain fruits)
• vinegar (5% acetic acid)
• carbonic acid (for
carbonation of soft drinks)
• lactic acid (in buttermilk)
What is a Base?
• Depending on the classification system,
bases produce OH- ions in aqueous
solutions (Arrhenius acid-base concept),
are proton acceptors (Bronsted-Lowry
theory), or are electron-pair donors
(Lewis theory).
• In other words, bases behave oppositely
from acids.
Common Bases
•
•
•
•
Detergents
Soap
Lye (NaOH)
Household
ammonia (aqueous)
• Baking soda
Acid Strength
• The strength of an acid refers to its ability or
tendency to lose a proton (H+).
• A STRONG acid completely ionizes (dissociates)
in a solution to produce hydronium ions (H3O+). It
ionizes 100%.
• A mole of a strong acid HA dissolves in water to
yield one mole of H3O+ and one mole of the
conjugate base, A-. Effectively there is no ionized
HA remaining. (This reversible reaction has
achieved chemical equilibrium.)
HA + H2O ↔ H3O+ + A-
Strong Acids
• These six acids are the strong acids:
1. hydrochloric acid (HCl)
2. hydroiodic acid (HI)
3. hydrobromic acid (HBr)
4. perchloric acid (HClO4)
5. nitric acid (HNO3)
6. sulfuric acid (H2SO4).
Weak Acids
• Weak acids do not ionize fully when dissolved in
water. An example is acetic (or ethanoic) acid.
CH3COOH + H2O ↔ CH3COO- + H3O+
At any one time, only about 1% of the acetic acid
molecules have converted into ions. The rest
remain as simple acetic acid molecules in solution.
• Hydrogen fluoride dissolves in water to produce
hydrofluoric acid (HF), a weak inorganic acid.
• Formic acid (HCOOH) is a weak organic acid (think
ants!).
• Most organic acids are weak (e.g., acetic, formic,
citric, lactic, carbonic).
Strong bases
• Strong bases are completely ionized in
solution to form hydroxide ions (OH-). For
example, KOH dissolves in water in the
reaction
KOH → K+ + OH• Group 1 and Group 2 metals form strong
bases. Examples are
Sodium hydroxide
Potassium hydroxide
Cesium hydroxide
Calcium hydroxide
NaOH
KOH
CsOH
Ca(OH)2
Weak bases
• A weak base is one which doesn't convert fully
into hydroxide ions in solution.
• Ammonia (NH3) is a typical weak base.
Ammonia itself obviously doesn't contain
hydroxide ions, but it reacts with water to
produce ammonium ions and hydroxide ions.
NH3 + H2O ↔ NH4+ + OH• This eqilibrium reaction is reversible, and at any
one time about 99% of the ammonia is still
present as ammonia molecules. Only about 1%
has actually produced hydroxide ions.
Don’t confuse…
• It is important that you don't confuse the
words strong and weak with the terms
concentrated and dilute.
• It is perfectly possible to have a
concentrated solution of a weak acid, or a
dilute solution of a strong acid.
• For instance, a 12M HF is a concentrated
solution of a weak acid, while 0.01M
HNO3 is a dilute solution of a strong acid.
Electrolytic Properties
• Strong acids and strong bases are good
conductors of electricity while weak acids
and weak bases are not.
• Why? Remember that strong acids and
strong bases ionize almost 100%, creating
many ions which act as mobile charge
carriers. Weak acids and weak bases
have few ions and few charge carriers.
Amphoteric Compounds
• The ability of some chemicals to act either
as an acid or a base is called
amphoterism. Whether an amphoteric
chemical acts as an acid or a base
depends on what other chemicals happen
to be around.
• Water is such an amphoteric substance.
We say that water has a conjugate base,
OH-, and a conjugate acid, H3O+.
Amphoteric Compounds
• If a base (like ammonia, NH3) is present,
water can act as an weak acid and react
by donating a proton to that base. In doing
so, water is changed into its conjugate
base, hydroxide ion.
H2O + NH3 → NH4+ + OH-
Amphoteric Compounds
• If an acid (like HCl) is present, water can
act as a weak base and react by
accepting a proton from that acid. In doing
so, water is changed into its conjugate
acid, hydronium ion.
H2O + HCl → Cl- + H3O+
Self-Ionization of Water
• The self-ionization of water is another example
of water being able to react either as an acid or a
base.
• The molecules in pure water continuously collide
and react with one another. In that reaction, one
water molecule can transfer a proton to
another water molecule. One water molecule
acts as an acid and the other acts as a base.
• The solution is neutral because equal quantities
of H3O+ and OH- are made.
H2O + H2O → OH- + H3O+
More Amphoteric Compounds
• Oxides of weakly electropositive metals
such as zinc, lead, aluminum, tin, and
beryllium (depends on their oxidation state)
• Amino acids and proteins, which have
amine and carboxylic acid groups
• Ammonia (NH3) is another self-ionizable
compound
• Bicarbonate ion, HCO3-
Standard Notation for
Acids and Bases
• To represent concentrations of ions in
moles per liter, the formula of the
particular ion or molecule is enclosed
in brackets [ ].
• For instance [H3O+] means “hydronium ion
concentration in moles per liter” or “molar
hydronium ion concentration.”
• Similarly, [OH-] means “hydroxide ion
concentration in moles per liter” or “molar
hydroxide ion concentration.”
Ion Concentrations as Molarity
• Did you notice that we already know a
measure of concentration that is
expressed in “moles of solute per liter of
solution”? Yes, you recognize this as
“molarity.”
• So both “molar hydronium ion
concentration” and “molar hydroxide ion
concentration” are expressed as molarity.
