B sc i chemistry i u ii ionic equilibria in aqueous solution a
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Transcript of B sc i chemistry i u ii ionic equilibria in aqueous solution a
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Course: B.Sc. I
Subject: Chemistry I
Unit: II
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Acids
Produce H+ (as H3O+) ions in water (the hydronium ion is a
hydrogen ion attached to a water molecule)
Taste sour
Corrode metals
React with bases to form a salt and water
pH is less than 7
Turns blue litmus paper to red “Blue to Red A-CID”
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Bases
Produce OH- ions in water
Taste bitter, chalky
Feel soapy, slippery
React with acids to form salts and water
pH greater than 7
Turns red litmus paper to blue “Basic Blue”
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Some Common Bases
NaOH sodium hydroxide lye
KOH potassium hydroxide liquid soap
Ba(OH)2 barium hydroxide stabilizer for plastics
Mg(OH)2 magnesium hydroxide “MOM” Milk of magnesia
Al(OH)3 aluminum hydroxide Maalox (antacid)
4
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Concepts of acid-base theory
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Acid/Base definitions
Definition #1: Arrhenius (traditional)
Acids – produce H+ ions (or hydronium ions H3O+)
Bases – produce OH- ions
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Arrhenius Concept of Acids and BasesAccording to the Arrhenius concept of acids
and bases,
an Acid is a substance that, when dissolved in water, increases the concentration of hydronium ion or H+ ion (H3O
+ ).• the aqueous hydrogen ion is actually chemically
bonded to water, that is, H3O+.
• Chemists often use the notation H+(aq) for the H3O
+(aq) ion.
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Arrhenius concept of basesA base, in the Arrhenius concept, is a
substance that, when dissolved in water, increases the concentration of hydroxide ion, OH-(aq).
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Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
3
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Limitations
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Brønsted-Lowry Concept of Acids and Bases• According to the Brønsted-Lowry concept,
• An acid is the species donating the proton
in a proton-transfer reaction.
• A base is the species accepting the proton in a
proton-transfer reaction
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Consider the reaction of NH3 and H20.
)aq(OH )aq(NH )l(OH )aq(NH 423
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Consider the reaction of NH3 and H2O.
– In the forward reaction, NH3 accepts a proton
from H2O. Thus, NH3 is a base and H2O is an
acid.
)aq(OH )aq(NH )l(OH )aq(NH 423
H+
base acid
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Consider the reaction of NH3 and H2O.
– In the reverse reaction, NH4+ donates a
proton to OH-. The NH4+ ion is the acid and
OH- is the base.
)aq(OH )aq(NH )l(OH )aq(NH 423
H+
baseacid
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Brønsted-Lowry Concept of Acids and Bases Consider the reaction of NH3 and H2O.
– A conjugate acid-base pair consists of two
species in an acid-base reaction, one acid and
one base, that differ by the loss or gain of a
proton.
)aq(OH )aq(NH )l(OH )aq(NH 423
base acid
– The species NH4+ and NH3 are a conjugate
acid-base pair.
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Brønsted-Lowry Concept of Acids and Bases Consider the reaction of NH3 and H2O.
– The Brønsted-Lowry concept defines a species
as an acid or a base according to its function in
the proton-transfer reaction.
)aq(OH )aq(NH )l(OH )aq(NH 423
base acid
– Here NH4+ is the conjugate acid of NH3 and
NH3 is the conjugate base of NH4+.
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Brønsted-Lowry Concept of Acids and Bases Some species can act as an acid or a base.
)l(OH)aq(CO)aq(OH)aq(HCO 2
2
33
–H+
– An amphoteric species is a species that can act
either as an acid or a base (it can gain or lose a
proton).
– For example, HCO3- acts as a proton donor (an acid) in
the presence of OH-
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Brønsted-Lowry Concept of Acids and Bases Some species can act as an acid or a base.
– An amphoteric species is a species that can act
either as an acid or a base (it can gain or lose a
proton).
– Alternatively, HCO3 can act as a proton acceptor
(a base) in the presence of HF.
)aq(F)aq(COH)aq(HF)aq(HCO 323
H+
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Brønsted-Lowry Concept of Acids and Bases The amphoteric characteristic of water is important in
the acid-base properties of aqueous solutions.
– Water reacts as an acid with the base NH3.
)aq(OH)aq(NH)l(OH)aq(NH 423
H+
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Brønsted-Lowry Concept of Acids and Bases The amphoteric characteristic of water is important in
the acid-base properties of aqueous solutions.
– Water can also react as a base with the acid
HF.
)aq(OH)aq(F)l(OH)aq(HF 32
H+
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Brønsted-Lowry Concept of Acids and Bases In the Brønsted-Lowry concept:
2. Acids and bases can be ions as well as molecular
substances.
3. Acid-base reactions are not restricted to aqueous
solution.
4. Some species can act as either acids or bases
depending on what the other reactant is.
1. A base is a species that accepts protons; OH- is only
one example of a base.
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Lewis Concept of Acids and Bases The Lewis concept defines an acid as an electron
pair acceptor and a base as an electron pair donor.
– This concept broadened the scope of acid-
base theory to include reactions that did not
involve H+.
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Lewis Concept of Acids and Bases The reaction of boron trifluoride with ammonia is an
example.
– Boron trifluoride accepts the electron pair, so it is a
Lewis acid. Ammonia donates the electron pair,
so it is the Lewis base.
