Acids & Bases
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Properties of Acids
• Sour taste
• Change color of acid-base indicators (red in pH
paper)
• Some react with active metals to produce
hydrogen gas
Ba(s) + H2SO4(aq)
BaSO4(s) + H2(g)
• Some react with bases to neutralize and form
salt and water
H2SO4 (aq) + 2NaOH(aq)
• Some are electrolytes
Na2SO4 (aq) + 2H2O(l)
Examples of Acids
• Lemons and oranges - citric acid
• Vinegar - 5% by mass acetic acid
• Pop and fertilizer - phosphoric acid
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QuickTi me™ and a
TIFF ( Uncompressed) decompressor
are needed to see thi s pi ctur e.
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
QuickTi me™ and a
TIFF ( Uncompressed) decompressor
are needed to see thi s pi ctur e.
Quic kTime™ and a
TIFF (Unc ompres sed) dec ompres sor
are needed to see this pic ture.
Properties of Bases
• Bitter taste
• Change color of acid-base indicators
(blue in pH paper)
• Dilute aqueous solutions feel slippery
Ex. Soap
• Some react with acids to neutralize and
form salt and water
• Some are electrolytes
Examples of Bases
• Soap - NaOH
• Household cleaners - NH3
• Antacids - Ca(OH)2, Mg(OH)2
Quic kTime™ and a
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are needed to see this pic ture.
QuickTime™ and a
TIFF (Uncompr essed) decompressor
are needed to see this picture.
QuickTi me™ and a
TIFF ( Uncompressed) decompressor
are needed to see thi s pi ctur e.
Arrhenius Acids
• Acids that increase the concentration of
hydronium (H3O+) in aqueous solutions
HNO3(aq) + H2O(l)
H3O+(aq) + NO3-(aq)
acid
H+ + NO3- + H2O
Why do acids produce H3O+?
• H+ is extremely attracted to the unshared pair of
electrons on the water molecule so it donates itself to
this molecule where it becomes covalently bonded.
The ion formed is known as the hydronium ion (H3O+)
H+
Arrenius Bases
• Bases that increase the concentration of
hydroxide ions (OH-) in aqueous
solutions
H2O
NaOH(s)
Na+(aq) + OH-(aq)
Strength of Acids & Bases
• Strong acids & bases completely ionize in
aqueous solutions
H2SO4 + H2O
NaOH
H3O+ + HSO4Na+ + OH-
• Strong acids & bases are strong electrolytes
• A list of strong acids & bases can be found on
pg. 460-461
• Weak acids & bases only partially break down
into ions when in aqueous solutions
HCN + H2O
NH3 + H2O
H3O+ + CNNH4+ + OH-
• Weak acids & bases are weak electrolytes
• A list of weak acids & bases can be found on
pg. 460-461
Why can we drink H2O?
• Water self ionizes to form equal
concentrations of H3O+ and OHH2O(l) + H2O(l)
H3O+(aq) + OH-(aq)
• A substance is considered “neutral”
when [H3O+] = [OH-]
• [H3O+] concentration = 1.0 x 10-7M
• [OH-] concentration = 1.0 x 10-7 M
When [H3O+] = [OH-]
• If [H3O+] > 1.0 x 10-7 M, the solution is acidic
• If [OH-] > 1.0 x 10-7 M, the solution is basic
• To find the concentration of [H3O+] or [OH-] in
acidic or basic solutions, the following
equation can be used:
1.0 x 10-14 M2 = [H3O+] [OH-]
1.0 x 10-14 M2 = ionization constant for H2O (Kw)
Sample Problem
• A 1.0 x 10-4 M solution on HNO3 has
been prepared for laboratory use.
a. Calculate the [H3O+] of this solution
b. Calculate the [OH-] of this solution
c. Is this solution acidic or basic?
Why?
d. Substitute H2SO4 as the acid. How
would the calculations change?
Sample Problem
• An aqueous 3.8 x 10-3 M NaOH solution has
been prepared for laboratory use.
a. Calculate the [H3O+] of this solution
b. Calculate the [OH-] of this solution
c. Is this solution acidic or basic?
Why?
d. Substitute Ca(OH)2 as the base. How
would the calculations change?
