Chapter 14 Section 1 Properties of Acids and Bases Section 2 Acid Base Theories Section 3 Acid Base Reactions 141 Properties of Acids and Bases List five general properties of aqueous acids and bases ID: 935513
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Slide1
Chapter 14
Acids and Bases
Slide2Chapter 14
Section 1 – Properties of Acids and Bases
Section 2 – Acid Base Theories
Section 3 – Acid Base Reactions
Slide314.1 Properties of Acids and Bases
List
five general properties of aqueous acids and bases.
Name
common binary acids and
oxyacids
, given their chemical formulas.
List
five acids commonly used in industry and the laboratory, and give two properties of each.
Define
acid and base according to Arrhenius’s theory of ionization.
Explain
the differences between strong and weak acids and bases.
Slide4Properties of:
Acids Bases
Sour taste
Conducts electricity
Turns litmus paper red
Reacts with bases to produce salts and water
Reacts with some metals and releases hydrogen gas
Bitter taste
Feels slipperyConducts electric currentTurns litmus paper blueReacts with acids to produce salts and water
Slide5Binary Acids
Contains only two different elements
Hydrogen & an electronegative, nonmetal
Nomenclature:
hydro - _________ -
ic
acid
Slide6Diatomic Nomenclature
Slide7Oxyacid
Contains hydrogen, oxygen, and a third element
(hydrogen with a polyatomic ion)
Nomenclature:
Slide8Acid Names
Slide9Oxyacids
Slide10Common Industrial Acids
Sulfuric Acid
Sulfuric acid is the most commonly produced industrial chemical in the world.
Nitric Acid
Phosphoric Acid
Hydrochloric Acid
Conc.
HCl
is commonly referred to as muriatic acid.Acetic AcidPure acetic acid is a clear, colorless, and pungent-smelling liquid known as glacial acetic acid.
Slide11Arrhenius Acids and Bases
Arrhenius Acids
:
Increases concentration of H
+
ions in solution
Arrhenius Bases
:Increases concentration of OH- ions in solution
Slide12Arrhenius Acid Base Video
Slide13Acid/Base Strength
Strong acid
:
Ionizes completely in solution and is an electrolyte
Higher the K
A
, the greater the strength as an acid
K reveals a greater extent of ionizationExample: HCl, HClO4, HNO3Weak acid
:Releases few hydrogen ions in solutionHydronium ions, anions and dissolved acid molecules presentExamples: HCN, Organic acids – HC2H3O2
Slide14Dissociation Constants
Strong vs. Weak Base
Strong bases ionizes completely in solution and is a strong electrolyte
K
B
= dissociation constant of a base
Higher the K
B , the greater the strength of a base
Slide15Aqueous Acids
Slide16Base Strength
Strong bases
:
Ionic compounds containing metal cation and hydroxide ion (OH-)
Dissociates in water
Weak bases
:
Molecular compounds do not follow Arrhenius definition:
Ammonia (NH
3
)
Produces hydroxide ions when it reacts with water molecules
Slide17Base Strength
Slide18Acidic solution has greater [H
3
O
+
]
Basic solution has greater [OH
–
]
Slide1914.2 Acid Base Theories
Define
and
recognize
Brønsted
-Lowry
acids and bases.
Define a Lewis acid and a Lewis base.Name compounds that are acids under the Lewis definition but are not acids under the Brønsted-Lowry definition.
Slide20Bronsted
-Lowry Acid
Bronsted
-Lowry Acid
:
Proton (H
+
) donor
Hydrogen chloride acts as a
Bronsted
-Lowry acid when it reacts with ammonia.
Water can also act as a
Bronsted
-Lowry acid
Slide21Bronsted
-Lowry Base
Bronsted
-Lowry Base
:
Proton acceptor
Ammonia accepts a proton from hydrochloric acid.
Slide22acid
base
Bronsted
-Lowry Acid Base Reactions
Protons are transferred from one reactant (the acid) to another (the base)
Slide23acid conjugate
base
Conjugate Acid – Base
Conjugate Base
:
The species that remains after a
Bronsted
-Lowry acid has given up a proton
Conjugate Acid
:
The species that remains after a
Bronsted
-Lowry base has accepted a proton
Slide24Conjugate Acid Base Pairs
Match up the acid-base pairs
(proton donor-acceptor pairs)
acid
1
base
2
conjugate base
1
conjugate
acid
2
Slide25strong
acid
base
acid
weak base
Strength of Acid Base Pairs
The stronger the acid, the weaker the conjugate base
The stronger the base, the weaker the conjugate acid
Slide26stronger acid stronger base
weaker acid weaker base
weaker acid weaker base
stronger acid stronger base
Proton transfer favors the production of the weaker acid and base.
