Chapter 18

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Chapter 18 Acids, Bases, and Salts

• Honors Chemistry
• Kings High School
• Mrs. Warren

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Chapter Outline

• Section 1 Properties and Theories
• Section 2 Hydrogen Ions and Acidity
• Section 3 Strengths of Acids Bases
• Section 4 Neutralization Reactions

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Section 18.1

• OBJECTIVES
• Define and identify the properties of acids and
  bases.
• Compare and contrast acids and bases as defined
  by the theories of a) Arrhenius
• b) Brønsted-Lowry
• c) Lewis.

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Learning Check from POGIL

• H2C2O4 (aq) H2O (l) ? H3O (aq) HC2O4 (aq)
• Identify the acid (WHY is this the acid?)
• Identify the base (WHY is this the base?)
• What is the name of this ion H3O (aq)? This ion
  consists of what and what?
• Why are H and proton transfer the same thing?

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Properties of Acids

• They taste sour (dont try this at home Chuckles
  and Giggles!!)
• They can conduct electricity.
• Can be _____or ______ electrolytes in aqueous
  solution
• React with metals to form H2 gas.
• Change the color of indicators.
• React with bases (metallic hydroxides) to form
  water and a salt.

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Properties of Acids

• They have a pH of _______(more on this concept of
  pH in a later lesson)
• How do you know if a chemical is an acid?
• It usually starts with ___________.
• HCl, H2SO4, HNO3, etc.

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Acids Affect Indicators, by changing their color
Blue litmus paper turns red in contact with an
acid (and red paper stays red).
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Acids have a pH less than 7
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Acids React with Active Metals
Acids react with active metals to form salts and
hydrogen gas
HCl(aq) Mg(s) ?
What type of reaction is this?
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Effects of Acid Rain on Marble(marble is calcium
carbonate)
George Washington BEFORE acid rain
George Washington AFTER acid rain
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Acids Neutralize Bases
HCl NaOH ?
-Neutralization reactions ALWAYS produce a salt
(which is an ionic compound) and water.
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Acids Conduct Electricity

• HCl conducts electricity when it is ___________
  in _______.
• It does this because HCl does what in water?

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Properties of Bases (metallic hydroxides)

• React with acids to form water and a salt. (What
  is this rxn called?)
• Taste bitter.
• Feel slippery (dont try this either).
• Can be strong or weak electrolytes in aqueous
  solution
• Change the color of indicators.

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Examples of Bases(metallic hydroxides)

• Sodium hydroxide (lye for drain cleaner soap)
• Potassium hydroxide (alkaline batteries)
• Magnesium hydroxide (Milk of Magnesia)
• Calcium hydroxide(limestone)

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Bases Affect Indicators
Red litmus paper turns blue in contact with a
base (and blue paper stays blue).
Phenolphthalein turns purple in a base.
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Bases have a pH greater than 7
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Bases Neutralize Acids
Milk of Magnesia contains magnesium hydroxide,
Mg(OH)2, which neutralizes stomach acid, HCl.
2 HCl Mg(OH)2
Magnesium salts can cause diarrhea (thus they are
used as a laxative) and may also cause kidney
stones.
MgCl2 2 H2O
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SECTION18.1
Properties of Acids and Bases
The usual solvent for acids and bases is
waterwater produces equal numbers of hydrogen
and hydroxide ions in a process called
self-ionization.
The hydronium ion is H3O.
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Acid-Base Theories
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Svante Arrhenius

• He was a Swedish chemist (1859-1927), and a Nobel
  prize winner in chemistry (1903)
• one of the first chemists to explain the chemical
  theory of the behavior of acids and bases

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The Arrhenius Model

• The Arrhenius model states that an acid is a
  substance that contains hydrogen and ionizes to
  produce hydrogen ions in an aqueous solution.
• A base is a substance that contains hydroxide
  ions and dissociates to produce a hydroxide ion
  in an aqueous solution.
• What does this mean??????

