Transcript 11/4

Common Acids
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Common Bases
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Indicators
• Indicators are chemicals that change color
depending on the solution’s acidity or
basicity.
• Many vegetable dyes are indicators.
– Anthocyanins
• Litmus
– From Spanish moss
– Red in acid, blue in base
• Phenolphthalein
– Found in laxatives
– Red in base, colorless in acid
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Definitions of Acids and Bases
• Arrhenius definition
– Based on H+ and OH-
• Brønsted–Lowry definition
– Based on reactions in which H+ is transferred
• Lewis definition
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Arrhenius Theory
• Acids: produce H+ ions in aqueous solution.
HCl(aq) → H+(aq) + Cl−(aq)
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Hydronium Ion
• The H+ ions produced by the acid are so reactive
they cannot exist in water.
– H+ ions are protons!
• Instead, they react with water molecules to produce
complex ions, mainly hydronium ion, H3O+.
H+ + H2O  H3O+
– There are also minor amounts of H+ with multiple water
molecules, H(H2O)n+.
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Arrhenius Theory
• Bases: produce OH− ions in aqueous solution.
NaOH(aq) → Na+(aq) + OH(aq)
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Arrhenius Acid–Base Reactions
• The H+ from the acid combines with the OH−
from the base to make a molecule of H2O.
– It is often helpful to think of H2O as H—OH.
• The cation from the base combines with the
anion from the acid to make a salt.
acid + base → salt + water
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
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Problems with Arrhenius Theory
• It does not explain why molecular substances, such
as NH3, dissolve in water to form basic solutions,
even though they do not contain OH– ions.
• It does not explain how some ionic compounds,
such as Na2CO3 or Na2O, dissolve in water to form
basic solutions, even though they do not contain
OH– ions.
• It does not explain why molecular substances, such
as CO2, dissolve in water to form acidic solutions,
even though they do not contain H+ ions.
• It does not explain acid–base reactions that take
place outside aqueous solution.
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Brønsted–Lowry Acid–Base Theory
• It defines acids and bases based on what happens
in a reaction.
• Any reaction involving H+ (proton) that transfers
from one molecule to another is an acid–base
reaction, regardless of whether it occurs in
aqueous solution or if there is OH− present.
• All reactions that fit the Arrhenius definition also fit
the Brønsted–Lowry definition.
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Brønsted–Lowry Theory
•
The acid is an H+ donor.
•
The base is an H+ acceptor.
–
•
Base structure must contain an atom with an unshared
pair of electrons.
In a Brønsted–Lowry acid–base reaction, the
acid molecule donates an H+ to the base
molecule.
H–A + :B  :A– + H–B+
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Brønsted–Lowry Acids
• Brønsted–Lowry acids are H+ donors.
– Any material that has H can potentially be a
Brønsted–Lowry acid.
– Because of the molecular structure, often one H in the
molecule is easier to transfer than others.
• When HCl dissolves in water, the HCl is the acid
because HCl transfers an H+ to H2O, forming
H3O+ ions.
– Water acts as base, accepting H+.
HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq)
acid
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base
Brønsted–Lowry Bases
• Brønsted–Lowry bases are H+ acceptors.
– Any material that has atoms with lone pairs can potentially
be a Brønsted–Lowry base.
– Because of the molecular structure, often one atom in the
molecule is more willing to accept H+ transfer than others.
• When NH3 dissolves in water, the NH3(aq) is the
base because NH3 accepts an H+ from H2O,
forming OH–(aq).
– Water acts as acid, donating H+.
NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq)
base
acid
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Amphoteric Substances
• Amphoteric substances can act as either an
acid or a base because they have both a
transferable H and an atom with lone pair
electrons.
• Water acts as base, accepting H+ from HCl.
HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq)
• Water acts as acid, donating H+ to NH3.
NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq)