Slide27Acid Base Strength
Slide28acid
1
base
2
acid
2
base
1
base
1
acid
2
acid
1
base
2
Amphoteric
Any species that can react as either an acid or a base
Example: water
Slide29Amphoteric
Water Video
Slide30Other Amphoteric Compounds
Covalently bonded –OH group in an acid is referred to as a hydroxyl group
Molecular compounds with hydroxyl groups can be acidic or
amphoteric
The behavior of the compound is affected by the number of oxygen atoms bonded to the atom connected to the –OH group
*The more oxygen’s in a polyatomic formula, the greater the strength of polyatomic as an acid
Slide31Oxyacids of Chlorine
Slide32Brønsted
-Lowry Acid Base Video
Slide33Monoprotic Acids
Can donate only one proton (hydrogen ion) per molecule
One ionization step
Slide34Monoprotic and Diprotic Acids
Slide35Polyprotic Acids
Donates more than one proton per molecules
Multiple ionization steps
Diprotic
– donates 2 protons Ex:
Triprotic
– donates 3 protons Ex:
Sulfuric acid solutions contain H
3
O
+
, HSO
4
-
, SO
4
-
ions
1.
2.
Slide36Lewis Acid
Lewis acid
:
Atom, ion, or molecule that ACCEPTS an ELECTRON PAIR to form a covalent bond
A proton (hydrogen ion) is a Lewis acid
Lewis base
:
Atom, ion, or molecule that DONATES an ELECTRON PAIR to form a covalent bond
Slide37Lewis Acid
A
lewis
acid might not include hydrogen
Silver as a
lewis
acid:
Slide38Lewis Acid Base Video
Slide39Acid and Base Definitions
Slide40Acid Base Definitions Video
Slide4114.3 Acid
Base Reactions
Describe
a conjugate acid, a conjugate base, and an
amphoteric
compound.
Explain
the process of neutralization.
Define acid rain, give examples of compounds that can cause acid rain, and describe effects of acid rain.
Slide42Neutralization Reactions
What does it mean to neutralize something?
Neutralization reactions:
Hydronium and hydroxide ions react to form water
The left over cation and anion in solution produce a salt (ionic compound)
Slide43Neutralization Reactions
Slide44Neutralization Reaction Video
Slide45Acid Rain
NO, NO
2
, CO
2
, SO
2
, and SO
3
gases from industrial processes can dissolve in atmospheric water to produce acidic solutions.
Very acidic rain is known as
acid rain.
Acid rain can erode statues and affect ecosystems.
Slide46Chapter 15
Acid Base Titration and pH
Slide47Chapter 15
Section 1 – Aqueous Solutions and the Concept of pH
Section 2 – Determining pH and Titrations
Slide4815.1 Aqueous Solutions and pH
Describe
the self-ionization of water.
Define
pH, and give the pH of a neutral solution at 25°C.
Explain
and use the pH scale.
Given
[H3O+] or [OH−], find pH.
Given
pH,
find
[H
3
O
+
] or [OH
−
].
Slide49Self Ionization of Water
Two water molecules produce a hydronium ion and hydroxide ion by proton transfer
In water at 25°C,
[H
3
O
+
] = 1.0 ×10
−7
M and [OH
−
] = 1.0 × 10
−7
M
The ionization constant of water,
K
w
K
w
= [H
3
O
+
][OH
−
]
Slide50At 25
O
C
K
w
= [H
3
O
+][OH−] = (1.0 × 10−7)(1.0 × 10
−7
) = 1.0 × 10
−14
K
w
= 1.0 x 10
-14
K
w
increases as temperature increases
Slide51Ion Concentration
[H
3
O
+
] = [OH
−
]
neutral[H3O+] > [OH−]
acidic
[H
3
O
+
] >
1.0 × 10
−7
M
[OH
−
] > [H
3
O
+
]
basic
[OH
−
] >
1.0 × 10
−7
M
Slide52Calculating Concentration
Strong acids and bases are considered
completely
ionized or dissociated in aqueous solutions.
1 mol 1 mol 1 mol
1.0 × 10
−2
M
NaOH
therefore,
[OH
−
] = 1.0 × 10
−2
M
[H
3
O
+
] is calculated using
K
w
Slide53Example Problem 1
Given:
[
HCl
] = 2.0 × 10
−4
M
[H
3
O
+
] = ______________
Unknown: [OH
-
] = ?
K
w
= [H
3
O
+
][OH
−
] = 1.0 × 10
−14
Slide54pH
Definition
of the
pH
of a solution: negative of the common logarithm of the hydronium ion concentration, [H
3
O+].pH = −log [H
3O+] Example: a neutral solution has a [H3O+
] = 1×10
−7
pH = −log [H
3
O
+
] = −log(1 × 10
−7
) = −(−7.0) = 7.0
Slide55pH Values as Specified [H
3
O
+
]
Slide56The pH Scale
Slide57pOH
The
pOH
of a solution is defined as the negative of the common logarithm of the hydroxide ion concentration, [OH
−
].
pOH
= −log [OH
–]
pH +
pOH
= 14.0
Example
:
a neutral solution has a [OH
–
] = 1×10
−7
the pH of this solution is?
Slide58Calculating [H
3
O
+
] from pH
Finding the [H
3
O
+] from pH requires taking the antilog of the negative pH [H3O+] = antilog (-pH)
You can find the [OH
−
] by also taking the antilog of the negative
pOH
.
