Transcript ACIDS and BASES - JH Rose

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Sour to taste
React with some metals to form Hydrogen
gas
Turn Litmus RED
Phenolphthalein stays colorless
Electrolytes (conduct)
Form H+ (H3O+ when attached to water
molecules)
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Bitter to taste
Slippery to touch
Turn Litmus BLUE
Phenolphthalein turns MAGENTA
Electrolytes
Many form OH- in water
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Arrhenius (traditional)
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Examples
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Acids: produce H3O+ in water
Bases: produce OH- in water
HCl (g) → H+(aq) + Cl-(aq)
NaOH (s) → OH-(aq) + Na+(aq)
Most common definition
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Bronsted-Lowry
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Acids: H+ donor (proton donor)
Bases: H+ acceptor (proton acceptor)
Examples
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HCl → Cl- (donates H+)
NH3 → NH4+ (accepts H+)
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When acids and bases donate or accept
hydrogen ions, conjugate acids and bases
are formed.
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Conjugate Acid: particle formed when a Base
gains a H+
Conjugate Base: particle formed when an Acid
donates a H+
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Conjugate acid-base pair: two substances
related by the loss or gain of a single H+
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Always paired with an acid and base
Examples- Label acid/base and conjugate
acid/base
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NH3 + H2O → NH4+ + OHHCl + H2O → H3O+ + Cl
Hint: Acid/Base is always Reactant; Conjugate
Acid/Base is always Product
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Amphoteric substances can behave as acids
or bases (water)
Monoprotic acids donate only one H+
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Ex. HCl, HNO3
Polyprotic acids donate more than one H+
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H2SO4, H2CO3
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The stronger the acid, the weaker the
conjugate base.
H+ is the strongest acid that can exist in
equilibrium in aqueous solution.
OH- is the strongest base that can exist in
equilibrium in aqueous solution.
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Any acid or base that is stronger than H+ or
OH- simply reacts stoichiometrically to
produce H+ and OH-.
The conjugate base of a strong acid (e.g. Cl-)
has negligible acid-base properties.
Similarly, the conjugate acid of a strong base
has negligible acid-base properties.
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In pure water, the following equilibrium is
established:
H2O(l) + H2O(l)
K eq 
H3O+(aq) + OH-(aq)
[H3O ][OH- ]
2
[ H 2O ]
K eq  [H 2O]2  [H3O ][OH- ]
K w  [H3O ][OH- ]  1.0  1014
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This is referred to as the autoionization of
water.
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Weak acids and bases
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Weak Acids
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Weak electrolytes
Partially ionize in water (much less than 100%)
Establish equilibria
H3PO4, HC2H3O2, H2CO3
Weak Bases
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NH3, low [OH-]
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Pouvoir hydrogene: “Hydrogen Power”
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Measure of the acidity of a solution
Uses [H3O+] or [H+]
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Concentrations usually expressed as powers of
10
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pH = -log [H+]
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pH scale 0-7 acid, 7 neutral, 7-14 base
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Can be lower than 0 or higher than 14
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“Hydroxide Power”
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Measures alkalinity of a solution
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Uses [OH-]
pOH = -log [OH-]
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pOH scale 0-7 base, 7 neutral, 7-14 acid
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Can be lower than 0 or higher than 14
pH + pOH = 14
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From the autoionization of water
Calculating Ka from pH
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Weak acids are simply equilibrium calculations.
The pH gives the equilibrium concentration of H+.
Using Ka, the concentration of H+ (and hence the pH) can be
calculated.
– Write the balanced chemical equation clearly showing
the equilibrium.
– Write the equilibrium expression. Find the value for Ka.
– Write down the initial and equilibrium concentrations for
everything except pure water. We usually assume that
the change in concentration of H+ is x.
– Substitute into the equilibrium constant expression and
solve. Remember to turn x into pH if necessary
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Percent ionization is another method to
assess acid strength.
For the reaction
HA(aq) + H2O(l)
% ionization 
H3O+(aq) + A-(aq)
[H3O ]eqm
[HA]0
 100
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Percent ionization relates the equilibrium H+
concentration, [H+]eq, to the initial HA
concentration, [HA]0.
The higher percent ionization, the stronger
the acid.
Percent ionization of a weak acid decreases
as the molarity of the solution increases.
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Polyprotic acids have more than one ionizable
proton.
The protons are removed in steps not all at
once:
H2SO3(aq)
H+(aq) + HSO3-(aq) Ka1 = 1.7 x 10-2
HSO3-(aq)
H+(aq) + SO32-(aq)
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Ka2 = 6.4 x 10-8
It is always easier to remove the first proton in a
polyprotic acid than the second.
Therefore, Ka1 > Ka2 > Ka3 etc.
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For a conjugate acid-base pair
K w  Ka  Kb
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Therefore, the larger the Ka, the smaller the
Kb. That is, the stronger the acid, the weaker
the conjugate base.
Taking negative logarithms:
pK w  pKa  pKb
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Strong electrolytes
Ionize (separate) 100% in water
Strong Acids: HCl, HNO3, H2SO4, HClO3, HClO4,
HI, HBr
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Usually only source of H+, so pH can be calculated
from the molarity of the acid (unless < 10-6)
Strong Bases: NaOH, KOH, Ca(OH)2
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Most Group 1 and 2 metal hydroxides are strong
bases
Ionic metal oxides, hydrides, and nitrides
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Most salts are strong electrolytes (ionize in
solution)
Acid-Base properties result from their ions.
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Many ions react with water to form H+ and OH(hydrolysis)
Anions from weak acids are basic.
Anions from strong acids are neutral.
All cations (except alkali/alkaline earth metals)
are weak acids
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SA + SB = Neutral Salt
SA + WB = Acidic Salt
WA + SB = Basic Salt
WA + WB = Acidic or Basic Salt
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Based on relative strength of Ka and Kb
Ka > Kb = acidic
Ka < Kb = basic
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For compound H-X,
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If H is partially positive, then it is an acid
If H is partially negative, then it is a base
Bond Strength and Polarity affects acid/base
strength
Bond strength used to determine strength in a group;
Bond polarity used to determine strength in a period
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Acid strength tends to increase down a group; Base
strength tends to decrease down a group
Acid strength tends to increase L to R; Base strength
tends to decrease L to R
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Acids that contain OH groups bound to the
central atom ( Y – O – H)
Strength depends on Y and the atoms
attached to Y
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Increasing electronegativity of Y = increasing
acidity
Increasing the number of O atoms attached to Y
increases polarity, which increases strength
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Ex. HClO < HClO2 < HClO3 < HClO4
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Contain –COOH
Additional oxygen atom on the carboxyl
group increase the polarity of the O-H bond
and stabilizes the conjugate base
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Lewis Acid – electron pair acceptor
Lewis Base – electron pair donor
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Do not need to contain protons – most general definition
of acids/bases
Many Lewis acids have an incomplete octet (BF3)
Transition metal ions are usually Lewis acids
Lewis acids must have an empty orbital
Compounds with multiple bonds can be Lewis acids
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Amphoteric
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Contains carboxyl group (acid) and ammine
group (basic)
Proton of the carboxyl group is transferred to
the basic nitrogen of the ammine
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Results in a zwitterion or dipolar ion