Transcript Acids and Bases - Livonia Public Schools

Acids and Bases
Properties
pH and
Titration
Nature: Electrolytes are classified
as Acids, Bases, or Salts
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Acids - react with H2O and produce H+
Bases - react with H2O and produce OHSalts - Ionic combinations of metal/nonmetal
ions.
Strong vs. Weak
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STRONG electrolytes show complete
ionization in water (all ions); good conductors
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Soluble salts, SA, SB
NaCl → Na+(aq) + Cl-(aq)
WEAK electrolytes show partial ionization in
water (mostly molecules); poor conductors
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WA, WB
NH3 + H+
NH4+
Properties of acids
1. Taste Sour (don’t try this except with foods).
2. Are electrolytes (conduct electricity).
–
Some are strong, some are weak.
3. Change indicator colors. (litmus →red).
4. React with metals to form hydrogen gas.
5. React with hydroxides to form water and a
salt.
Acid’s Reaction with Metals
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Metals: Dissolves; Problem: bridges, cars,
buildings
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Magnesium:
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Iron:
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2HCl + Mg ---> H2 + MgCl2
2HCl + Fe ---> H2 + FeCl2
Copper:
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2HCl + Cu ---> H2 + CuCl2
Properties of Bases
1.
2.
3.
4.
5.
React with acids to form water and a salt.
Taste bitter.(Don’t try this)
Feel slippery (Don’t try this either).
Can be strong or weak electrolytes.
Change indicators (litmus → blue).
Common Acids
Fruits – citric acid
 Milk – lactic acid
 Vinegar – acetic acid
 Soda pop – carbonic and
phosphoric acid
And lots more!!!!
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Common Bases
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Windex – ammonia
Baking soda – sodium
bicarbonate
Drain cleaner – NaOH
Milk of Magnesia –
Mg(OH)2
And more…..
Organic acids
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found in living things (fruits, etc)
contain -COOH a carboxyl group
weak acids are only slightly ionized to -COOCalled carboxylic acids
Mineral acids
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from inorganic materials (rocks)
traditional acids - used industrially
Common Industrial Acids
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Sulfuric acid - H2SO4 – petroleum, fertilizer, metallurgy,
paper, paints,batteries, etc
Nitric acid – HNO3 – explosives, rubber, plastics,
pharmaceuticals, etc.
Phosphoric – H3PO4 – fertilizer, flavoring agent,
detergents, etc.
Hydrochloric – HCl – pickling metal, cleaning,
chlorination (pools)
Acetic Acid – CH3COOH – plastics, food supplements,
etc.
Nomenclature
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Two basic types of acids
1. Binary acids – 2 elements only
 hydro + stem + ic acid
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HCl – hydrochloric acid
HI – hydroiodic acid
Nomenclature
2. Oxyacid names –
 anion stem + ous (ite anions)
–
HNO2 nitrous acid
Or anion stem + ic (ate anions)
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HNO3 nitric acid
HClO4 perchloric acid
Aqueous acids
Arrhenius definition: acids ionize in water to form H+ ions - are polar covalent compounds and all have H.
- may ionize in more than 1 step. (ex H2SO4)
 Strong acids show complete ionization (100%)
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HA → H+1 + A-1
Weak acids produce few ions (less than 5%); are
dissolved intact as molecules.
HA → H+1 + A-1
Arrhenius Base
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Bases dissociate and produce OH- ions.
Strong bases – 100% dissociation
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Group I and II hydroxides
Weak bases – less than 5% dissociation
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Ammonia, aniline, carbonates are not included.
All other hydroxides are.
Memorize the Strong Acids
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HCl - hydrochloric
HBr – hydrobromic
HI - hydroiodic
H2SO4 - sulfuric
HClO4 – perchloric
HNO3 - nitric
Memorize the Strong Bases
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NaOH - sodium hydroxide
KOH - potassium hydroxide
LiOH – lithium hydroxide
RbOH - rubidium hydroxide
Ba(OH)2 – barium hydroxide
Sr(OH)2 – strontium hydroxide
Ca(OH)2 - calcium hydroxide
Mg(OH)2 – magnesium hydroxide
Acid definitions
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Bronsted Lowry
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Acids are proton donors
Bases are proton acceptors
Acids and bases occur in conjugate pairs
Come in Pairs
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General equation
HA(aq) + H2O(l)
Acid + Base
H3O+(aq) + A-(aq)
Conjugate acid +
Conjugate base
Conjugate pairs
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This is an equilibrium.
B(aq) + H2O(l)
BH+(aq) + OH-(aq)
Base + Acid
Conjugate acid +Conjugate base
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NH3(aq)+H2O(l)
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NH4+(aq)+OH-(aq)
In Bronsted-Lowry theory, bases do
not require OH
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Bases are able to accept protons
Allows ammonia and carbonate ions to be
considered bases, others as well.
