Chapter 14
Acids and Bases
Acid/Base Theories
• Arrhenius Theory– Acids produce H+ ions in solution– Bases produce OH- ions in solution– Downside
• Must be in solution and must have those ions
• Bronsted-Lowry Theory– Acids are H+ donors (Proton donors)– Bases are H+ acceptors (Proton acceptors)
Vocabulary
• H+ is the hydrogen ion– Just a proton
• H3O+ is a hydronium ion
– It is the way H+ exists in water
• Water accepts a hydrogen ion and becomes H3O+
• Either way is fine the first is just easier
Conjugates Acid/Base Pairs
• Conjugate Acid is formed when a base gains a proton
• Conjugate Base is what remains after an acid donates a proton
• Ex – HNO3 + H2O H3O+ + NO3-
General Form for Acids and Bases
HA + B A- + BH+
• HA is an acid
• B is a base
• A- is the conjugate base– Just the negative ion of the acid
• BH+ is the conjugate acid– Just the base plus a hydrogen
Strong and Weak Acids
• Strong acids completely ionize in solution– Nitric, Perchloric, Sulfuric, Hydrochloric,
Hydrobromic, Hydroiodic– Weak Conjugate bases
• Weak Acids only partially ionize in solution– Every other acid– Equilibrium is established in the ionization– Weak acids have Ka values – Strong Conjugate bases
Acid Dissociation Constant, Ka
HA(aq) + H2O A- + H3O+
• For weak acids equilibrium is established
• Equilibruim constant is Ka
• Ka=[A-][H3O+]/[HA]
• Values tend to be small– Because CB is fairly strong
• Strong acids do not have Ka values.
Acid Terms
• Monoprotic – One acidic hydrogen
• Polyprotic – Many acidic hydrogens
• Diprotic – Two acidic hydrogens
• Triprotic – Three acidic hydrogens
• Oxyacid – Acid that has an acidic hydrogen attached to an oxygen
• Organic Acid – Acid that has the acidic hydrogen attached to the carboxyl group
Water As An Acid and Base
• Water is amphoteric – Both an acid and base
H2O + H2O H3O+ + OH-
• Equilibrium system that always has the same value
• Called Kw
Autoionization of Water, Kw
H2O + H2O H3O+ + OH-
• Kw = [H3O+] [OH-]
• In pure water the products have the same concentration, 1.00x10-7M
• Value of Kw = 1.00x10-14 at 25ºC
• Concentrations can change is acid or base is added
Acid, Base, or Neutral
• If the concentration of H+ = OH-
– Neutral
• If the concentration of H+ > OH-
– Acidic
• If the concentration of H+ < OH-
– Basic
What is the hydrogen ion concentration when the hydroxide ion concentration is 1.00x10-5M
What is the hydroxide ion concentration when the when the concentration of nitric acid is 0.0010M?
Homework
• P. 704 30, 32, 33, 35, 39ab,40ab
Logarithms
• The logarithm of a number to a given base (commonly 10) is the power or exponent to which the base must be raised in order to produce the number.
• SAY WHAT!
Examples
• If your question says
log 100 = x
It is saying to what power must 10 be raised to equal 100
10x = 100
102 = 100
So x = 2
Examples
log 1 = x x = 0
log 10 = xx = 1
log 1000 = x
x = 3
log 1x106 = x
x = 6
Examples
• Logarithms can also be used for numbers smaller than 1
log 0.1 = x
x = -1
log 0.01 = x
x = -2
log 1x10-5
x = -5
Examples
• If they are not easy to calculate you can do it on your calculator
log 15 = x
You can approximate it between . . .
–Type log 15 on your calculator
x = 1.18
Examples
• If your question says
log x = 7
It is saying 10 to the 7th power is what number
107 = 1x107
log x = 3
x = 1000
log x = -3
x = 0.001
Tougher Examples
log 234 = x
x = 2.37
log x = -3.3
x = 5.0x10-4
-log 9.1x10-5 = x
x = 4.0
-log x = 12.1
x = 7.9x10-13
pH
• Negative logarithm of the concentration of hydrogen ions in solution
pH = -log [H+]
• pH means power of Hydrogen
• Measures how acidic or basic a solution is
• pH scale typically goes from 0 to 14
• pH < 7 Acidic
• pH > 7 Basic
• Highly acidic = low pH
• Highly basic = high pH
Significant Figures and pH
• Digits after the decimal are the only ones that are significant in pH values
pH = 4.44
2 Significant Figures
pH = 10.874
3 Significant Figures
• If your [H+] is 0.088 M your pH is 1.06
Determining [H+] in solution
• The concentration of a strong acid is equal to the H+ concentration.
• 0.010 M HCl has an [H+] of 0.010 M
• To obtain the [H+] for weak acids you must use equilibrium
– Need Ka data
• Discuss hydroxide later
Other Info
Turn Kw into a log equation
pH + pOH = 14.00
pH’s You MUST Know
• When the [H+] is _______ the pH is _____0.10 M
1.000.0010M
3.001.0x10-6M
6.001.0x10-10M
10.00
• Find the pOH, [H+], [OH-] of lemon juice that has a pH of 2.48
• Determine the pH of 0.150M HCl
• Determine the pH of 2.3x10-3 M Hydrocyanic Acid HCN. Ka = 6.2x10-10
Homework
• Page 705 #’s 45,47,50,52,53,58
Mixtures of Weak Acids
• When there are mixtures of weak acids in solution determining the pH could be a difficult problem.
