Chapter 16: Acids and Bases

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• Chapter 16 Acids and Bases

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Experimental Definitions

• Acids Bases(alkalis)
• Turn blue litmus red Turn red litmus blue
• taste sour taste bitter
• react with active metals slippery feel
• react with bases to form react with acids to
  salt and water form a salt and water

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Arrhenius Definition

• Acids release hydrogen ions in water
• HCl

H2O
H Cl-

• Bases release hydroxide ions in water

H2O
Na OH-
NaOH
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The Hydronium Ion

• The H ion produced by an acid in water does not
  exist in isolation. It is bonded to a water
  molecule
• H H2O H3O

H3O is called the hydronium ion.
The Lewis dot structure of the hydronium ion is
shown below
The hydronium ion is also hydrated by other water
molecules

HOH

H
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Types of Acids

• Monoprotic acids
• release one hydrogen (proton) ion per formula
  unit.
• HNO3 H NO3-
• Diprotic acids
• Release two protons per formula unit.
• H2SO4 2H SO42-
• Triprotic acids
• Release three protons per formula unit.
• H3PO4 3H PO43-

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Carboxylic Acids

• Not all hydrogens in a compound are acidic.
• In acetic acid, only the hydrogen attached to the
  carboxyl group is acidic

Carboxyl group
H
Acidic hydrogen
Non-acidic hydrogens
Acetate Ion
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Strong and Weak Acids and Bases

• A strong acid or base is 100 ionized
• HCl H Cl-
• NaOH Na OH-
• A weak acid or base is partially ionized,
  equilibrium
• HC2H3O2 H C2H3O2-
• NH3 H2O NH4 OH-

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Aqueous NH3 is sometimes referred to as NH4OH

• NH3 H2O NH4 OH- NH4OH

ammonia
Ammonium ion
Ammonium hydroxide
Hydroxide ion
Although some ammonium hydroxide is formed in
this reaction, the equilibrium lies to the left
and the solution is better referred to as
aqueous ammonia.
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Acids React With Bases.(Neutralization)

• H OH- H2O
• HCl NaOH H2O NaCl
• H NH3 NH4
• HCl NH3 NH4Cl
• HC2H3O2 OH- C2H3O2- H2O
• HC2H3O2 NaOH NaC2H3O2 H2O

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Reactions of Acids

• With active metals
• Form H2(g) and a salt
• H2SO4(aq) Zn(s) H2(g) ZnSO4(aq)
• With metal oxides
• Form H2O(l) and a salt
• 2HCl(aq) CaO(s) H2O(l) CaCl2(aq)
• With carbonates and bicarbonates
• Form CO2(g), H2O(l), and a salt
• CaCO3(s) 2HCl(aq) CO2(g) H2O(l)
  CaCl2(aq)

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Acidic Reactions

• Acids react with metal sulfides
• To form H2S(g) and a salt
• H2SO4(aq) FeS(s) H2S(g) FeSO4(aq)
• Non-metal oxides react with water to form acid
• SO3(g) H2O(l) H2SO4(aq)
• The acids formed this way in the atmosphere may
  fall back to earth as acid precipitation.

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Reactions of Bases

• With transition metal salts
• Form insoluble metal hydroxides and a salt
• 2KOH(aq) Ni(NO3)2(aq) Ni(OH)2(s)
  2KNO3(aq)
• With amphoteric hydroxides
• Amphoteric hydroxides will react with acids or
  bases
• Al(OH)3(s) NaOH(aq) NaAl(OH)4(aq)
• Al(OH)3(s) 3HCl(aq) AlCl3(aq) 3H2O(l)
• With amphoteric metals
• Form hydrogen gas and soluble metal complexes.
• 2Al(s) 2NaOH(s) 6H2O(l) 2NaAl(OH)4(aq)3H2(g
  )

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Definitions of Acids and Bases

• Brønsted - Lowry
• Acid proton (H) donor
• Base proton acceptor
• HF H2O H3O F-
• acid base c.a. c.b.

• NH3 H2O OH- NH4
• base acid c. b. c. a.

