Transcript Chapter 19 Acid

Chapter 15
Acid-Base Theories
Properties of Acids and Bases
• Acids
– Give foods a tart or sour taste
• What acidic foods might you eat?
– Aqueous solutions of acids are electrolytes
• Conduct Electricity
– Some are strong electrolytes (strong acids)
– Some are weak electrolytes (weak acids)
• Cause indicator dyes to change colors
• Many metals react with acids producing hydrogen gas
• React with compounds containing hydroxide ions to form
water and a salt
Properties of Acids and Bases
• Bases
– Have bitter taste, and slippery feel
– Aqueous solutions of bases are also
electrolytes
• Conduct Electricity
– Some are strong electrolytes (strong bases)
– Some are weak electrolytes (weak bases)
• Cause indicator dyes to change colors
• Water and salt are formed when a base that
contains hydroxide ions react with an acid
Arrhenious Acids and Bases
• Acids are hydrogen-containing compounds
that ionize to yield hydrogen ions (H+) in
aqueous solution
• Bases are compounds that ionize to yield
hydroxide ions (OH-) in aqueous solution
Arrhenious Acids
• Can be monoprotic, diprotic, or triprotic
– Monoportic: HNO3 → H+ + NO3• Ionization yields one hydrogen ion
– Diprotic: H2SO4 → 2H+ + SO42• Complete ionization yields 2 hydrogen ions
– Triprotic: H3PO4 → 3H+ + PO43• Complete ionization yields 3 hydrogen ions
• Not all the hydrogens in an acid may be released as hydrogen ions
• Not all hydrogen containing compounds are acids
– Only hydrogens joined to very electronegative elements, and thus have
very polar bonds, are ionizable in water
H O
H C C
O-
H
Ethanoic Acid
H+
Nonionizable
Hydrogen
Ionizable
Hydrogen
Arrhenious Bases
• NaOH → Na+(aq) + OH-(aq)
• KOH → K+(aq) + OH-(aq)
– Bases formed with group one metals are very soluble
and caustic
– Can be made by reacting group one metals with
water
• Na + H2O → Na+(aq) + OH-(aq) H2 (g)
• Bases of group 2 metals are very weak resulting
low solubility
– Examples are Ca(OH)2 and Mg(OH)2
Bronsted-Lowry Acids and Bases
• Arrhenious definition of acids and bases is not very
comprehensive and does not explain why certain
substances have basic or acidic properties
– Ammonia (NH3) is a base, but there is no hydroxide (OH-) in the
compound to ionize
• The Bronsted-Lowry theory defines an acid as a
hydrogen-ion donor, and a base as a hydrogen-ion
acceptor
– Why ammonia is a base
NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)
Hydrogen ion
aceptor, BronstedLowry Base
Hydrogen ion
donar, BronstedLowry Acid
Makes the
solution basic
Conjugate Acids and Bases
conjugate acid-base pair
NH4+(aq) + OH-(aq)
NH3(aq) + H2O(l)
Base
Conjugate
Acid
Acid
Conjugate
Base
conjugate acid-base pair
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A conjugant acid is the particle formed when a base gains a hydrogen ion
A conjugant base is the particle that remains when an acid has donated a
hydrogen ion
A conjugate acid-base pair consists of two substances related by the loss or
gain of a single hydrogen ion
conjugate acid-base pair
HCl(aq) + H2O(l)
Acid
Base
H3O+(aq) + Cl-(aq)
Conjugate
Acid
conjugate acid-base pair
Conjugate
Base
Conjugate Acids and Bases
conjugate acid-base pair
HCl(aq) + H2O(l)
Acid
Base
H3O+(aq) + Cl-(aq)
Conjugate
Acid
Conjugate
Base
conjugate acid-base pair
• A water molecule that gains a hydrogen ion becomes a positively
charged hydronium ion H3O+
– In above equation, what is the hydrogen ion donor (acid) and which is
the hydrogen ion acceptor (base)
• Notice, water can both accept and donate a hydrogen ion and thus
act as an acid and a base
– A substance that can act as both an acid and a base is said to be
amphoteric
• Amino Acids as an example
Lewis Acids and Bases
• Acids accept a pair of electrons during a
reaction while a base donates a pair of
electrons
– Lewis acid – a substance that can accept a
pair of electrons to form a covalent bond
– Lewis base – a substance that can donate a
pair of electrons to form a covalent bond
NH3 + BF3 → NH3BF3
Identify the Lewis Acid and the Lewis Base in the above equation
Acid-Base Definitions Review
Type
Acid
Base
Arrhenius
H+ producer
OH- producer
Bronsted-Lowry
H+ donor
H+ acceptor
Lewis
electron-pair acceptor
electron-pair donor
Hydrogen Ions and Acidity
• Occasionally collusions between water molecules
cause them to react forming hydroxide ions and
hydronium ions
– The reaction in which water molecules produce ions is called
the self ionization of water
H2O(l)
H2O (l) + H2O (l)
H+(aq) + OH-(aq)
H3O+(aq) + OH-(aq)
• In aqueous solution, hydrogen ions H+ are always joined