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Brønsted–Lowry Acid–Base Reactions
• One of the advantages of Brønsted–Lowry theory is that it
illustrates reversible reactions to be as follows:
H–A + :B  :A– + H–B+
• The original base has an extra H+ after the reaction, so it
will act as an acid in the reverse process.
• And the original acid has a lone pair of electrons after the
reaction, so it will act as a base in the reverse process:
:A– + H–B+  H–A + :B
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Conjugate Acid–Base Pairs
• In a Brønsted-–Lowry acid–base reaction,
– the original base becomes an acid in the reverse reaction.
– the original acid becomes a base in the reverse process.
• Each reactant and the product it becomes is called
a conjugate pair.
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Conjugate Pairs
A base accepts a proton and becomes a conjugate acid.
An acid donates a proton and becomes a conjugate base.
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Arrow Conventions
• Chemists commonly use two kinds of
arrows in reactions to indicate the degree
of completion of the reactions.
• A single arrow indicates that all the
reactant molecules are converted to
product molecules at the end.
• A double arrow indicates the reaction
stops when only some of the reactant
molecules have been converted into
products.
–  in these notes
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Strong or Weak
• A strong acid is a strong electrolyte.
– Practically all the acid molecules ionize.
• A strong base is a strong electrolyte.
– Practically all the base molecules form OH– ions,
either through dissociation or reaction with water.
• A weak acid is a weak electrolyte.
– Only a small percentage of the molecules ionize, .
• A weak base is a weak electrolyte
– only a small percentage of the base molecules form OH–
ions, either through dissociation or reaction with water, .
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Strong Acids
• Strong acids donate practically all their H’s.
 100% ionized in water
 Strong electrolyte
• [H3O+] = [strong acid]
 [X] means the molarity of X
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Weak Acids
• Weak acids donate a small fraction of their H’s.
•
 Most of the weak acid molecules do not
donate H to water.
 Much less than 1% ionized in water
[H3O+] << [weak acid]
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Ionic Attraction and Acid Strength
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General Trends in Acidity
• The stronger an acid is at donating H, the weaker
the conjugate base is at accepting H.
• Higher oxidation number = stronger oxyacid
– H2SO4 > H2SO3; HNO3 > HNO2
• Cation stronger acid than neutral molecule; neutral
stronger acid than anion
– H3O+ > H2O > OH−; NH4+ > NH3 > NH2−
– Trend in base strength opposite
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Autoionization of Water
• Water is amphoteric; it can act either as an acid or a base.
– Therefore, there must be a few ions present.
• About 2 out of every 1 billion water molecules form ions
through a process called autoionization.
H2O  H+ + OH–
H2O + H2O  H3O+ + OH–
• All aqueous solutions contain both H3O+ and OH–.
– The concentration of H3O+ and OH– are equal in water.
– [H3O+] = [OH–] = 10−7M at 25 °C
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Measuring Acidity: pH
• The acidity or basicity of a
•
solution is often expressed
as pH.
pH = −log[H3O+]
Exponent on 10 with a
positive sign
pHwater = −log[10−7] = 7
Need to know the [H3O+]
concentration to find pH
• pH < 7 is acidic; pH > 7 is
•
basic. pH = 7 is neutral.
[H3O+] = 10−pH
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What Does the pH Number Imply?
• The lower the pH, the more acidic the solution;
the higher the pH, the more basic the solution.
1 pH unit corresponds to a factor of 10 difference
in acidity.
• Normal range of pH is 0 to 14.
pH 0 is [H3O+] = 1 M; pH 14 is [OH–] = 1 M.
pH can be negative (very acidic) or larger than 14
(very alkaline).
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pOH
• Another way of expressing the acidity/basicity of
•
a solution is pOH.
pOH = −log[OH], [OH] = 10−pOH
pOHwater = −log[10−7] = 7
Need to know the [OH] concentration to find pOH
• pOH < 7 is basic; pOH > 7 is acidic; pOH = 7 is
•
neutral.
pH + pOH = 14.0
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Relationship between pH and pOH
• pH + pOH = 14.00 at 25 °C.
– You can use pOH to find the pH of a
solution.
[H3O ][OH ]  K w  1.0  1014



