1 the chemistry of acids and bases. 2 3 some properties of acids þ react with certain metals to...
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1The Chemistry of Acids and Bases
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Some Properties of Acids
þ React with certain metals to produce hydrogen gas.
þ React with carbonates and bicarbonates to produce carbon dioxide gas
þ Taste sour
þ Corrode metals
þ Electrolytes
þ React with bases to form a salt and water
þ pH is less than 7
þ Turns blue litmus paper to red “Blue to Red”
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Some Properties of Bases
Taste bitter, chalky
Are electrolytes
Feel soapy, slippery
React with acids to form salts and water
pH greater than 7
Corrosive
Turns red litmus paper to blue “Basic Blue”
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Acid and Bases
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Acid and Bases
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Acid and Bases
8Indicators
• Indicators are dyes that can be added in small amounts that will change color in the presence of an acid or base.
• Some indicators only work in a specific range of pH
9Examples: Indicators
• Litmus paper• Phenolpthalein• Bromothymol blue• Methyl Orange
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Universal Indicator
• Universal indicator is a pH indicator composed of a blend of several compounds that changes colour over a wide range of pH values from 0-14.
Cabbage juice
pH paper
12• A universal indicator is typically composed of water, methanol, proan-1-ol, phenolpthalein, sodium salt, methyl red, bromothylmol blue, monosodim salt and thymol blue monosodium salt.
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pH• Measure of H+ ions in
solution. • High [H+] ions = more
acidity and low pH.• pH meter tests the
voltage of the electrolyte
• Converts the voltage to pH
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The pH scale is a way of expressing the strength of acids and bases. Under 7 = acid
7 = neutral Over 7 = alkaline (base)
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pH of Common Substances
16pH is a logarithmic function
• pH 2 ____ more acidic than a pH of 3• pH 2 ____ more acidic than a pH of 4• pH 2 ____ more acidic than a pH of 5
pH = - log [H+](Remember that the [ ] mean
Molarity)
17Strong and Weak Acids/Bases
The strength of an acid (or base) is determined by the amount of IONIZATION (dissociation).
18Strong Acids
• A strong acid is one that completely 100% ionizes in water.
Example:
HCl (aq) + H2O (l) -H+ (aq) + Cl- (aq)
HNO3, HCl, H2SO4 and HClO4 are among the only known strong acids. Most acids are weak.
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• Strong Base: 100% dissociated into its respective ions water.
NaOH (aq) ---> Na+ (aq) + OH- (aq)
Strong Bases
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• Weak acids ionize less than 100% in water.
Weak Acids
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HClO (aq) H+(aq) + ClO-
(aq)
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Weak Acid Dissociation(only partially
ionizes)
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Weak base: less than 100% ionized in waterOne of the best known weak bases is ammonia
NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)
Weak Bases
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Weak Bases
25Conductivity
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More About Water
In pure water there can be AUTOIONIZATION
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Equilibrium constant for water = Kw
It can be seen from the above equation that H3O+ and
OH¯ concentrations are in the molar ratio of one-
to-one. This means that [H3O+] = [OH¯].
Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC
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Acid/Base definitions
Arrhenius (Swedish chemist)
29Arrhenius acid is a substance that produces H+ in water
Arrhenius base is a substance that produces OH- in water
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POLYPROTIC ACIDS
Monoprotic
Diprotic
Triprotic
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Limitations to Arrhenius’ Theory
An acid is expected to be an acid in any solvent.
But that’s not the case nowadays. For example HCL acts as an Arrhenius acid when dissolved in water. However when HCL is dissolved in benzene there is no dissociation. This is against Arrhenius theory; Arrhenius states that dissociation occurs in any aqueous solution. The properties of acid and bases play a critical role. Water does not have to be the only solvent.
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Limitations to Arrhenius’ Theory
2- In Arrhenius theory all salts produced in the neutralization reaction should produce solutions that are neither acidic nor basic. But there are some exceptions against this theory. For example if equal amounts of HCl and ammonia react, the solution is slightly acidic. If equal amounts of acetic acid and sodium hydroxide react, the resulting solution is basic. Arrhenius theory does not include any explanation for this.
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Limitations to Arrhenius’ Theory
3- The proton H+ liberated from the acid actually combines with water to form hydronium ions in a water solution. Therefore this reaction: H2O + H+ H3O+ Occurs most of the time.
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Limitations to Arrhenius’ Theory
4- The need for hydroxide as the base led Arrhenius to propose the formula NH4OH as the formula for ammonia in water. This led to the misunderstanding that NH4OH is the actual base. But the actual base is NH3.
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Acid/Base DefinitionsBrønsted – Lowry
Acids – any species that can donate a proton (H+ ions) in solution ‘proton donor’
Bases – any species that accepts a proton (H+ ions) in solution proton acceptor
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A Brønsted-Lowry acid is a proton donorA Brønsted-Lowry base is a proton acceptor
acidconjugate
basebase conjugate
acid
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ACID-BASE THEORIES
The Brønsted definition means NH3 is a BASE in water — and water is itself an ACID
BaseAcidAcidBaseNH4
+ + OH-NH3 + H2O
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Conjugate Pairs
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Learning Check!
Label the acid, base, conjugate acid, and conjugate base in each reaction:
HCl + OH- Cl- + H2O
H2O + H2SO4 HSO4- + H3O
+
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Amphoteric
• A substance that can act as an acid in one situation and a base in another.
• Ex• H20 + HCl H3O+ + Cl-• H20 + NH3 NH4 + OH-
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ACID-BASE REACTIONSTitrations
H2C2O4(aq) + 2 NaOH(aq) --->
acid base
Na2C2O4(aq) + 2 H2O(liq)
Carry out this reaction using a TITRATION.
Oxalic acid,
H2C2O4
42Setup for titrating an acid with a base
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Titration1. Add solution from the buret.2. Reagent (base) reacts with
compound (acid) in solution in the flask.
3. Indicator shows when exact stoichiometric reaction has occurred. (Acid = Base)
This is called NEUTRALIZATION.