to a water molecule as hydronium ions
• In pure (neutral) water, the self-ionization of water results
in 1 x 10-7 M of H+ ions and 1 x 10-7 M of OH- ions
– Any aqueous solution in which H+ and OH- ions are equal is
described as a neutral solution
Ion Product Constant for Water
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For aqueous solutions, the product of the hydrogen-ion concentration and
the hydroxide-ion concentration equals 1.0x10-14
Kw = [H+] x [OH-] = 1.0x10-14 M
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The product of the concentrations of the hydrogen ions and hydroxide ions
in water is called the ion-product constant
Aqueous acids and bases sift the ratio of hydrogen ions to hydroxide ions in
solution causing it to become either acidic or basic
– In a basic solution aka alkaline solution, the hydroxide ion (OH-) is greater than
1x10-7 M and the hydrogen ion (H+) is less 1x10-7 M
– In a acidic solution, the hydrogen ion (H+) is greater than 1x10-7 M and the
hydroxide ion (OH-) is less 1x10-7 M
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Regardless of the acidity or alkalinity of the solution, the product of the
Molarity (M) concentration of H+ and OH- always equals 1x10-14 at 25ºC
– If the hydrogen ion H+ concentration in a soft drink is 1 x 10-5 M, what is the
concentration of the hydroxide ion OH-?
– Is the solution basic, neutral, or acidic?
The pH Concept
• Expressing hydrogen-ion concentrations in molarity is cumbersome
– Soren Sorensen suggested that the hydrogen ion concentration be
expressed as the negative log of the hydrogen-ion concentration giving
us much smaller numbers to work with
pH = -log[H+]
• The pH of a solution is the negative logarithm of the hydrogen ion
concentration
• A neutral solution H+ = 1x10-7 has a pH = -log[1x10-7]= 7
• A solution in which the [H+] is greater than 1x10-7 M and has a pH less than
7.0 is acidic
• The pH of pure water or a neutral solution has a pH of 7
• A solution with a pH greater than 7 is basic and has a [H+] concentration of
less than 1x10-7 M
– You can also calculate pOH which is the negative logarithm of
hydroxide ion concentration
pOH = -log[OH-]
Work some pH problems (pH of 1 x10-5 M H+?)
Relationship between pH and pOH
pH + pOH = 14
pH = 14 – pOH
pOH = 14 - pH
pH and significant figures
• Hydrogen ion concentrations should
always be reported to two significant
figures
• pH and pOH calculations should always
be reported to two decimal places
– Rules are due to the sensitivity of pH meters
Acid-Base Indicators
Dyes
• An indicator (HIn) is an acid or base that undergoes
dissociation in a known pH range
– An indicator is a valuable tool for measuring pH because its
acid form and base form have different colors in solution
HIn (aq)
Acid Form
H+
OH-
H+ (aq) + In- (aq)
Base Form
– The acid form of the indicator dominates the disassociation
equilibrium at low pH
– The basic form of the indicator dominates the disassociation
equilibrium at high pH
– Color change of an indicator occurs in a narrow pH range ≈ 2
pH units
– Thus it takes many indicators to span the entire pH spectrum
• Indicator dyes have limitations
Acid-Base Indicators
pH Meter
• Makes rapid, accurate pH
measurements
• Can record pH continuously over time
when performing a reactions
• Measures pH to two decimal places
• Color and cloudiness of solution does
not interfere with reading
• Are many different types specialized
for different jobs were pH
measurements are required
http://www.vittbi.com/photogallery/biotech/PH-Meter.jpg
http://personals.galaxyinternet.net/tunga/Meter.jpg
Strengths of Acids and Bases
Strong and Weak Acids and Bases
• Acids are classified as strong or weak
depending on the degree to which they ionize in
water
– A strong acid completely ionizes in water
– Weak acids ionize only slightly in aqueous solution
– What are some example of strong acids and weak
acids
HCl(q) + H2O(l)
CH3COOH(aq) + H2O(l)
H3O+(aq) + Cl-(aq) 100% ionized
H3O+(aq) + CH3COO-(aq) <1% ionized
pH of 0.10 M Solutions of Common Acids and Bases
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Compound
HCl (hydrochloric acid)
H2SO4 (sulfuric acid)
NaHSO4 (sodium hydrogen sulfate)
H2SO3 (sulfurous acid)
H3PO4 (phosphoric acid)
HF (hydrofluoric acid)
CH3CO2H (acetic acid)
H2CO3 (carbonic acid)
H2S (hydrogen sulfide)
NaH2PO4 (sodium dihydrogen phosphate)
NH4Cl (ammonium chloride)
HCN (hydrocyanic acid)
Na2SO4 (sodium sulfate)
NaCl (sodium chloride)
NaCH3CO2 (sodium acetate)
NaHCO3 (sodium bicarbonate)
Na2HPO4 (sodium hydrogen phosphate)
