Transcript Acids and Bases - Cathedral High School

Acids and
Bases
Chapter 15
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Acids and Bases
• The concepts acids and bases were loosely defined
as substances that change some properties of
water.
• One of the criteria that was often used was taste.
• Substances were classified
– salty-tasting, sour-tasting, sweet-tasting, bittertasting
• Sour-tasting substances would give rise to the word
'acid', which is derived from the Greek word oxein,
which mutated into the Latin verb acere, which
means 'to make sour'
• Vinegar is a solution of acetic acid. Citrus fruits
contain citric acid.
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Acids
•
•
React with certain metals to
produce hydrogen gas.
React with carbonates and
bicarbonates to produce carbon
dioxide gas
Bases
• Have a bitter taste
• Feel slippery.
• Many soaps contain bases.
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Properties of Acids
þ Produce H+ (or H3O+) ions in water (the hydronium ion is
a hydrogen ion attached to a water molecule)
þ Taste sour
þ Corrode metals
þ Good Electrolytes
þ React with bases to form a salt and water
þ pH is less than 7
þ Turns blue litmus paper to red
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Properties of Bases
 Generally produce OH- ions in water
 Taste bitter, chalky
 Are electrolytes
 Feel soapy, slippery
 React with acids to form salts and water
 pH greater than 7
 Turns red litmus paper to blue
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Arrhenius Definition
Arrhenius
Acid - Substances in water that increase the
concentration of hydrogen ions (H+).
Base - Substances in water that increase
concentration of hydroxide ions (OH-).
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Bronsted-Lowry Definition
Acid - neutral molecule, anion, or cation
that donates a proton.
Base - neutral molecule, anion, or cation
that accepts a proton.
HA + :B
HCl + H2O
Acid
Base


HB+
+
H3O+ +
Conj Acid
:AClConj Base
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Conjugate Acid Base Pairs
Conjugate Base - The species remaining after an
acid has transferred its proton.
Conjugate Acid - The species produced after base
has accepted a proton.
HA & :A:A-
- conjugate acid/base pair
- conjugate base of acid HA
:B & HB+
- conjugate acid/base pair
HB+
- conjugate acid of base :B
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Examples of Bronsted-Lowry Acid
Base Systems
Note: Water can act as acid or base
Acid
Conjugate Acid
Base
Conjugate Base
HCl
+
H2O 
H3O+ +
Cl-
H2PO4-
+
H2O
H3O+ +
HPO42-
NH4+
+
H2O 
H3O+ +
NH3
Base
NH3
+
Acid Conjugate Acid Conjugate Base :
H2O 
NH4+ +
OH-
PO43-
+


