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Acids, Bases, and Salts
Chapter 19
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Acid-Base Theories
Essential Question:
What are the properties of acids and bases, and what distinguishes the Arrhenius, Bronsted-Lowry and Lewis theories of
acids and bases?
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Properties of Acids
• Acids taste tart or sour.
• Aqueous solutions of acids are electrolytes.
• Acids change the color of acid-base indicators.
• Acids react with metals to produce hydrogen gas.
• Acids react with bases to form water and a salt.
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Properties of Bases
• Bases have a bitter taste.
• Bases have a slippery feel.
• Aqueous solutions of bases are electrolytes.
• Bases change the color of acid-base indicators.
• Bases react with acids to form water and a salt.
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Three Major Theories
• Arrhenius
• Bronsted-Lowry
• Lewis
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Arrhenius Acids and Bases
• Acids are hydrogen-containing compounds that ionize to produce H+ ions in aqueous solution.
• Bases are hydroxide-containing compounds that ionize to produce OH– ions in aqueous solution.
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Mono- Di- and Tri-protic Acids
• HNO3 is a monoprotic acid.
• H2SO4 is a diprotic acid.
• H3PO4 is a triprotic acid.
• Not all substances that contain hydrogen are acids.
• Not all hydrogens in acids are necessarily released as H+ ions.
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Hydrochloric Acid
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Ethanoic Acid (Acetic Acid)
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Arrhenius Bases
• The most commonly known is NaOH (lye).
• Another common base is KOH
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Bronsted-Lowry Acids and Bases
• Defines an acid as a hydrogen-ion donor.
• Defines a base as a hydrogen-ion acceptor.
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Why Ammonia is a Base
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Conjugate Acids and Bases
• When a substance donates a hydrogen ion, what remains has the ability to accept it back.
• When a substance accepts a hydrogen ion, what remains has the ability to donate the ion.
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Conjugate Acids and Bases
• A conjugate acid is formed when a base gains a hydrogen ion.
• A conjugate base is 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 H+.
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Consider Ammonia in Water
NH3 + H2O NH4+ + OH–
Base Acid Conjugate Conjugate
Acid Base
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Hydronium ion ( H3O+ )
• A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion.
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Amphoteric
• Sometimes water accepts a hydrogen.
HCl + H2O H3O+ + Cl–
• Other times, water donates a hydrogen.
NH3 + H2O NH4+ + OH–
• A substance that can act as either an acid or a base is said to be amphoteric.
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Lewis Acids and Bases
• A Lewis acid accepts a pair of electrons during a reaction.
• A Lewis base donates a pair of electrons during a reaction.
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Hydrogen Ions and Acidity
Essential Question:
How are [H+] and [OH–] related and how do they relate to acidity?
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Hydrogen Ions from Water
• The reaction in which water molecules produce ions is called the self-ionization of water.
H2O (l) H+ (aq) + OH– (aq)
• In water, H+ ions are always joined to water molecules to form H3O+ hydronium ions.
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Self-Ionization of Water
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Self-Ionization of Water
• This happens to a very small extent.
• In pure water, the concentration of H+ and OH– are equal.
• [H+] and [OH–] both equal 1.0 x 10-7 M.
• This is called a neutral solution.
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Ion Product Constant for Water
• In any aqueous solution, when [H+] increases, [OH–] decreases.
• When [H+] decreases, [OH–] increases.
• The product of the hydrogen-ion concentration and the hydroxide ion concentration always equals 1.0 x 10-14.
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Ion Product Constant for Water
• Kw = [H+] x [OH–] = 1.0 x 10-14
• Remember, as [H+] goes up, [OH–] goes down.
• But the product of [H+] x [OH–] will always be 1.0 x 10-14.
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Acidic vs Basic
• An acidic solution is one in which the [H+] is greater than the [OH–].
• A basic solution is one in which the [H+] is less than the [OH–].
• Basic solutions are also called alkaline solutions.