 
 ]  14.00

log
[H
O
]


log
[OH
     
 log [H3O ][OH ]   log 1.0  1014

3
pH  pOH  14.00
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[H3O+] and [OH−] in a Strong Acid or
Strong Base Solution
• There are two sources of H3O+ in an aqueous
solution of a strong acid—the acid and the water.
• There are two sources of OH− in an aqueous
solution of a strong base—the base and the water.
• For a strong acid or base, the contribution of the
water to the total [H3O+] or [OH−] is negligible.
– The [H3O+]acid shifts the Kw equilibrium so far that [H3O+]water is
too small to be significant.
• Except in very dilute solutions, generally < 1 × 10−4 M
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Finding pH of a Strong Acid or Strong
Base Solution
• For a monoprotic strong acid [H3O+] = [HAcid].
– 0.10 M HCl has [H3O+] = 0.10 M and pH = 1.00
• For a strong ionic base, [OH−] = (number OH−) ×
[Base].
– For molecular bases with multiple lone pairs available,
only one lone pair accepts an H; the other reactions
can generally be ignored.
– 0.10 M Ca(OH)2 has [OH−] = 0.20 M and pH = 13.30.
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Example 15.1 Identifying Brønsted–Lowry Acids and Bases and
Their Conjugates
In each reaction, identify the Brønsted–Lowry acid, the Brønsted–Lowry base, the conjugate acid, and the
conjugate base.
a.
b.
Solution
a. Because H2SO4 donates a proton to H2O in this reaction, it is the acid (proton donor). After H2SO4 donates the
proton, it becomes HSO4−, the conjugate base. Because H2O accepts a proton, it is the base (proton acceptor).
After H2O accepts the proton, it becomes H3O+, the conjugate acid.
H2SO4(aq) + H2O(l) ⟶ HSO4−(aq) + H3O+(aq)
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Example 15.1 Identifying Brønsted–Lowry Acids and Bases and
Their Conjugates
Continued
Solution
b. Because H2O donates a proton to HCO3− in this reaction, it is the acid (proton donor). After H 2O donates the
proton, it becomes OH−, the conjugate base. Because HCO3− accepts a proton, it is the base (proton acceptor).
After HCO3− accepts the proton, it becomes H2CO3, the conjugate acid.
HCO3−(aq) + H2O(l)
H2CO3(aq) + OH−(aq)
For Practice 15.1
In each reaction, identify the Brønsted–Lowry acid, the Brønsted–Lowry base, the conjugate acid, and the
conjugate base.
a.
b.
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Example 15.3 Calculating pH from [H3O+] or [OH−]
Calculate the pH of each solution at 25 °C and indicate whether the solution is acidic or basic.
a. [H3O+] = 1.8 × 10−4 M
b. [OH−] = 1.3 × 10−2 M
Solution
a. To calculate pH, substitute the given [H3O+] into the pH equation.
Since pH < 7, this solution is acidic.
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Example 15.3 Calculating pH from [H3O+] or [OH−]
Continued
b. First use Kw to find [H3O+] from [OH−].
Then substitute [H3O+] into the pH expression to find pH.
Since pH > 7, this solution is basic.
For Practice 15.3
Calculate the pH of each solution and indicate whether the solution is acidic or basic.
a. [H3O+] = 9.5 × 10−9 M
b. [OH−] = 7.1 × 10−3 M
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Example 15.4 Calculating [H3O+] from pH
Calculate the [H3O+] for a solution with a pH of 4.80.
Solution
To find the [H3O+] from pH, start with the equation that defines pH. Substitute the given value of pH and then
solve for [H3O+]. Since the given pH value was reported to two decimal places, the [H 3O+] is written to two
significant figures. (Remember that 10log x = x (see Appendix I ) . Some calculators use an inv log key to represent
this function.)
For Practice 15.4
Calculate the [H3O+] for a solution with a pH of 8.37.
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