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1. Arrhenius Definition - 1887

• Acids produce hydrogen ions (H1) in aqueous
  solution (HCl ? H1 Cl1-)
• Bases produce hydroxide ions (OH1-) when
  dissolved in water.
• (NaOH ? Na1 OH1-)
• Limited to aqueous solutions.
• Only one kind of base (hydroxides)
• NH3 (ammonia) could not be an Arrhenius base no
  OH1- produced.

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2. Brønsted-Lowry - 1923

• A broader definition than Arrhenius
• Acid is hydrogen-ion donor (H or proton) base
  is hydrogen-ion acceptor.
• Acids and bases always come in pairs.
• When it dissolves in water, it gives its proton
  to water.
• HCl(g) H2O(l) ? H3O(aq) Cl-(aq)

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Johannes Brønsted Thomas Lowry
(1879-1947) (1874-1936)
Denmark England
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Why Ammonia is a Base

• Ammonia can be explained as a base by using
  Brønsted-Lowry
• NH3(aq) H2O(l) ? NH41(aq) OH1-(aq)

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Acids and bases come in pairs

• A conjugate base is the species produced when
  an acid donates H.
• the remainder of the original acid, after it
  donates its hydrogen ion
• A conjugate acid is the species produced when a
  base accepts H.
• the particle formed when the original base gains
  a hydrogen ion

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Acids and bases come in pairs

• General equation is
• HA(aq) H2O(l) ? H3O(aq) A-(aq)
• Acid Base ? Conjugate acid Conjugate base
• NH3 H2O ? NH41 OH1-
• base acid c.a. c.b.
• HCl H2O ? H3O1 Cl1-
• acid base c.a. c.b.
• Amphoteric a substance that can act as both an
  acid and base- as water shows

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Learning Check

• Identify the conjugate acid-base pairs in each
  reaction
• NH41 (aq) OH-1 (aq) ? NH3 (aq) H2O (l)
• HBr (aq) H2O (l) ? H3O1 (aq) Br-1 (aq)
• CO3-2 (aq) H2O (l) ? HCO3-1 (aq) OH-1(aq)
• Challenge
• The products of an acid-base reaction are H3O1
  and SO4-2. Write a balanced equation for the
  reaction and identify the conjugate acid-base
  pairs.

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Monoprotic and Polyprotic Acids

• HCl and HF
• How many hydrogen ions can these acids donate?
• We call these ____________ _______.
• Ionizable Hydrogen Atoms
• Acetic Acid (CH3COOH)
• An ionizable atom is bonded to an electronegative
  atom.
• An acid must have an ionizable hydrogen.

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Organic Acids (those with carbon)
CH3COOH of the 4 hydrogen, only 1 ionizable
The carboxyl group is a poor proton donor, so it
is considered a weak acid.
(due to being bonded to the highly
electronegative Oxygen)
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Polyprotic Acids

• Some compounds have more than one ionizable
  hydrogen to release
• HNO3 nitric acid - monoprotic
• H2SO4 sulfuric acid - diprotic - 2 H
• H3PO4 phosphoric acid - triprotic - 3 H
• Having more than one ionizable hydrogen does not
  mean stronger!

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Steps of Polyprotic Acids

• H3PO4 (aq) H2O (l) ? H3O1 (aq) H2PO4-1 (aq)
• H2PO4-1 (aq) H2O (l) ? H3O1 (aq) HPO4-2 (aq)
• Write the last one.

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3. Lewis Acids and Bases

• Gilbert Lewis focused on the donation or
  acceptance of a pair of electrons during a
  reaction
• Lewis Acid An acid is an ion or molecule with a
  vacant atomic orbital that can accept an electron
  pair i.e. electron pair acceptor
• Lewis Base A base is an ion or molecule with a
  lone electron pair that it can donate i.e.
  electron pair donor

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Examples of Electron Pair Donors and Acceptors
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Review of Theories
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Section 18.2

• OBJECTIVES
• Describe how H1 and OH1- are related in an
  aqueous solution.
• Classify a solution as neutral, acidic, or basic
  given the H1 or OH1- .
• Convert H1 into pH values and OH1- into pOH
  values.
• Describe the purpose of an acid-base indicator.