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The pH Concept
• Expressing hydrogen concentration in molarity can be cumbersome.
• A more popular method is the pH scale.
• The pH scale ranges from 0 to 14, with the most acidic = 0 and the most basic = 14.
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Calculating pH
• pH = the negative of the logarithm of the hydrogen ion concentration.
pH = -log [H+]
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Calculating pOH
• pOH is the red-headed stepchild…
• pOH is the negative logarithm of the [OH–].
pOH = -log [OH–]
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pH and pOH
• You can calculate the pH or the pOH of a solution using the log function key on a calculator.
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pH and Acidity
• Acidic solution: pH < 7.0 and[H+] is greater than 1 × 10−7M
• Neutral solution: pH = 7.0 and[H+] is equal to 1 × 10 −7M
• Basic solution: pH > 7.0 and[H+] is less than 1 × 10 −7M
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pH and Significant Figures
• Express [H+] and [OH–] in scientific notation.
• Express pH and pOH with the same number of digits to the right of the decimal place as the number of significant digits in the scientific notation.
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Sample Problem 19.2
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The Solution
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Measuring pH
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Measuring pH
• Two common methods are used:
– Acid base indicators which change color
– pH meters which measure electrical conductivity
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Acid-Base Indicators
• Acid-Base indicators have different colors based upon acidity.
• For each indicator, the change takes place over a relatively narrow range of about 2 pH units.
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Effects of Acidity on Plant Color
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Strengths of Acids and Bases
Essential Question:
How does the value of an acid dissociation constant relate to the strength of the acid,
and how are those values calculated?
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Strong and Weak Acids
• Strong acids are completely ionized in aqueous solution.
HCl + H2O H3O+ + Cl– (100% ionized)
• Weak Acids only slightly ionized in aqueous solution.
• CH3COOH + H2O H3O + CH3COO–
(partially ionized)
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Acid Dissociation – Strong
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Acid Dissociation – Weak
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Acid Dissociation – Very Weak
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Acid Dissociation Constants• Takes the same form as the equilibrium-constant
expression from a balance chemical equation.
• For ethanoic acid, for instance, the acid dissociation constant is calculated as follows:
[H3O+] x [CH3COO–]
Ka =[CH3COOH] x [H2O]
These are sometimes called ionization constants.
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Acid Dissociation Constants
• Weak acids have small Ka values.
• Strong acids have large Ka values.
• The Ka value for HCl (aq) is ∞ (infinite).
• Why do you think this is the case?
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Base Dissociation Constants
• Strong bases and weak bases refer to the degree of dissociation just like acids.
• For sodium hydroxide, for instance, the base dissociation constant is calculated as follows:
[Na+] x [OH–]
Kb =[NaOH] x [H2O]
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Neutralization Reactions
Essential Question:
What are the products of the reaction of an acid and a base when the endpoint of a
titration is reach?
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Acid—Base Reactions
• A mixture of a strong acid with an equal amount of a strong base results in a neutral solution.
• These reactions are called neutralization reactions.
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Neutralization Reactions
• Consider these examples:
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
H2SO4(aq) + 2KOH(aq) K2SO4(aq) + H2O(l)
What would the net ionic equations for each of these reactions be?
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Titration
• Acids and bases sometimes react 1:1
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
However, the ratio can vary
H2SO4(aq) + NaOH(aq) Na2SO4(aq) + 2H2O(aq)
2HCl(aq) + Ca(OH)2(aq) CaCl2 + 2H2O(l)
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Titration
• When an acid and a base are mixed, the equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.
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Titration
• You can determine the concentration of an acid in a solution by performing a neutralization reaction.
• You must select an appropriate acid-base indicator.
• Phenolphthalein turns from colorless to pink as the pH changes from acidic to basic.
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Titrations
• A measured volume of an acid solution of unknown concentration is added to a flask.
• Several drops of the indicator are added to the solution.
• Measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color.