Transcript Acids and Bases Chapter 19 Describing Acids and Bases Mini-Project Work with a partner.

Acids and Bases
Chapter 19
Describing Acids and Bases
Mini-Project
Work with a partner. Organize the following
formulas into two groups with four formulas in
each group:
HNO3, NaOH, H2SO4, H2CO3, Ca(OH) 2, KOH,
H8PO4, Mg(OH) 2
One way to organize them into groups is:
Group One
Group Two
HNO3
NaOH
H2SO4
Ca(OH) 2
H2CO3
KOH
H8PO4
Mg(OH) 2
Group One formulas represent acids.
Group Two formulas represent bases.
ACIDS
•
•
•
•
Electrolytes in solution
Taste sour (lemon, vinegar)
React with metal (corrosion)
React with carbonates (makes bubbles of
CO2
• Turns blue litmus RED
• In Water forms Hydrogen ION
HCl
Water
H+ + Cl-
BASES
•
•
•
•
•
Electrolyte in solution
Taste Bitter (soap, tonic water)
Feel Slippery (soap)
Turns Red Litmus Blue
React with Acids to make water
NaOH
Water
Na+ + OH-
Why are Some Solutions Acid &
Others Base?
• Acid solutions contain more H+ ions than
OH- ions.
• Base solutions contain more OH- ions than
H+ ions.
• Water is the standard for Acid/Base and is
defined as NEUTRAL
• Water has equal amounts of H+ and OHions
Arrhenius Model of Acids/Bases
• Substance is an acid if it contains
hydrogen and dissociation causes
hydrogen ions to form in solution
• Substance is a base if it contains a
hydroxide and dissociates to produce
hydroxide ions in solution
Bronsted-Lowry Model
• Acid is a proton (hydrogen ion) donor
• Base is a proton (hydrogen ion) receptor
• This is a broader definition than Arrhenius
model because there are substances that
cause donation or reception without
having hydrogen in them.
Example
• When an acid dissolves in water, it donates an
H+ ion to a water molecule forming H3O+.
• The water molecule acts as a base and accepts
the H+ ion
HX + H2O ⇄ H3O+ + X– Conjugate acid = species produced when a
base accepts a hydrogen ion from an acid
– Conjugate base = species produces when an
acid donates a hydrogen ion to a base
– Conjugate base pair = 2 substances related to
each other by donating and accepting a single
hydrogen ion
Historical views on acids
• O (e.g. H2SO4) was originally thought to cause
acidic properties. Later, H was implicated, but
it was still not clear why CH4 was neutral.
• Arrhenius made the revolutionary suggestion
that some solutions contain ions & that acids
produce H3O+ (hydronium) ions in solution.
Ionization
+
H
H
HO
+ Cl
Cl H + O
H
H
• The more recent Bronsted-Lowry concept is
that acids are H+ (proton) donors and bases
are proton acceptors
The Bronsted-Lowry concept
• In this idea, the ionization of an acid by water
is just one example of an acid-base reaction.
H
Cl H
acid
+
O
H
base
+
H
HO
+
Cl
H
conjugate acid conjugate base
conjugate acid-base pairs
• Acids and bases are identified based on
whether they donate or accept H+.
• “Conjugate” acids and bases are found on the
products side of the equation. A conjugate
base is the same as the starting acid minus H+.
Practice problems
Identify the acid, base, conjugate acid,
conjugate base, and conjugate acid-base pairs:
HC2H3O2(aq) + H2O(l)  C2H3O2–(aq) + H3O+(aq)
acid
base
conjugate base conjugate acid
conjugate acid-base pairs
OH –(aq) + HCO3–(aq)  CO32–(aq) + H2O(l)
base
acid
conjugate base conjugate acid
conjugate acid-base pairs
Answers: question 18
(a) HF(aq) + SO32–(aq)  F–(aq) + HSO3–(aq)
acid
base
conjugate base conjugate acid
(b)
conjugate acid-base pairs
CO32–(aq) + HC2H3O2(aq)  C2H3O2–(aq) + HCO3–(aq)
base
acid
conjugate base conjugate acid
(c)
conjugate acid-base pairs
H3PO4(aq) + OCl –(aq)  H2PO4–(aq) + HOCl(aq)
acid
base
conjugate base conjugate acid
conjugate acid-base pairs
More
• Amphoteric = substances like water that
can act like either an acid or a base
Monoprotic and Polyprotic Acids
• Monoprotic – acids based on formula that can
donate only one hydrogen ions
– CH3COOH + H2O ⇄ H3O+ + CH3COO-
• Polyprotic – acids that can donate multiple
hydrogen ions
– H3PO4 + H2O ⇄ H3O+ + H2PO4+
– H2PO4+ + H2O ⇄ H3O+ + HPO4+2
– HPO4+2 + H2O ⇄ H3O+ + PO4+3
• Anhydride = oxides that can become acids or
bases by adding elements contained in water
Acid Rain
• Acid rain comes from rain collecting
gasses from the air to create acids:
– Carbon Dioxide = carbonic acid
– Sulfur oxides = sulfuric acid
– Nitrogen oxides = nitric acid
• Damages statues, buildings, kills forests,
kills fish
Acids and Bases in Solution
Chapter 19.2
Acid/Base Strength
• In strong acids, almost all molecules ionize.