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Hydrogen Ions from Water

• Water ionizes, or falls apart into ions
• H2O ? H1 OH1-
• Called the self ionization of water
• Since they are equal, a neutral solution results
  from water
• Kw H1 OH1- 1 x 10-14 M2
• H1 OH1- 1 x 10-7 M
• Kw is called the ion product constant for water

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Ion Product Constant

• H2O ? H1 OH1-
• Kw is constant in every aqueous solution H
  x OH- 1 x 10-14 M2
• If H gt 10-7 then OH- lt 10-7
• If H lt 10-7 then OH- gt 10-7
• If we know one, other can be determined
• If H gt 10-7 , it is acidic and OH- lt 10-7
• If H lt 10-7 , it is basic and OH- gt 10-7
• Basic solutions also called alkaline

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Example 1

• The hydrogen ion concentration in a cup of coffee
  is 1.0 x 10-5M. What is the hydroxide
  concentration in the coffee? If the coffee
  acidic, basic, or neutral?

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Example 2 and 3

• Calculate the hydrogen ion or hydroxide ion
  concentration and state whether the solution is
  acidic, basic, or neutral.
• H 1.0 x 10-13 M
• OH- 1.0 x 10-3 M

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The pH concept from 0 to 14

• pH hydrogen power
• definition pH -logH
• in neutral pH -log(1 x 10-7) 7
• in acidic solution
• H gt 10-7
• pH lt -log(10-7)
• pH lt 7 (from 0 to 7 is the acid range)

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Calculating pOH

• pOH -log OH-
• Thus, a solution with a pOH less than 7 is basic
  with a pOH greater than 7 is an acid
• Not greatly used like pH is.
• Whats the relationship between pH and pOH?
• H x OH- 1 x 10-14 M2
• pH pOH 14

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pH and pOH Relationship
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Examples

• What is the pH of a neutral solution at 298K?
• Calculate the pH of a solution having OH- 8.2
  x 10-6M
• Calculate the pH of aqueous solutions with the
  following H
• H 0.0055M
• H 0.000084M

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Examples

• Calculate the pH and pOH of aqueous solutions
  with the following concentrations
• OH- 6.5 x 10-4 M
• H 3.6 x 10-9 M
• What are H and OH- in a healthy persons
  blood that has a pH of 7.40?

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Some of the many pH Indicators and theirpH range
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Acid-Base Indicators

• Although useful, there are limitations to
  indicators
• A pH meter may give more definitive results

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Section 18.3

• OBJECTIVES
• Define strong acids and weak acids.
• Describe how an acids strength is related to the
  value of its acid dissociation constant.
• Calculate an acid dissociation constant (Ka) from
  concentration and pH measurements.
• Order acids by strength according to their acid
  dissociation constants (Ka).
• Order bases by strength according to their base
  dissociation constants (Kb).

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Strength

• Acids and Bases are classified according to the
  degree to which they ionize in water
• Strong are completely ionized in aqueous
  solution this means they ionize 100
• Weak ionize only slightly in aqueous solution
• Strength is very different from Concentration

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Strength

• Strong means it forms many ions when dissolved
  (100 ionization)
• Mg(OH)2 is a strong base- it falls completely
  apart (nearly 100 when dissolved).
• But, not much dissolves- so it is not concentrated

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Strong Acid Dissociation
(makes 100 ions)
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Weak Acid Dissociation (only partially
ionizes)
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Measuring strength

• Ionization is reversible
• HA H2O ? H A-
• This makes an equilibrium
• Acid dissociation constant Ka
• Ka H A-
  HA
• Stronger acid more products (ions), thus a
  larger Ka

(Note that the arrow goes both directions.)
(Note that water is NOT shown, because its
concentration is constant, and built into Ka)
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What about bases?

• Strong bases dissociate completely.
• MOH H2O ? M OH- (M a metal)
• Base dissociation constant Kb
• Kb M OH- MOH
• Stronger base more dissociated ions are
  produced, thus a larger Kb.