• However
– The acid with the largest Ka will control the pH of the solution
• Solutions of 0.10M HF (Ka = 7.2x10-4) and 0.10M HCN (Ka = 6.2x10-10) are mixed A) Which acid will control the pH of the solution? Why? B) What is the pH of the resulting solution.
Percent Dissociation
• Ratio of the concentration of the dissociated ions to the initial concentration
• Found just like the doing the 5% check
• Can be used to find pH and Ka
100*%centrationInitialCon
rationIonConcentHonDissociati
• A 0.25M solution of HClO is 20.% dissociated. A) What is the pH? B) What is the Ka value of the acid?
Strong Bases
• Strong bases are any compound containing the hydroxide ion
– Group 1 hydroxides are very soluble
– Group 2 less soluble but still strong
• Group 2 have two hydroxides per mole
–Be Careful
• Determine the pH of 0.022M Sr(OH)2
Weak Bases
• Organic Bases are weak bases. (Ammonia)
– Contain Nitrogen (Amines)
• CH3NH2 – Methylamine
– Must have a lone pair of electrons
– The hydroxide will come from water
(CH3)3N + H2O
Cont.
• Weak bases have Kb values
B + H2O BH+ + OH-
• The conjugate acid of a weak base is stronger than water.
][
]][[
B
OHBHKb
• Determine the pH of 15.0 M ammonia. The Kb is 1.8x10-5
Homework
• Page 705 #’s 62,63a,72,73,77ab,87
Polyprotic Acids
• Acids with more than one acidic hydrogen• Dissociate in a “stepwise” process
H3PO4 H+ + H2PO4- Ka1 =7.5x10-3
H2PO4- H+ + HPO4
-2 Ka2 =6.2x10-8
HPO4-2 H+ + PO4
-3 Ka3 =4.8x10-13
• Ka1 > Ka2 > Ka3
• Successive dissociation do not effect pH (except sulfuric acid)
Determine the pH of 5.0M H3PO4 (Ka=7.5x10-3) and the [H2PO4
-], [HPO4-2], [HPO4
-3]
Sulfuric Acid
• Sulfuric acids has two dissociations
– The first is strong
– The second is weak
• The second dissociation is quite strong, but it is not complete
H2SO4 H+ + HSO4- (Strong)
HSO4- H+ + SO4
-2 Ka = 1.2x10-2
Continued
• When you write sulfuric acid in net ionic equations only use the first equation
H2SO4 H+ + HSO4-
• The second dissociation only needs to be considered in dilute solutions.
• A 1.0M solution of sulfuric acid will have a lower pH than a 1.0M solution of HCl
Acid / Base Properties of Salts
• Some salts have acid base properties
– Make a solution have a pH below or above 7
• Some have no acid base properties
– Make a solution with a pH of 7
Neutral Salts
• Contain the following
• A metallic ions– (Except Aluminum)
• A conjugate base of a strong acid– Such AS . . . .
• Conjugate base of strong acids are weaker than water.
Basic Salts
• Contain the following
• A metallic ion
• A conjugate base of a weak acid– Such AS . . .
• Conjugate base of weak acids are stronger than water.
Acidic Salts
• Contain the following
• A conjugate base of a strong acid
• A conjugate acid of a weak base– Such AS . . .
• Conjugate acid of weak bases are stronger than water.
Acidic Aluminum Ions
• Highly charged cations polarizes the O-H bond in water
• Hydrogens in water become acidic
• Ion becomes hydrated
• Al(H2O)6+3
OMGosh I Have Both Types of Ions
• If a salt has acidic and basic ions – Such AS . . .
• If Ka > Kb the solution is acidic
• If Ka < Kb the solution is basic
• If Ka = Kb the solution is neutral
Converting Ka to Kb
• If you know an acids Ka value you can find its Kb value as a salt.
• If you know a bases Kb value you can find its Ka value as a salt
Kw = Ka * Kb
Calculate the pH of 0.33 M NaHCO3
Homework
• P 707 #’s 94,95b, 98,99,102,105a, 107,112
Effect of Structure On Acids
• Different bonding patterns lead to differences in the strength of acids
• Two different types of acids
– Hydrohalic
– Oxyacids
Hydrohalic Acids
• Acids that contain a hydrogen and a halogen
• The weaker the bond the stronger the acid
Bond Enthalpy
H—F 567
H—Cl 433
H—Br 366
H—I 299
Oxyacids
• Tend to be of the form H – O – X
• The greater the electronegativity of X the stronger the acid
• Pulls electron density away and weakens the H – O bond
• Which is a stronger acid?
• H – O – Cl or H – O – Br – H – O – Cl because Cl has a greater E.N.
Oxyacids
• Bases tend to be of the form X – O – H too
• NaOH
• Why
• Metal has a low electronegativity and the O – H bond is not weakened
Oxyacids
• Acids with the same X element can have different number of oxygens
• As the number of oxygens increases so does the strength of the acid
• The extra oxygens weaken the H – O bond
Oxides
• Nonmetalic oxide in water gives an acid
CO2 + H2O H2CO3
• Acidic anhydrides (Acid w/o water)
• Nonmetal must keep its oxidation #
• Metallic oxides in water gives a base
Na2O + H2O 2NaOH
• Basic anhydrides
Lewis Acids and Bases
• Lewis acids are electron pair acceptors
• H+ for example has no electrons
• Lewis bases are electron pair donors
• NH3 has an electron pair to share
Homework
• P 708 #’s 113,114,115ab,116ab,119
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