Practice Problems
• Complete practice problems on pg. 484
#1-4
The pH scale
• The pH scale measures the power of the
hydronium ion [H3O+] in a solution
• The scale typically goes from 1-14 (although
it can extend below or above it under extreme
conditions)
• The following equations can be used to
determine the pH or [H3O+] of a solution:
pH = -log [H3O+] [H3O+] = antilog (-pH)
[H3O+] = 1 x 10-pH
pH > 7 basic
pH = 7 neutral
pH < 7 acidic
The pOH scale
• The pOH scale measures the power of the
hydroxide ion [OH-] in a solution
• The scale typically goes from 1-14 (although
it can extend below or above it under extreme
conditions)
• The following equations can be used to
determine the pOH or [OH-] of a solution:
pOH = -log [OH-] [OH-] = antilog (-pOH)
[OH-] = 1 x 10-pOH
pH + pOH = 14
Sample Problems
•
Calculate the pH of each of the
following. Classify as acidic or basic.
a. 1.3 x 10-5 M NaOH
b. 1.0 x 10-4 M HCl
Sample Problems
•
What is the [H3O+] for each of the
following? Classify as acidic or basic.
a. pH = 5.8
b. pOH = 8.9
Sample Problems
•
What is the [OH-] for each of the
following? Classify as acidic or basic.
a. [H3O+] = 9.5 x 10-10 M
b. pOH = 1.3
Practice Problems
• Complete practice problems on
pg. 487 #1
pg. 488 #1-4
pg. 490 #1-4
Expansion of the Acid-Base
Theory
• Substances can still act as an acid or
base if they are not dissolved in water to
make a solution
Bronsted-Lowry Acids
• A molecule or ion that is a proton (H+)
donor
HCl(g) + NH3(g)
H+ donor
NH4(g)+ + Cl-(g)
Bronsted-Lowry Bases
• A molecule or ion that is a proton (H+)
acceptor
HCl(g) + NH3(g)
H+ acceptor
NH4+(g) + Cl-(g)
• In a Bronsted-Lowry acid-base reaction,
protons (H+) are transferred from one
reactant (the acid) another (the base)
Monoprotic versus Polyprotic
Acids
• Monoprotic acids can only donate 1
proton per molecule
HCl(g) + H2O(l)
Monoprotic
H3O+(aq) + Cl-(aq)
• Polyprotic acids can donate more than
one proton per molecule
H2SO4(aq) + H2O(l)
H3O+(aq) + HSO4-(aq)
Polyprotic
HSO4-(aq) + H2O(l)
H3O+(aq) + SO4-2(aq)
One additional proton can still be donated
Conjugate acids & bases
• A conjugate acid is the species that is
formed when a Bronsted-Lowry base
gains a proton
• A conjugate base is the species that
remains after a Bronsted-Lowry acid
has given up a proton
HF(aq) + H2O(l)
F-(aq) + H3O+(aq)
acid
base
Conjugate
base
Conjugate
acid
More examples
CH3COOH(aq) + H2O(l)
acid
base
HCl(aq) + H2O(l)
H3O+(aq) + CH3COO-(aq)
CB
CA
H3O+(aq) + Cl-(aq)
acid
bas
CA
CB
e
Proton transfer reactions favor the production of the weaker
acid and base.
Use table 15-6 on pg. 471 in your text to compare the relative
strengths of acids and bases
Is H2O an acid or a base?
• H2O is amphoteric, it can react as either an acid or a
base
• If H2O reacts with a compound that is a stronger acid
than itself, it acts as a base
• If H2O reacts with a weaker acid, it will act as the acid
H2SO4(aq) + H2O(l)
H3O+(aq) + HSO4-(aq)
Base
H+ acceptor
NH3(aq) + H2O(l)
Acid
H+ donor
NH4+(aq) + OH-(aq)
OH- in a molecule
• When an OH- group is covalently
bonded in a molecule, it is referred to as
a hydroxyl group
• Hydroxyl groups are present in many
organic compounds
Ex. Acetic acid (HC2H3O2) or CH3COOH
Hydroxyl
group
How does the OH- make
something acidic?
• In order for a compound with an OHgroup to be acidic, H2O must be able to
attract the H atom from the OH- group
and act as a proton donor
CH3COOH(aq) + H2O(l)
H3O+(aq) + CH3COO-(aq)
The more O atoms bonded to the OH- group, the more acidic
the compound is likely to be.
Oxygen is highly electronegative and will attract electrons
closer to it, making the OH- bond more polar. This will allow
H2O to “steal” the H atoms more easily.
Why are substances with OH- covalently
bonded to it sometimes not acidic?
• Ex. Acetic acid (CH3COOH) versus
ethanol (C2H5OH)
Acetic acid
Ethanol
Acetic acid- the 2 O atom on the C atom draws electron
density away from the OH- group, making the bond more
polar. This allows the H+ to be donated more easily
Ethanol- this compound is essentially neutral. It does not
have a second O atom to make the bond as polar. It
would be classified as a very weak acid because it is
harder to donate H+.