Standard Values
• At room temperature (25 °C), pure water has
[H3O+] = 1.0 x 10-7 M and [OH-] = 1.0 x 10-7 M.
• The mathematical product of [H3O+] and [OH-]
remains constant in water and dilute aqueous
solutions at constant temperature.
• We call this product the ionization constant
of water, Kw, which is expressed as:
Kw = [H3O+] [OH-]
• Therefore, at 25 °C , the ionization constant of
water has this value:
Kw = [H3O+] [OH-] = 1.0 x 10-14
Acidity and Basicity
• pH (pouvoir hydrogène or ‘power of
hydrogen’) is a measure of the acidity or
basicity of an aqueous solution. The molar
concentration of hydronium ions in the
solution [H3O+] is used in the formula for pH:
pH = -log[H3O+]
(Note: This is log10, not ln, the natural log.)
• Solutions with a pH less than 7 are said to
be acidic and solutions with a pH greater
than 7 are basic or alkaline.
• Pure water has a pH very close to 7 (neutral).
pH Scale
Pure
water
Relationship between pH and pOH
• Similarly, the concentration of OH ions is
expressed as pOH
pOH = -log[OH-]
• This leads us to a relationship between pH
and pOH at 25 °C:
pH + pOH = 14.0
Sample Calculations of pH
1. What is the pH of a solution that contains
a [H3O+] of 1.0 x 10-4 M?
pH = -log[H3O+]
= -log[1.0 x 10-4]
= -(-4)
= 4
Sample Calculations of pH
2. What is the pH of a 0.001 M solution of
hydrochloric acid (HCl)? (Here we use
molarity equivalently to concentration.)
pH = - log [.001]
=
(-)
LOG
= - (-3.00)
= 3.00
(.001)
Sample Calculations of pH
3. What is the pH of a solution if the
hydronium concentration, [H3O+], is
3.4 x 10-5 M?
pH = - log [H3O+] = - log [3.4 x 10-5 ]
=
(-)
LOG
= - (- 4.47)
= 4.47
(3.4
2nd
EE
(-)
5)
Sample Calculations of pH
4. What is the pH of 735 L of a solution
holding 0.34 moles of nitric acid (HNO3)?
To find the concentration, we need to find
the molarity:
M=
𝒎𝒐𝒍𝒆𝒔 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒆
𝒗𝒐𝒍𝒖𝒎𝒆 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏 (𝑳)
nitric acid
= 0.34 moles
735 L
M = 0.000463 M HNO3
Substitute this into the equation for pH:
pH = - log(0.000463) = 3.33
Sample Calculations of pH
5. What is the pH of a 1.0 x 10-3 M NaOH
solution at room temperature?
Hmm, NaOH is a strong base, so it will disassociate
completely into Na+ and OH- ions. Therefore the
concentration of OH- must be the same as the
solution’s molarity: [OH-] = 1.0 x 10-3
We use concentration of OH- to find pOH:
pOH = -log[OH-] = -log[1.0 x 10-3]
pOH = 3.0
We know that pH + pOH = 14.0, so
pH = 14.0 – pOH = 14.0 – 3.0 = 11.0
We expected a high pH for a strong base!
Calculating [H3O+] from pH
If you know the pH of a substance, you can calculate
the ion concentration of [H3O+] from this formula:
[𝐻3 𝑂+ ] = 10−𝑝𝐻
For a substance with a pH of 7.52, and
using your graphing calculator,
2nd
10x
LOG
(-)
7.52 = 3.01995 E -8
( - 7.52)
= 3.01995 x 10-8 = [𝐻3 𝑂+ ]
Classifying Acids
• Most acids used in the laboratory can be
classified as either binary acids or oxyacids.
• Binary acids are acids that consist of two
elements, usually hydrogen and one of the
halogens.
• Oxyacids are acids that contain hydrogen,
oxygen, and a THIRD element, usually a
NONMETAL.
• ‘Acid’ usually refers to a solution in water of
one of these special compounds.
NAMING ACID RULES:
Binary Acids
1. Anion does NOT contain oxygen (O)
use prefix ‘hydro-’ AND suffix ‘–ic’
added to the root word
Examples:
HCl
HF
HBr
hydrochloric acid
hydrofluoric acid
hydrobromic acid
NAMING ACID RULES:
Oxyacids
2. Anion DOES contain oxygen (O)

change suffix ‘–ate’ to ‘–ic ‘
(“You ate something icky.”)
Examples: H2SO4 sulfuric acid
HNO3 nitric acid

change suffix ‘–ite’ to ‘–ous’
(“A snake bite is poisonous.”)
Examples: H2PO3 phosphorous acid
HNO2 nitrous acid
Naming More Complicated Acids
• For example, H2SO5, H2SO4, H2SO3, and
H2SO2 are all acids. How do we name
them?
• Our point of reference is the ‘-ic’ acid
made from the ‘-ate’ ion. (Here H2SO4 or
sulfuric acid)
• Then the acid with one more oxygen than
the ‘-ic’ acid is called the per-_________ic acid. (Persulfuric acid)
Naming More Complicated Acids
• The acid with one less oxygen than the
‘-ic’ acid is called the ______-ous acid.
(This is the ‘-ite’ to ‘-ous’ change.)
• If the acid has one less oxygen than the
‘-ous’ acid, it is called the
hypo-_______-ous acid.
Examples:
H2SO5 = persulfuric acid
HNO4 = pernitric acid
H2SO4 = sulfuric acid
HNO3 = nitric acid
H2SO3 = sulfurous acid
HNO2 = nitrous acid
H2SO2 = hyposulfurous acid HNO = hyponitrous
acid