Introduction to Acids and Bases
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Transcript of Introduction to Acids and Bases
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Introduction to
Acids and Bases
IB Chemistry Power Points
Topic 08
Acids and Bases
www.pedagogics.ca
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In aqueous solutions, a proportion of the water molecules dissociate;
The ions formed are H+ or positively charged hydrogen ions and negatively charged hydroxide ions (OH-)
Technically
2 H2O(l) H3O+(aq) + OH-
(aq)
Kw = [H + ][OH − ] = 1 x 10-14
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Some chemical compounds contribute additional H+ to make the solution more acidic. Other compounds remove H+ ions.
A compound that increases H+ is called an acid
Examples: HCl, H2SO4, HNO3, CH3COOH
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A compound that removes H+ ions from an aqueous solution is called a base. This reaction is called a neutralization.
Often this is done by adding OH- ions for example NaOH, KOH, Ca(OH)2. Soluble bases are called alkalis.
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Neutralization Reactions hydroxides
acid + base water + saltHCl + NaOH H2O + NaCl (aq)
• metal oxidesacid + base water + salt2 HCl + Cu2O H2O + CuCl2 (aq)
• ammoniaacid + base salt HCl + NH3 NH4Cl (aq)
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Three theories of acidsArrhenius (most common)
Bronsted-LowryLewis
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Arrhenius (most common): an acid dissociates to yield H+ and a base dissociates to yield
OH-
Hydrochloric acid H+ + Cl-
Sodium hydroxide Na+ + OH-
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Bronsted-Lowry:
an acid is a proton (H+) donor
and a base is a proton acceptor
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amphiprotic
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Lewis: An acid is an electron pair
acceptor
and a base is an electron pair donor
A dative covalent bond is formed
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Example of Lewis Acid
Lewis Acid Lewis
Base
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This is a common example that is not an obvious acid/base rxn
Boron trifluoride acts as a Lewis Acid. The boron has only 6 electron in valence shell so the lone pair of electrons forms a dative bond and fills up the valence shell of the boron
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IndicatorsAcids and bases are substances with specific physical and chemical properties.
We can determine if substances are acidic or basic by testing their reaction with indicators.
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Indicators are organic substances that change color in the presence of an acid or a base.Some common indicators
in acid in baseLitmus red bluePhenolphthalein colorless pinkMethyl orange red yellow
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Reactions of acidsReact with active metals (above copper in
reactivity series)2 HCl + Ca CaCl2 + H2
Reaction with carbonatesH2SO4 + Na2CO3 Na2SO4 + CO2 + H2O
Reaction with bicarbonatesHNO3 + NaHCO3 NaNO3 + CO2 + H2O
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Acid/base properties of Period 3 oxides (topic 3)Metal oxides Na2O and MgO react with water to
form hydroxides (basic solutions)Na2O + H2O 2 NaOH (aq)
Aluminum oxide is amphoteric (will react as a base with an acid or vice versa)
Al2O3 + 6 HCl 2 AlCl3 + 3 H2O
Other period 3 oxides (non-metal S, P, Cl oxides) react with water to form acidic solutions
SO3 + H2O H2SO4 (aq)
see page 15 in study guide
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Acid/base properties of Period 3 chlorides (topic 13)Chlorides across Period 3 become more acidic
across the periodNaCl (aq) is neutral
MgCl2 (aq) is weakly acidic
Chlorides of Al, Si, P, S and Cl2 react with water to produce HCl (aq) solutions
see Study guide page 16
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Strong Acids vs Weak Acids
The strength of an acid or base depends on how easily it dissociates in water.
The dissociation of an acid or base is an equilibrium.
HA(aq) H+(aq) + A-
(aq)
BOH(aq) B+(aq) + OH-
(aq)Strong acids or bases dissociate (ionize) easily – the equilibrium favors the ionic products : kc >> 1
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Strong vs Weak
When the strength of an acid or base is discussed, it is very important NOT to confuse “strength” with “concentration”
A 5M acid solution contains 5 mol of acid per dm3 but its strength is determined by how much of that acid is ionized.
Strong acids : HCl, H2SO4, HNO3 (mono vs diprotic)Strong bases : NaOH, KOH, Ba(OH) 2
Weak acids: CH3COOH, H2CO3, carbonic acid CO2(aq)Weak bases: NH3, ethylamine CH3CH2NH2
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Strong vs Weak
How to tellStrong acids and bases are mostly ionized and therefore solutions are good electrolytes (high conductivity). The pH of the solution can also be measured.
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What is the pH scale?pH is a measurement of hydrogen ion
concentration
It tells you how acids or basic (or alkaline) something is
Ranges from 0 (most acidic) to 14 (most basic
log[ ]pH H
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How does scale work?
The scale is logarithmic. As you go up or down, the concentration is changed by a power of ten
Example pH 3 is 100 times more concentrated than pH 5
neutral
pH 10 is 100 times less concentrated than pH 8
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Strong Acid
example HCl
HCl(aq) H+(aq) + Cl-(aq)
+ -[H ][Cl ]k = >> 1
[HCl]
• completely dissociated
• pH of 0.1 M soln = 1• strong electrolyte• reacts vigorously • note simplified “net
ionic” equation
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Weak Acid
example CH3COOHCH
3COOH (aq) H+
(aq) + CH3COO-
(aq)
+ -3
3
[H ][CH COO ]k = << 1
[CH COOH]
• partially dissociated• pH of 0.1 M soln = 2.9• weak electrolyte• reacts slowly