• In weak acids, fewer molecules ionize.
Conjugate Pairs Strength
• If an acid is a strong acid, its conjugate
pair base is a weak base
• Why?
– If HX is strong acid, it ionizes completely.
– The conjugate base must be a weak base
because it has a greater attraction to the H+
than HX
– The reaction equilibrium lies far to the right of
the equation.
Conjugate Pair Strength
• For a weak acid, the equation equilibrium
lies to the right (reactant side)
– Conjugate base (Y-) has a stronger attraction
for the H+ ion than the base H2O
– HY + H2O
H3O+ + Y-
Acid Ionization Constants
• An “ionization constant” is the tendency of
an item to make ions in solution.
– Higher the constant, the higher the amount of
ions.
– Acid ionization constant is value of the
equilibrium constant expression for a weak
acid
– Value Ka indicates whether reactants or
products are favored at equilibrium
– Weak acids have low Ka values
Acid Ionization Constants
Acid Ionization Constants
Substance
Formula
Ka
Acetic Acid
HC2H3O2
1.7 x 10-5
Boric Acid
H3BO3
5.9 x 10-10
H2CO3
4.3 x 10-7
HCO3-
4.8 x 10-11
H2S
8.9 x 10-8
HS-
1.2 x 10-13
Hypochlorous Acid
HClO
3.5 x 10-8
Nitrous Acid
HNO2
4.5 x 10-4
H2C2O4
5.6 x 10-2
HC2O4-
5.1 x 10-5
H3PO4
6.9 x 10-3
HSO3
6.3 x 10-8
Carbonic Acid
Hydrogen Sulfide
Oxalic Acid
Phosphoric Acid
Base Ionization Constant
• Same Basic Principle as Acid
– Measures OH- concentrations
pH Scale
Chapter 19.3
Ionization Constant for Water
• The ionization constant for water is:
– 1.0 x 10-14
– [Ka][Kb]
– = [1.0x10-7] [1.0 x 10-7]
• Experiments show that the product of [H+]
and [OH-] always equals 1.0 x 10-14 at
298°C
• pH scale is a way of showing this
relationship of ionization constants
The pH Scale
•
•
•
•
pH stands for ‘per Hydrion’
Low pH is Acid
High pH is base
Water is neutral (7.0)
pH
• There are many ways to consider acids and
bases. One of these is pH.
• [H+] is critical in many chemical reactions.
• A quick method of denoting [H+] is via pH.
• By definition pH = –log [H+], [H+] = 10-pH
• The pH scale, similar to the Richter scale,
describes a wide range of values
• An earthquake of “6” is 10 as violent as a “5”
• Thus, the pH scale condenses possible
values of [H+] to a 14 point scale (fig. 2, p370)
• Also, it is easier to say pH=7 vs. [H+]=1x10–7
pH
• pH = -log [H+]
– [H+] = 10-pH
• pOH = -log [OH-]
• pH + pOH = 14
Calculations with pH
Q: What is the pH if [H+]= 6.3 x 10–5?
pH = –log [H+]
‘(-)’, ‘log’, ‘6.3’, ’10x’, ‘(-)’, ‘5’, ‘)”, ‘)”, ‘ENTER’)
Q: What is the [H+] if pH = 7.4?
[H+] = 10–pH mol/L
(’10x’, ‘(-)’, ‘7.4’, “)” ‘ENTER‘)
Ans: 4.2
3.98x10–8 M
Calculating pH from Strong Acid
Solutions
• Strong acids are 100% ionized
• For monoprotic acids, concentration of the
acid IS the concentration of the H+ ion
• Use Acid concentration as substitute for
H+ ion concentration.
– Use Base concentration as substitute for OHconcentration
Calculating pH from Strong Acid
Solutions
• Example: What is the pH of a 0.1M solution of
HCl?
– 0.1 M HCl = 1 x 10-1 M
– Calculate pH = 1
• Example: What is pH of solution that is 7.5 x 10-4
M Ca(OH)2?
– (7..5 x 10-4) x 2 = 1.5 x 10-3M
• There are 2 OH- ions per molecule
– Calculate pOH = -log[OH-]
• = 2.8
• pH = 14-2.8 = 11.2
Calculating Molarity from pH
• Example: what is the molarity of an acid
solution with a pH of 2.37?