+ N
H
H
H:
::
: B
F
F
F
: :
::
::
::
: B
F
F
F
: :
::
:: N
H
H
H
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Relative Strength of Acids and Bases The Brønsted-Lowry concept introduced the idea of
conjugate acid-base pairs and proton-transfer reactions.
– The stronger acids are those that lose their
hydrogen ions more easily than other acids.
– Similarly, the stronger bases are those that hold
onto hydrogen ions more strongly than other
bases.
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Relative Strength of Acids and Bases The Brønsted-Lowry concept introduced the idea of
conjugate acid-base pairs and proton-transfer reactions.
– If an acid loses its H+, the resulting anion is now in
a position to reaccept a proton, making it a
Brønsted-Lowry base.
– It is logical to assume that if an acid is considered
strong, its conjugate base (that is, its anion) would be
weak, since it is unlikely to accept a hydrogen ion.
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Relative Strength of Acids and Bases Consider the equilibrium below.
– In this system we have two opposing Brønsted-
Lowry acid-base reactions.
– In this example, H3O+ is the stronger of the two
acids. Consequently, the equilibrium is skewed
toward reactants.
(aq)OHC(aq)OH 2323
)l(OH)aq(OHHC 2232 acid acidbase base
conjugate acid-base pairs
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Relative Strength of Acids and Bases Consider the equilibrium below.
– This concept of conjugate pairs is fundamental to
understanding why certain salts can act as acids or
bases.
(aq)OHC(aq)OH 2323
)l(OH)aq(OHHC 2232 acid acidbase base
conjugate acid-base pairs
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Self-ionization of Water Self-ionization is a reaction in which two like
molecules react to give ions.
In the case of water, the following equilibrium
is established.
)aq(OH)aq(OH )l(OH)l(OH 322
– The equilibrium-constant expression for this
system is:
22
3c
]OH[
]OH][OH[K
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Self-ionization of Water Self-ionization is a reaction in which two like
molecules react to give ions.
– The concentration of ions is extremely
small, so the concentration of H2O remains
essentially constant. This gives:
]OH][OH[K]OH[ 3c2
2
constant
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Self-ionization of Water– We call the equilibrium value for the ion product
[H3O+][OH-] the ion-product constant for water,
which is written Kw.
]OH][OH[K 3w
– At 25 oC, the value of Kw is 1.0 x 10-14.
– Like any equilibrium constant, Kw varies with
temperature.
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Self-ionization of Water– Because we often write H3O
+ as H+, the ion-
product constant expression for water can be
written:
]OH][H[Kw
– Using Kw you can calculate the concentrations of
H+ and OH- ions in pure water.
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Self-ionization of Water These ions are produced in equal numbers in pure
water, so if we let x = [H+] = [OH-]
– Thus, the concentrations of H+ and OH- in pure
water are both 1.0 x 10-7 M.
– If you add acid or base to water they are no longer
equal but the Kw expression still holds.
C 25at )x)(x(100.1o14
714100.1100.1x
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The pH of a Solution Although we can quantitatively describe the acidity of
a solution by its [H+], it is often more convenient to give acidity in terms of pH.
– The pH of a solution is defined as the negative
logarithm of the molar hydrogen-ion concentration.
]Hlog[pH
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The pH of a Solution For a solution in which the hydrogen-ion
concentration is 1.0 x 10-3, the pH is:
– Note that the number of decimal places in
the pH equals the number of significant
figures in the hydrogen-ion concentration.
00.3)100.1log( 3 pH
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The pH of a Solution In a neutral solution, whose hydrogen-ion concentration
is 1.0 x 10-7, the pH = 7.00.• For acidic solutions, the hydrogen-ion
concentration is greater than 1.0 x 10-7, so the
pH is less than 7.00.
• Similarly, a basic solution has a pH greater
than 7.00.
• Figure shows a diagram of the pH scale and the
pH values of some common solutions.
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The pH Scale
3
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A Problem to Consider A sample of orange juice has a hydrogen-ion
concentration of 2.9 x 10-4 M. What is the pH?
]Hlog[pH
)109.2log(pH4
54.3pH
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A Problem to Consider The pH of human arterial blood is 7.40. What is the
hydrogen-ion concentration?
)pHlog(anti]H[
)40.7log(anti]H[
M100.410]H[840.7
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The pH of a Solution A measurement of the hydroxide ion concentration,
similar to pH, is the pOH.
– The pOH of a solution is defined as the
negative logarithm of the molar hydroxide-
ion concentration.
]OHlog[pOH
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The pH of a Solution A measurement of the hydroxide ion concentration,
similar to pH, is the pOH.
– Then because Kw = [H+][OH-] = 1.0 x 10-14
at 25 oC, you can show that
00.14pOHpH
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A Problem to Consider An ammonia solution has a hydroxide-ion
concentration of 1.9 x 10-3 M. What is the pH of the solution?
You first calculate the pOH:
72.2)109.1log(pOH3
Then the pH is:
28.1172.200.14pH
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The pH of a Solution The pH of a solution can accurately be measured using
a pH meter
– Although less precise, acid-base indicators are
often used to measure pH because they usually
change color within a narrow pH range.
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The pH Meter
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THE COMMON–ION EFFECT
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SELECTIVE PRECIPITATION
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1. General chemistry by Ebbing Darell,3rd
edition
2. Essentials of physical chemistry by bahland tuli ,3 rd edition
3.https:// sites.google.com
4. https://www.mhhe.com
Presentation
of Lecture
Outlines,
16–49