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CHAPTER 2

ACIDS AND BASES

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Arrhenius Acids and Bases

• Arrhenius proposed these definitions: – Acid: A substance that dissolves in water to

produce hydronium ions ions (H3O+).– Base: A substance that dissolves in water to

produce hydroxide ions (OH–).– Today we know that H+ reacts immediately

with a water molecule to give a hydronium ion H3O+.

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Arrhenius Acids and Bases

• When HCl, for example, dissolves in water, it reacts with water to give a hydronium ion and a chloride ion.

H2O(l) + HCl(aq) H3O+(aq) + Cl–(aq)

• We use curved arrows to show the change in position of electron pairs during this reaction.

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Arrhenius Acids and Bases

• With bases, the situation is slightly different.– Many bases are metal hydroxides such as

KOH, NaOH, Mg(OH)2, and Ca(OH)2.– These compounds are ionic solids and, when

they dissolve in water, their ions merely separate.

– Other bases are not hydroxides; these bases produce OH– by reacting with water molecules.

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Arrhenius Acids and Bases

– We use curved arrows to show the transfer of a proton from water to ammonia.

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Brønsted-Lowry Acids & Bases

• Acid: A proton donor.• Base: A proton acceptor.• Conjugate base: The species formed from an

acid when an acid donates a proton to a base.• Conjugate acid: The species formed from a

base when the base accepts a proton from an acid.– Acid-base reaction: A proton-transfer reaction.– Conjugate acid-base pair: Any pair of molecules or

ions that can be interconverted by transfer of a proton.

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Conjugate Acids & Bases

• We illustrate these relationships by the reaction of hydrogen chloride with water:

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Brønsted-Lowry Acids & Bases

• Brønsted-Lowry definitions do not require water as a reactant.

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Brønsted-Lowry Acids & Bases

• We use curved arrows to show the flow of electrons that occurs in the transfer of a proton from acetic acid to ammonia.

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Brønsted-Lowry Acids & Bases

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Brønsted-Lowry Acids & Bases

• Note the following about the conjugate acid-base pairs in Table 2.1:An acid can be : positively charged H3O+ , neutral H2CO3 negatively charged H2PO4

– A base can be :negatively charged OH–, Cl–, neutral NH3

Acids are classified depending on the number of protons that each may give up; examples are Monoprotic HCl Diprotic H2CO3

Triprotic H3PO4

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Brønsted-Lowry Acids & Bases

• Carbonic acid, for example, can give up one proton to become bicarbonate ion, and then the second proton to become carbonate ion.

Several molecules and ions appear in both the acid and conjugate base columns; that is, each can function as both an acid and as a base.

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Brønsted-Lowry Acids & Bases

There is an inverse relationship between the strength of an acid and the strength of its conjugate base.

– The stronger the acid, the weaker its conjugate base.– HI, for example, is the strongest acid in Table 2.1 and

its conjugate base, I–, is the weakest base in the table.

– CH3COOH (acetic acid) is a stronger acid than H2CO3 (carbonic acid); conversely, CH3COO– (acetate ion) is a weaker base than HCO3

– (bicarbonate ion).

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Acid and Base Strength

• Strong acid: One that ionizes completely or almost completely with water to form H3O+ ions.

• Strong base: One that ionizes completely or almost completely with water to form OH– ions.– Here are the six most common strong acids and the

four most common strong bases.

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Acid and Base Strength

• Weak acid: A substance that only partially dissociates in water to produce H3O+ ions.– Acetic acid, for example, is a weak acid; in water, only 4

out every 1000 molecules are converted to acetate ions.

• Weak base: A substance that only partially dissociates in water to produce OH– ions.– ammonia, for example, is a weak base.

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Acid-Base Reactions

• Acetic acid is incompletely ionized in aqueous solution.

• The equation for the ionization of a weak acid, HA, is

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pKa Values for Some Organic and Inorganic Acids

Larger the value of pKa ,weaker is the acid

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Acid-Base Equilibria

• To determine the position of equilibrium in an acid-base reaction:– Identify the two acids in the equilibrium; one on the left

and one on the right.– Use the information in Table 2.2 to determine which is the

stronger acid and which is the weaker acid.– Remember that the stronger acid gives the weaker

conjugate base, and the weaker acid gives the stronger conjugate base.

– The stronger acid reacts with the stronger base to give the weaker acid and weaker base.

– Equilibrium lies on the side of the weaker acid and the weaker base.

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Acid-Base Equilibrium

• Equilibrium in the following acid-base reaction lies to the right, on the side of the weaker acid and the weaker base.

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Structure and Acidity

• Most important factor in determining relative acidity of an organic acid is :

relative stability of the anion, A–, formed when the acid, HA, transfers a proton to a base.

• We consider these four factors:1. Electronegativity of the atom bonded to H in

HA.2. Resonance stabilization of A–.3. Inductive effect.4. Size and delocalization of charge in A–.

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Structure and Acidity

• Electronegativity of the atom bearing the negative charge.

• Within a period – The greater the electronegativity of the atom bearing the

negative charge, the more strongly its electrons are held.– The more strongly electrons are held, the more stable the anion

A–.– The more stable the anion A–, the greater the acidity of the acid

HA.