Acids and Bases - SharpSchoollpschurchill.ss5.sharpschool.com/UserFiles/Servers/Server...Common...
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Acids and Bases
Properties,
Reactions, pH, and
Titration
C-19
2017
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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.
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Acid’s Reaction with Metals
Metals: Dissolves; Problem: bridges, cars,
buildings
– Magnesium:
2HCl + Mg ---> H2 + MgCl2
– Iron:
2HCl + Fe ---> H2 + FeCl2
– Copper:
2HCl + Cu ---> H2 + CuCl2
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Common Acids
Fruits – citric acid
Milk – lactic acid
Vinegar – acetic acid
Soda pop – carbonic and
phosphoric acid
And lots more!!!!
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Properties of Bases
1. React with acids to form water and a salt.
2. Taste bitter.(Don’t try this)
3. Feel slippery (Don’t try this either).
4. Can be strong or weak electrolytes.
5. Change indicators (litmus → blue).
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Common Bases
Windex – ammonia
Baking soda – sodium
bicarbonate
Drain cleaner – NaOH
Milk of Magnesia –
Mg(OH)2
And more…..
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Organic acids
found in living things (fruits, etc)
contain -COOH a carboxyl group
weak acids are only slightly ionized to -COO-
Called carboxylic acids
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Mineral acids
from inorganic materials (rocks)
traditional acids - used industrially
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Common Industrial Acids
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.
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Nomenclature
Two basic types of acids: binary and oxyacids
1. Binary acids – 2 elements only
hydro + stem + ic acid
– HCl – hydrochloric acid
– HI – hydroiodic acid
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Nomenclature
2. Oxyacid names –
anion stem + ous (ite anions) NO2-1 (nitrite)
– HNO2 nitrous acid
Or anion stem + ic (ate anions) NO3-1 (nitrate)
– HNO3 nitric acid
– HClO4 perchloric acid
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Nature: Electrolytes are classified as Acids, Bases, or Salts
Acids - react with H2O and produce H+
– The H+ ion combines with water and forms H3O+
– called the “hydronium ion”
Bases – dissociate with H2O and produce OH-
Salts - Ionic combinations of metal/nonmetal
ions.
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Strong vs. Weak
STRONG electrolytes show complete
ionization in water (all ions); good conductors
– Soluble salts, SA, SB
– NaCl → Na+(aq) + Cl-(aq)
WEAK electrolytes show partial ionization in
water (mostly molecules); poor conductors
– WA, WB
– NH3 + H+ NH4+
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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%)
HA → H+1 + A-1
Weak acids produce few ions (less than 5%); are
dissolved intact as molecules.
HA H+1 + A-1
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Arrhenius Base
Bases dissociate and produce OH- ions.
Strong bases – 100% dissociation
– Group I and II hydroxides
Weak bases – less than 5% dissociation
– Ammonia, aniline, carbonates are not included.
– All other hydroxides are.
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Memorize the Strong Acids
HCl - hydrochloric
HBr – hydrobromic
HI - hydroiodic
H2SO4 - sulfuric
HClO4 – perchloric
HNO3 - nitric
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Memorize the Strong Bases
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
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Acid definitions
Bronsted Lowry
– Acids are proton donors
– Bases are proton acceptors
Acids and bases occur in conjugate pairs
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Come in Pairs
General equation
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
Acid + Base Conjugate + Conjugate
acid base
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Conjugate pairs
This is an equilibrium.
B(aq) + H2O(l) BH+(aq) + OH-(aq)
Base + Acid Conjugate acid +Conjugate base
NH3(aq)+H2O(l) NH4+(aq)+OH-(aq)
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In Bronsted-Lowry theory, bases do not require OH-
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
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Acid & Base Reactions
Neutralization Reaction:
– Acid + Base → “salt” + H2O (usually)
– “Salt” = general term for an ionic compound
– Example:
HCl + NaOH → NaCl + H2O
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Acid-Base reactions
Are equilibrium reactions (reversible)
Compare strength of the two acids (charts)
Equilib. shifts away from the stronger acid.
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HClO4 + H2O ⇆ H3O+ + ClO4
-
Acid + base ⇆ cong.acid + cong. Base
HClO4 is a stronger acid than H3O+ so….
Equilibrium shifts to the right →
– away from HClO4
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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
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Amphoteric substances
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-
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Lewis Theory
Lewis Acid – accepts an electron pair
Lewis Base – donates an electron pair
Not frequently used for chemists
Most general definition
(same G. Lewis that made e-dot diagrams)
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Lewis Acids and Bases
Lewis Acid
– A species (atom, ion or molecule) that is an
electron pair acceptor.
Lewis Base
– A species that is an electron pair donor.
base acid adduct
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Showing Electron Movement
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Focus On Acid Rain
CO2 + H2O → H2CO3
H2CO3 + H2O → HCO3- + H3O
+
HCO3- + H3O
+ → CO2 + H2O
3 NO2 + H2O → 2 HNO3 + NO
or
SO2 + H2O → H2SO3
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Acid Rain
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Acid rain
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
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Acid/Base Titration - a lab process
Basic Concepts:
– 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”.
– 5. Concentration of the “unknown” is calculated.
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7 Steps
1. Fill Burette with NaOH (known)
2. Place 20ml HCl in flask (unknown)
– 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.
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Titration calculation
Use the equation: Ma x Va = Mb x Vb
Example: 25 ml of HCl is neutralized by 20 ml of 0.5 M NaOH. Find conc. of HCl.
Solution: Ma = Mb x Vb / Va
Ma = 0.5 M x 20mL / 25mL = 0.4 M HCl
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Water
• Self ionization of water. (very small amount)
• H2O H+ + OH- • [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
• H2O H+ + OH-
• Kw 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-] • Kw/ [H+] = [OH-]
• 1 x 10-14/1 x 10-3 = [OH-] = 1 x 10-11
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Logarithms
• 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
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pH and pOH equations
• pH = -log[H+]
• pOH = - log [OH-]
• [H+] x [OH-] = 1 x 10-14
• pH + pOH = 14
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0 1 3 5 7 9 11 13 14
0 1 3 5 7 9 11 13 14
Basic
100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14
Basic 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14
Acidic Neutral
[OH-]
pH
[H+]
pOH