H2O 
HPO42- +
OH9
G.N. Lewis Definition
Lewis
Acid - an electron pair acceptor
Base - an electron pair donor
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pH and acidity
HC2H3O2 + H2O  H3O+ + C2H3O2NH3 + HCl  NH4+ + ClHSO4- + HPO4-2  H2SO4 + PO4-3
HSO4- + HPO4-2  SO4-2 + H2PO4-1
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pH and acidity
The pH values of several
common substances are
shown at the right.
Many common foods are
weak acids
Some medicines and many
household cleaners are
bases.
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The pH Scale
pH
[H3O+ ]
[OH- ]
pOH
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pH and acidity
1. Acidity or Acid Strength depends on Hydronium Ion
Concentration [H3O+]
2. The pH system is a logarithmic representation of the
Hydrogen Ion concentration (or OH-) as a means of avoiding
using large numbers and powers.
pH
= - log [H3O+]
pOH = - log [OH-]
In pure water [H3O+] = 1 x 10-7 M
 pH = - log(1 x 10-7) = 7
4. pH range of solutions:
0 - 14
pH < 7 (Acidic) [H3O+] > 1 x 10-7 M
pH > 7 (Basic) [H3O+] < 1 x 10-7 M
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pH and acidity
Important Equations
pH = - log [H3O+]
[H3O+] = 10-pH
pOH = - log [OH-]
[OH-] = 10-pOH
pH + pOH = 14
Kw = [H3O+] [OH-] = 1.0 x10-14
In pure water: [H3O+] = [OH-] = 1.0 x10-7
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Calculating the pH
pH = - log [H3O+]
Example 1: If [H3O+] = 1 X 10-10
pH = - log 1 X 10-10
pH = - (- 10)
pH = 10
Example 2: If [H3O+] = 1.8 X 10-5
pH = - log 1.8 X 10-5
pH = - (- 4.74)
pH = 4.74
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pH and acidity
If a substance has a pH of 3.5, determine:
[H3O+], [OH-], pOH and whether it is acidic,
basic, or neutral.
If a substance has an [OH-] of 8.7 x 10-5,
determine: pH, pOH, [H3O+] and whether it is
acidic, basic, or neutral.
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Indicators
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Acid Strength
Strong Acid - Transfers all of its protons to water;
- Completely ionized or dissociated;
- Strong electrolyte;
- The conjugate base is very weak
Weak Acid
- Transfers only a fraction of its protons
to water;
- Only partly ionizes or dissociates;
- Weak electrolyte;
- The conjugate base is stronger
 As acid strength decreases, base strength
increases.
 The stronger the acid, the weaker its conjugate base
 The weaker the acid, the stronger its conjugate base
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Base Strength
Strong Base - all molecules accept a proton;
- completely ionizes or dissociates;
- strong electrolyte;
- conjugate acid is very weak
Weak Base - fraction of molecules accept proton;
- partly ionizes or dissociates
- weak electrolyte;
- the conjugate acid is stronger.
 As base strength decreases, acid strength increases.
 The stronger the base, the weaker its conjugate acid.
 The weaker the base the stronger its conjugate acid.
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Common Strong Acids
Strong Acids
Hydrochloric Acid, HCl
Hydrobromic Acid, HBr
Hydroioidic Acid, HI
Nitric Acid, HNO3
Sulfuric Acid, H2SO4
Chloric Acid, HClO3
Perchloric Acid, HClO4
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Dissociation of Strong Acids
• Strong acids completely dissociation (broke
down) to the ions that make them up.
• HCl  H+1 + Cl-1
• HNO3  H+1 + NO3-1
• If you know the concentration or molarity of a
strong acid, you also know the amount of H+
ions and can find the pH (-log [H+])
• 0.1 M HCl = 0.1 M H+, pH = -log 0.1 = 1.0
Common Strong Bases
Strong Bases
Any Group 1 Hydroxide and any Group 2
Hydroxide below Mg.
Sodium Hydroxide, NaOH
Potassium Hydroxide, KOH
*Barium Hydroxide, Ba(OH)2
*Calcium Hydroxide, Ca(OH)2
*While strong bases they are not very soluble
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Dissociation of Strong Bases
• Like strong acids, strong bases also completely
dissociate to the ions that make them up.
• KOH  K+1 + OH-1
• Ca(OH)2  Ca+2 + 2 OH-1
• Just like with strong acids, if you know the
molarity of a strong base, you can determine
the pOH and therefore the pH.
• 0.01 M NaOH = 0.01 M OH-,
• pOH = -log 0.01 = 2.0 and pH = 12.0
pH of Strong Acids and Strong Bases
• Determine the pH of 0.25 M KOH.
• Determine the pH of 0.0050 M H2SO4.
• Determine the pH of 0.0075 M Ca(OH)2.
Dissociation of Weak Acids
• Weak acids only dissociate to a very small
degree, most break down to ions less than 5%.
So an equilibrium is established where there is
a small amount of product (H+ and conjugate
base) but most of the acid does not break
down.
• These acids are weak electrolytes versus
strong acids are strong electrolytes.
Dissociation of Weak Acids
• To determine the pH of a weak acid, an “ICE”
chart must be used.
• I = Initial concentrations
• C = Change
• E = Equilibrium concentrations
• Unless otherwise noted, the initial
concentrations of the products are zero.
Dissociation Constants
• For a generic weak acid dissociation,
HA(aq) + H2O(l)

A−(aq) + H3O+(aq)
the equilibrium expression would be
Ka =
[H3O+] [A−]
[HA]
• This equilibrium constant is called the aciddissociation constant, Ka.
Dissociation of Weak Acids
• Equilibrium expressions are written as
products/reactants, and liquids and solids are
left out of the expressions.
• Ka are used to tell the relative strength of the
acid.
• You will be given the Ka value for each weak
acid.
Dissociation of Weak Acids
• Determine the pH of 0.025 M HC2H3O2.
Ka = 1.8 X 10-5.
Dissociation of Weak Acids
• Determine the pH of 1.0 M hypoiodous acid,
HIO. Ka of HIO = 2.3 x 10-11.
Determination of Ka for a Weak Acid
• If you are given the initial concentration of the
weak acid, and the equilibrium concentration
of the conjugate base or H+, you can
determine the Ka of the acid.
• Remember the initial concentrations of both
products are zero.
• If you know the equilibrium concentration of
the conjugate base or H+, you also know the
value of “x” in the “C” step of the “ICE” chart.
Determination of Ka for a Weak Acid
• Determine the Ka of a weak acid when the
initial concentration of HA as 0.015 M and the
equilibrium concentration of H+ = 0.0035 M.
• Equation: HA + H2O  A-1 + H3O+1
Titration
• Titration – technique for determining the
unknown concentration of an acid or base.
• Titration involves delivery (from a buret) of
a measured volume of a solution of known
concentration (the titrant, usually a base)
into a volume of solution of unknown
concentration (the analyte, usually an acid).
Titration
• Equivalence or stoichiometric point – point in
a titration where enough titrant has been
added to react exactly with the analyte.
• Equivalence pt: moles of acid = moles of base
Titration
Titration
• The equivalence point is often marked by an
indicator (commonly phenolphthalein), which
is added at the beginning of the titration. It
changes color at (or just after) the equivalence
point.
• The point where the indicator actually
changes color is called the endpoint of the
titration.
Titration
Titration Calculations
• 2.00 mL of an unknown concentration of HCl
was titrated with 0.25 M NaOH. If 30.23 mL of
NaOH was needed to reach the stoichiometric
point, what is the concentration of the HCl?
• Equation: HCl + NaOH  NaCl + H2O
Titration Calculations
• 2 drops of phenolphthalein is added to
10.0 mL of an 0.75 M of HCl. If 16.82 mL of
NaOH is needed to completely react with all of
the HCl, what is the concentration of the
NaOH?