Further expansion of acidbase theory
• Substances can still act like an acid or
base if they do not contain hydrogen at
all
Lewis acids & bases
• A Lewis acid is an atom, ion, or molecule that
accepts an electron pair to form a covalent
bond
+
+
Ag
+
2NH
[H
N-Ag--NH
]
(aq)
3(aq)
3
3
e pair
acceptor
e pair
donator
• A Lewis base is an atom, ion, or molecule that
donates an electron pair to form a covalent
bond
Sample Lewis acid-base
problem
• For the following equation, which
reactant is the Lewis acid? Lewis
base?
BF3(aq) + F-(aq)
BF4- (aq)
• BF3 is the Lewis acid because it is the epair acceptor
• F- is the Lewis base because it is the epair donor
Review of acid-base
categorization
Type
Acid
Base
Arrhenius
H3O+
producer
OHproducer
BronstedLowry
Proton (H+)
donor
Proton (H+)
acceptor
Lewis
e- pair
acceptor
e- pair
donor
Strong Acid-Base
Neutralization
• When equal parts of acid and base are
present, neutralization occurs where a
salt and water are formed
HCl(aq) + NaOH(aq)
NaCl(aq) + H2O(l)
Sample Problems
•
•
•
•
H2CO3 + Sr(OH)2
HClO4 + NaOH
HBr + Ba(OH)2
NaHCO3 + H2SO4
Titrations
• When you have a solution with an unknown
concentration, you can find it by reacting it
completely with a solution of known
concentration
• This process is known as titrating
• To perform a titration, an instrument called a
buret can be used to precisely measure
amounts of solution, drop by drop
Titration Termonology
• Equivalence point - the point at which the
known and unknown concentration solutions
are present in chemically equivalent amounts
moles of acid = moles of base
Indicator - a weak acid or base that is added
to the solution with the unknown
concentration before a titration so that it will
change color or “indicate” when in a certain
pH range (table 16-6 on pg. 495 in your text
will show various indicators and their color
ranges)
• End point - the point during a titration
where an indicator changes color
• The 2 most common indicators we will
use in our chemistry class will be:
• Phenolphthalein - turns very pale pink at
a pH of 8-10
• Bromothymol blue - turns pale green at
a pH of 6.2-7.6
Quic k Ti me™ and a
T IFF (Unc om pres s ed) dec om pres s or
are needed to s ee t his pic t ure.
Phenolpthalein is clear at pH<8,
pale pink at pH 8-10 and
magenta at pH >10
QuickTime™ and a
TIFF (U ncompressed) decompressor
are needed to see t his picture.
Bromothymol blue
Practice Titration for
an unknown acid
QuickTi me™ and a
T IFF (Uncom pressed) decom pressor
are needed to see t his pict ure.
• 1. Titrate 5.0 of mL of unknown HCl into a 250 mL
erlenmeyer flask - *remember to document the
starting amount and ending amount of acid on the
buret to prevent error
• 2. Add 2 drops of indicator (phenolphthalein) to the
flask - the color of the solution should be clear
• 3. Titrate with .5M NaOH, continuously swirling the
flask, until the solution turns very pale pink for 30
seconds - *remember to document the starting
amount and ending amount of base on the buret
• 4. Mathematically determine the concentration of the
unknown HCl solution by using the following
equation:
Titration Equation
MAVA = MBVB
MA = molarity (mol/L) of acid
VA = volume in L of acid
MB = molarity (mol/L) of base
VB = volume in L of base
molesA = molesB
5. After calculating the molarity of the unknown
acid experimentally, get the theoretical
molarity and calculate % error
Practice titration for an
unknown base
• 1. Titrate 5.0 of mL of unknown NaOH into a 250 mL
erlenmeyer flask - *remember to document the starting
amount and ending amount of base on the buret to prevent
error
• 2. Add 2 drops of indicator (phenolphthalein) to the flask the color of the solution should be magenta
• 3. Titrate with .5M HCl, continuously swirling the flask, until
the solution turns very pale pink for 30 seconds - *remember
to document the starting amount and ending amount of acid
on the buret
• 4. Mathematically determine the concentration of the
unknown NaOH solution by using MAVA = MBVB
• 5. After calculating the molarity of the unknown base
experimentally, get the theoretical molarity and calculate %
error
How do pH indicators work?
• Acid-base indicators are usually weak acids
or bases that are in equilibrium and show
color changes when a stress is applied
HIn
H+ + Inred
blue
In acidic solutions, the H+ concentration increases.
The stress will cause a shift to the left (red color).
In basic solutions, the OH- concentration increases.
These ions will combine with H+ which will cause a
shift to the right (blue color)