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Examples

• See Handout

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Strength vs. Concentration

• The words concentrated and dilute tell how much
  of an acid or base is dissolved in solution -
  refers to the number of moles of acid or base in
  a given volume
• The words strong and weak refer to the extent of
  ionization of an acid or base
• Is a concentrated, weak acid possible?

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Practice

• Write the Ka expression for HNO2
• Equation HNO2 ? H1 NO21-
• Ka H1 x NO21-
  HNO2
• Write the Kb expression for NH3 (as NH4OH)

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- Page 610
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Section 18.4

• OBJECTIVES
• Define the products of an acid-base reaction.
• Explain how acid-base titration is used to
  calculate the concentration of an acid or a base.
• Explain the concept of equivalence in
  neutralization reactions.
• Describe the relationship between equivalence
  point and the end point of a titration.

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Acid-Base Reactions

• Acid Base ? Water Salt
• Properties related to every day
• antacids depend on neutralization
• farmers adjust the soil pH
• human body kidney stones from insoluble salts

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Acid-Base Reactions

• Neutralization Reaction - a reaction in which an
  acid and a base react in an aqueous solution to
  produce a salt and water
• HCl(aq) NaOH(aq) ? NaCl(aq) H2O(l)
• H2SO4(aq) 2KOH(aq) ? K2SO4(aq) 2 H2O(l)

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Titration

• Titration is the process of adding a known amount
  of solution of known concentration to determine
  the concentration of another solution
• Remember? - a balanced equation is a mole ratio
• The equivalence point is when the moles of
  hydrogen ions is equal to the moles of hydroxide
  ions ( neutralized!)

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- Page 614
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Titration

• The concentration of acid (or base) in solution
  can be determined by performing a neutralization
  reaction
• An indicator is used to show when neutralization
  has occurred
• Often we use phenolphthalein- because it is
  colorless in neutral and acid turns pink in base

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Steps - Neutralization reaction

• 1. A measured volume of acid of unknown
  concentration is added to a flask
• 2. Several drops of indicator added
• 3. A base of known concentration is slowly
  added, until the indicator changes color measure
  the volume

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Neutralization

• The solution of known concentration is called the
  standard solution
• added by using a buret
• Continue adding until the indicator changes color
• called the end point of the titration
• Or equivalence point

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Titration Curves
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Salt Hydrolysis

• A salt is an ionic compound that
• comes from the anion of an acid
• comes from the cation of a base
• is formed from a neutralization reaction
• some neutral others acidic or basic
• Salt hydrolysis - a salt that reacts with water
  to produce an acid or base

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Salt Hydrolysis

• Hydrolyzing salts usually come from
• a strong acid a weak base, or
• a weak acid a strong base
• Strong refers to the degree of ionization
• A strong Acid a strong Base Neutral Salt

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Salt Hydrolysis

• To see if the resulting salt is acidic or basic,
  check the parent acid and base that formed it.
  Practice on these
• HCl NaOH ?
• H2SO4 NH4OH ?
• CH3COOH KOH ?

NaCl, a neutral salt
(NH4)2SO4, acidic salt
CH3COOK, basic salt
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Buffers

• Buffers are solutions in which the pH remains
  relatively constant, even when small amounts of
  acid or base are added
• made from a pair of chemicals a weak acid and
  one of its salts or a weak base and one of its
  salts

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Buffers

• A buffer system is better able to resist changes
  in pH than pure water
• Since it is a pair of chemicals
• one chemical neutralizes any acid added, while
  the other chemical would neutralize any
  additional base
• AND, they produce each other in the process!!!

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Buffers

• The buffer capacity is the amount of acid or base
  that can be added before a significant change in
  pH

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Buffers

• The two buffers that are crucial to maintain the
  pH of human blood are
• 1. carbonic acid (H2CO3) hydrogen carbonate
  (HCO31-)
• 2. dihydrogen phosphate (H2PO41-) monohydrogen
  phoshate (HPO42-)