– [H+] = 10-pH
– [H+] = 10-2.37 = 4.27 x 10-3 M
Neutralization
Chapter 15.4
Acid-Base Reactions
• Neutralization reaction is a reaction between an
acid and a base
– Makes Water + Salt
– Solution becomes Neutral (not acid or base)
HCl + NaOH
H2O + Na+ + Cl-
Acid-Base Reactions
• Mg(OH)2 + 2 HCl → MgCl2 + 2H2O
• Note:
– Cation from base (Mg) is combined with anion
from acid (Cl)
– The salt is MgCl2
– The H+ and OH- always combine to form
water
Acid-Base Titration
• Acid/Base Titration is the stoichiometry of
acid/base reactions.
– Titration is a method for determining the
concentration of a solution by using another
solution of known concentration
– Uses an INDICATOR to show when the
acid/base reaction is complete (neutral)
• Indicator is a chemical that changes color as
determined by acid or base conditions
• There are many indicators with different pH points.
Acid/Base Titration Curve
pH Indicators
pH Range of
Color Change
Name
Acid Color
Base Color
Methyl violet
Yellow
0.0 - 1.6
Blue
Thymol blue
Red
1.2 - 2.8
Yellow
Methyl orange
Red
3.2 - 4.4
Yellow
Bromocresol green
Yellow
3.8 - 5.4
Blue
Methyl red
Red
4.8 - 6.0
Yellow
Litmus
Red
5.0 - 8.0
Blue
Bromothymol blue
Yellow
6.0 - 7.6
Blue
Thymol blue
Yellow
8.0 - 9.6
Blue
Phenolphthalein
Colorless
8.2 - 10.0
Pink
Thymolphthalein
Colorless
9.4 - 10.6
Blue
Alizarin yellow R
Yellow
10.1 - 12.0
Red
Calculating Molarity from Titration
1. Write the balanced equation
2. Calculate the number of moles used in
the ‘known’ solution
3. Use the mole ratio from the balanced
equation to calculate moles of reactant in
the ‘unknown’ solution
4. Calculate the molarity of the ‘unknown’
solution based on moles used and liters
used.
Salt Hydrolysis
• When you put salts in water, the resulting
solution can be either acid, base, or
neutral
– Salts will dissolve to form ions
– The anions will accept hydrogens from water
– The cations will accept hydroxides from water
– Which way it goes depends upon the strength
of the conjugate acids/bases
– If conjugate acid is strong, it will be acid
– If conjugate base is strong, it will be base
– If both are strong, it will be neutral
What are Buffers?
• Buffers are solutions that resist changes in
pH when limited amounts of acid or base
are added.
– Buffer is a weak acid and its conjugate base
or a weak base and it’s conjugate acid
– Buffer can accept or donate Hydrogen ions
and shift its equilibrium point left or right
– Buffers have limits, but the work for a while
– Heavily used in human body, especially blood
Acid Names
1. Binary or hydrohalic acids – HF, HCl, HBr, HI,
etc. “hydro____ic acid” are usually strong
acids
• If name ends in ‘-ide’
– Acid name will be “hydro ____ic acid”
• HF and H2S are weak hydrohalic acid. Although
the H-F bond is very polar, the bond is so strong
(due to the small F atom) that the acid does not
completely ionize.
Acid Naming
2. Oxyacids – contain a polyatomic ion
a. Most common form (MCF) “ic” ending –
strong acids (contain 2 oxygen per
hydrogen)
b. If chemical name ends in “-ate”
a.
–
–
–
–
Acid name will be “___IC Acid”
HNO3 – nitric
from nitrate
H3PO4 - phosphoric from phosphate
H2SO4 - sulfuric
from sulfate
HClO3 - chloric
from chlorate
b. Acids with 1 less oxygen than the MCF “ous”
ending- weaker acids
a. Chemical name ends in “-ite”
b. Acid name is “___OUS Acid”
HNO2 – nitrous
from nitrite
H3PO3 - phosphorous from phosphite
H2SO3 - sulfurous
from sulfite
HClO2 - chlorous from chlorite
c. Acids with 2 less oxygen than the MCF
“hypo___ous” – very weak acids
HNO - hyponitrous
H3PO2 - hypophosphorus
HClO - hypochorous
d. Acids with 1 more oxygen than the MCF
“per______ic” – very strong acids
HClO4 – perchloric acid
HNO4 - pernitric acid
e. Organic acids – have carboxyl group COOH - usually weak acids
•
Acid names are based on the base organic
name or common name
HC2H3O2 - acetic acid
C7H5COOH - benzoic acid