NH3 + H+ → NH4+
Base + H+ → Conjugate acid
Most accepted theory
Acid & Base Reactions
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Neutralization Reaction:
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Acid + Base  “salt” + H2O (usually)
“Salt” = general term for an ionic compound
Example:
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HCl + NaOH  NaCl + H2O
Acid-Base reactions
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Are equilibrium reactions (reversible)
Compare strength of the two acids (charts)
Equilib. shifts away from the stronger acid.
HClO4 + H2O ⇆ H3O+ + ClO4
Acid + base ⇆ cong.acid + cong. Base
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HClO4 is a stronger acid than H3O+ so….
Equilibrium shifts to the right →
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–
away from HClO4
Protons are Hydrogen ions
Monoprotic acids have one proton to donate
ex. HCl
Diprotic acids have two protons to donate
ex. H2SO4 (one step at a time)
Polyprotic – two or more protons to donate
ex. H3PO4
Amphoteric substances
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Substances which can either accept or donate
a proton.
Water is an example
H2O + H+ → H3O+ (water as a base)
H2O → H+ + OH- (water as an acid)
Other examples are NH3 and HSO4-
Lewis Theory
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Lewis Acid – accepts an electron pair
Lewis Base – donates an electron pair
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Not frequently used for chemists
Most general definition
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(same G. Lewis that made e-dot diagrams)
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17-9 Lewis Acids and Bases
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Lewis Acid
–
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A species (atom, ion or molecule) that is an
electron pair acceptor.
Lewis Base
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A species that is an electron pair donor.
base
acid
adduct
Showing Electron Movement
Focus On Acid Rain
CO2 + H2O → H2CO3
3 NO2 + H2O → 2 HNO3 + NO
H2CO3 + H2O → HCO3- + H3O+
or
HCO3- + H3O+ → CO2 + H2O
SO2 + H2O → H2SO3
Acid Rain
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Acid rain
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Gases like sulfur dioxide and nitrogen dioxide
are produced from burning coal, oil, and other
fuels.
These gases react with water vapor in the
atmosphere to form acids.
Acid rain can be stopped with govt. regulations.
Less in US/Canada now, but more in
China/India
Acid/Base Titrations cont.
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Basic Concepts:
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1. Acids & bases neutralize each other
2. From the balanced equation, the number of
moles needed of the “known” reactant & the
“unknown” reactant are given.
3. An indicator is selected based on the strength of
the “known” reactant.
4. The indicator will change color when the “known”
reactant equals the “unknown”.
7 Steps
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1. Fill Burette with NaOH (known)
2. Place 20ml HCl in flask (unknown)
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The amount may be different, but record
3. Place indicator in HCl
4. Slowly add NaOH until the endpoint is reached (color change).
5. Record amount of NaOH used (let’s pretend 19.9ml)
6. Use the factor label method to find the number of moles of
NaOH.
7. Look at the balanced equation to determine the ratio of moles
between the “Known” NaOH & “unknown” HCl.
Titration calculation
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Use the equation: Ma x Va = Mb x Vb
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Example: 25 ml of HCl is neutralized by 20 ml
of 0.5 M NaOH. Find conc. of HCl.
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Solution: Ma = Mb x Vb / Va
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Ma = 0.5 M x 20mL / 25mL = 0.4 M HCl
Water
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Self ionization of water. (very small amount)
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H2O  H+ + OH-
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[H+ ] = [OH-] = 1 x 10-7M
A neutral solution.
In water: Kw = [H+ ] x [OH-] = 1 x 10-14
Kw is called the ion product constant.
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Ion Product Constant
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H2O
H+ + OHKw is constant 1 x 10-14
If [H+] > 10-7 then [OH-] < 10-7 (acidic)
If [H+] < 10-7 then [OH-] > 10-7 (basic)
If we know one, we can determine the other.
If [H+] = 1x 10-3 Find [OH-]
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Kw/ [H+] = [OH-]
1 x 10-14/1 x 10-3 = [OH-] = 1 x 10-11
Logarithms
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Powers of ten. A shorthand form
pH = -log[H+]
in neutral pH = - log(1 x 10-7) = 7
in acidic solution [H+] > 10-7
pH < -log(10-7)
pH < 7
in base pH > 7
pH and pOH equations
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pH = -log[H+]
pOH = - log [OH-]
[H+] x [OH-] = 1 x 10-14
pH + pOH = 14
[H+]
100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14
pH
0
1
Acidic
14 13
10-14 10-13
3
11
5
7 9
Neutral
9
7 5
11
3
13
14
Basic
1
0
pOH
10-11 10-9Basic
10-7 10-5 10-3 10-1 100
[OH-]