Na2SO3 (sodium sulfite)
NaCN (sodium cyanide)
NH3 (aqueous ammonia)
Na2CO3 (sodium carbonate)
Na3PO4 (sodium phosphate)
NaOH (sodium hydroxide, lye)
pH
1.1
1.2
1.4
1.5
1.5
2.1
2.9
3.8 (saturated solution)
4.1
4.4
4.6
5.1
6.1
6.4
8.4
8.4
9.3
9.8
11.0
11.1
11.6
12.0
13.0
http://www.cartage.org.lb/en/themes/sciences/chemistry/Inorganicchemistry/AcidsBases/Common/Common.htm
Acid Disassociation Constant
• The equilibrium constant for weak acids (HA) can be written as:
Acid
Conjugate base
H3O+(aq) + A-(aq)
HA(aq) + H2O(l)
Keq =
[H3O+] X [A-]
[HA] X [H2O]
• For dilute solutions, the concentration of water is a constant, and
can be combined with Keq to give the acid dissociation constant
(Ka)
Keq X H2O = Ka=
[H3O+] X [A-]
[HA]
• Ka reflects the fraction of an acid in the ionized form and thus is
sometimes referred to as the ionization constant
• Weak acids have small Ka values, while stronger acids have
larger Ka values; why?
Base Disassociation Constant
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The equilibrium constant for weak Bases (B) can be written as:
base
Conjugate acid
BH+(aq) + HO-(aq)
B(aq) + H2O(l)
Keq =
[BH+] X [HO-]
[B] X [H2O]
• For dilute solutions, the concentration of water is a constant, and
can be combined with Keq to give the base dissociation constant
(Kb)
Keq X H2O = Kb=
[BH+] X [OH-]
[B]
• Kb is the ratio of the concentration of the conjugate acid times the
concentration of the hydroxide ion to the concentration of the
base
• The magnitue of Kb indicates the ability of a weak base to
compete with the very strong base OH- for hydrogen ions
• The smaller the Kb the weaker the base
Concentration and Strength
• Remember, the word strong and weak acids and bases
refers to the number particles of the acid or base that
completely dissociate into their respective ions in
solution
• Concentration and dilute refer to how many moles of an
acid or base is diluted in a constant volume of solution
• Even though an acid may be “weak”, if it is highly
concentrated, it will result in much lower pH of the
solution it is dissolved in than a dilute solution of the
same weak acid
Calculating Dissociation Constants
• Disassociation constants are calculated from
experimental data
• To find the Ka of weak acid or the Kb of a weak
base, substitute the measured concentrations of
all the substances present at equilibrium into the
expression for Ka or Kb
• A 0.1000M solution of ethanoic acid is only
partially ionized and has a pH of 2.87. What is
the acid dissociation constant (Ka) or ethanoic
acid
Neutralization Reactions
Acid Base Reactions
• Reactions in which an acid and a base react in
an aqueous solution to produce a salt and water
are called neutralization reactions
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l)
Titrations
• Titration - the process used to determine the concentration of
solution (often an acid or base) in which a solution of known
concentration (the standard) is added to a measured amount of the
solution of unknown concentration until an indicator signals the end
point
– In titrations, it is important to know the mole ratios that the acid and
base in question react
– When an acid and base mix, the equivalence point is when the number
of moles of hydrogen ions equals the number of moles hydroxide ions
giving a pH of 7
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Mole ratio 1:1
H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l)
Mole ratio 1:2
2HCl(aq) + CaOH(aq) → CaCl2(aq) + 2H2O(l)
Mole ratio 2:1
Steps in Titrating a Neutralization Reaction
1. A measured volume of an acid solution of unknown
concentration is added to a flask
2. Several drops of the indicator are added to the
solution while the flask is gently swirled
3. Measured volumes of a base of known
concentration are mixed into the acid until the
indicator just barely changes color
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The solution of known concentration is called
the standard solution
The point at which the indicator changes color
is the end point
The point of neutralization is the end point aka
equivalence point of the titration
Titration of Strong Acid with a Strong Base
http://www.bwsd.k12.wi.us/astitt/Chemistry/AcidsandBases/TitrationCurve-StrongAcid.gif
Problem
• A 25 ml solution of H2SO4 is completely
neutralized by 18 ml of 1.0M NaOH. What
is the concentration of the H2SO4 solution?
Salts in Solution
Salt hydrolysis
• Salt consist of anion from an acid and a
cation from a base
– Solutions of many salts are neutral while other
salt solutions are not