[OH
-
] = antilog (-
pOH
)
Slide59The Circle of pH
pH
pOH
[ H
3
O
+
]
[ OH
-
]
-log [H
3
O
+
]
antilog (-pH
)
antilog (-pOH)
-log [OH
-
]
[ H
3
O
+
]
[ OH
-
]
=
1.0x10
-14
pH
+ pOH
= 14
Slide60pOH
Video
Slide61pH Values of Some Common Materials
Slide62Approximate pH Range of Common Materials
Slide63Comparing pH and
pOH
Video
Slide64pH of Weak Acids and Bases
The pH of solutions of weak acids and weak bases must be measured experimentally.
The [H
3
O
+
] and [OH
−
] can then be calculated from the measured pH values.
Slide65Significant Figures
There must be as many significant figures to the right of the decimal as there are in the number whose logarithm was found.
Example
: [H
3
O
+
] = 1 × 10
−7 one significant figure pH = 7.0
Slide6615.2 Determining pH and Titrations
Describe
how an acid-base indicator functions.
Explain
how to carry out an acid-base titration.
Calculate
the molarity of a solution from titration data.
Slide67Indicators
Acid-base indicators:
compounds whose colors are sensitive to
pH.
The pH range over which an indicator changes color is called its
transition interval
.
Slide68pH Meters
pH meter
determines the pH of a solution by measuring the voltage between the two electrodes that are placed in the solution.
The voltage changes as the hydronium ion concentration in the solution changes.
Measures pH more precisely than indicators
Slide69Color Ranges of Indicators
Slide70Color Ranges of Indicators
Slide71Color Ranges of Indicators
Slide72Antacids Video with Methyl Orange
Slide73Titration
Neutralization occurs when hydronium ions and hydroxide ions are supplied in equal numbers by reactants.
H
3
O
+
(
aq
) + OH−
(
aq
) 2H
2
O(
l
)
Titration:
the controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration.
Slide74Titration Points
equivalence point:
point at which the two solutions used in a titration are present in chemically equivalent amounts
end point:
point in a titration at which an indicator changes color
Slide75Which indicator do I choose?
pH less than 7
Indicators that change color at pH lower than 7 are used to determine the equivalence point of strong-acid/weak-base titrations.
strong-acid/weak-base titration = acidic.
pH at 7
Indicators that undergo transition at about pH 7 are used to determine the equivalence point of strong-acid/strong base titrations.
strong acids/strong bases = salt solution with a pH of 7.
Slide76Which indicator do I choose?
pH greater than 7
Indicators that change color at pH higher than 7 are used to determine the equivalence point of weak-acid/strong-base titrations.
weak-acid/strong-base = basic
Slide77Titration Curve
Strong Acid and a Strong Base
Equivalence Point:
pH at 7
Slide78Titration Curve
Weak Acid and a Strong Base
Equivalence Point:
pH higher than 7
Slide79Titration Curve
Strong Acid and a Weak Base
Equivalence Point:
pH less than 7
Slide80Titration Problems:
* Can be used to determine concentration of unknown solution or volume of added standard
Start with the balanced equation for the neutralization reaction
Make amount of acid and base chemically equivalent to each other (1 to 1 mol ratio).
Determine the molarity of the unknown solution.
Equation: M
1
V1
= M
2
V
2
1: starting solution
2: added standard
Slide81Molarity and Titration
standard solution
: solution that contains the precisely known concentration of a solute
primary standard:
highly purified solid compound used to check the concentration of the known solution
The standard solution can be used to determine the molarity of another solution by titration.
Slide82Performing a Titration – Set up
Slide83Performing a Titration – Set up Acid
Slide84Performing a Titration – Starting Amount
Slide85Performing a Titration – Set up Base
Slide86Performing a Titration - Titrating
Slide87Performing a Titration – End Point
Slide881 mol 1 mol 1 mol 1 mol
Molarity and Titration
Determine the molarity of an acidic solution,
10
mL
HCl
, by titration
Titrate acid with a standard base solution
20.00
mL
of 5.0 × 10
−3
M
NaOH
was titrated
Write the balanced neutralization reaction equation.
HCl
(
aq
) +
NaOH
(
aq
)
NaCl
(
aq
) + H
2
O(
l
)
Slide89Molarity and Titration
Calculate the number of moles of
NaOH
used in the titration.
20.0
mL
of 5.0 × 10
−3
M
NaOH
is needed to reach the end point
mol of
HCl
= mol
NaOH
= 1.0 × 10
−4
mol
Calculate the molarity of the
HCl
solution
Slide90Example Problem
In a titration, 27.4
mL
of 0.0154 M
Ba
(OH)
2
is added to a 20.0 mL sample of HCl
solution of unknown concentration until the equivalence point is reached. What is the molarity of the acid solution?
Slide91Ba
(OH)
2
+ 2HCl BaCl
2
+ 2H
2O 1 mol 2 mol 1 mol 2 mol
Example Problem SolutionGiven: 27.4 mL of 0.0154 M Ba
(OH)
2
Unknown
: ? M
HCl
of 20.0
mL
Solution
:
Write balanced equation:
Slide921. Calculate Moles of Given
Slide932. Write a mole ratio:
moles of base used to moles of acid produced
Slide943. Calculate Unknown Molarity