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Brønsted - Lowry

• The driving force of the Brønsted Lowry
  acid-base reaction is the formation of weaker
  acids and weaker bases.
• The products formed are conjugate acids (c. a.)
  and conjugate bases (c. b.)
• The Brønsted Lowry acid and its conjugate base
  and the Brønsted Lowry base and its conjugate
  acid are called a conjugate acid-base pair.

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Conjugate acids and bases

• Identify conjugate Acid-Base pairs in an
  acid-base reaction. Complete reactions
  containing conjugate acid-base pairs.
• Take a hypothetical reaction between an acid HA
  and a base B-.
• HA B- A- HB
• In the forward reaction HA acts as an acid by
  donating a proton to the base, B-.
• In the reverse reaction, HB acts as an acid by
  donating the proton back to A-, the base.

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Conjugate Acid-Base Pairs

• The acid from the forward reaction, HA becomes
  A-, the base for the reverse reaction.
• The base from the forward reaction, B- becomes
  the acid, HB for the reverse reaction.
• HA and A- are a conjugate acid-base pair.
• B- and HB are a conjugate acid-base pair.

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Conjugate Acids and Bases

• Every acid has a conjugate base, and every base
  has a conjugate acid.
• The conjugate base of the pair has one fewer H
  and one more negative charge than the acid.
• The conjugate acid of the pair has one more H and
  one fewer negative charge than the base.

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Conjugate acid-base reactions

• A BrØnsted-Lowry acid-base reaction occurs when
  an acid and a base react to form their conjugate
  base and conjugate acid, respectively
• acid1 base2 base1 acid2
• For example, in the reaction
• HF H2O H3O F-
• acid base c.a. c.b.
• HF F- are a conjugate pair.
• H2O H3O are a conjugate pair.

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Definitions of Acids and Bases Continued

• Lewis
• Acid electron pair acceptor
• Base electron pair donor
• with previous examples
• H H
• H-N H- Ö ? H - N - H
  Ö
• H H
  H H
• L.B. L.A.
• e p donor e p acceptor

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Another Example of a Lewis Acid-Base Reaction

• Cl H Cl
  H
• Cl-Al N-H Cl-Al N-H
• Cl H Cl
  H
• L. A. L. B. Adduct
• ep acceptor ep donor

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Self-Ionization of Water

• Write a chemical equation for the autoionization
  of water.
• Water itself is slightly ionized in the liquid
  state The autoionization may be written two
  different ways
• 2 H2O (l) H3O(aq) OH-(aq)
• or
• H2O (l) H (aq) OH-(aq)

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Extent of Ionization of Water

• The extent of ionization of water is given by the
  expression Kw
• Kw H3O(aq)OH-(aq)
• at 25oC, Kw 1.0 x 10-14.
• H3O(aq) Kw/OH-(aq)
• OH-(aq) Kw/ H3O(aq)

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Ion Product Constant of Water, Kw

• Kw HOH or
• Kw H3OOH-
• In a neutral solution, H OH- and
• H3O OH-
• At 25oC,, Kw 1 x 10-14 HOH-
• In a neutral solution, substituting H for
  OH-
• 1 x 10-14 HH H2 ?1 x 10-14
  ?H2

H 1 x 10-7 OH-
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Balance between H3O OH-

• In a neutral solution,
• H3O(aq) OH-(aq) 1.0 x 10-7 M
• In basic solution OH-(aq) gt H3O(aq)
• OH-(aq) gt 1.0 x 10-7 M
• H3O(aq) lt 1.0 x 10-7 M
• In acidic solution H3O(aq) gt OH-(aq)
• OH-(aq) lt 1.0 x 10-7 M
• H3O(aq) gt 1.0 x 10-7 M

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Relation between OH- H3O

• When H3O(aq) increases, OH-(aq) decreases.
• When H3O(aq)decreases, OH-(aq) increases.
• When OH-(aq) increases, H3O(aq) decreases.
• When OH-(aq)decreases, H3O(aq) increases.

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pH and pOH

• Given one of the following, calculate the others
  H(aq) H3O(aq) , OH-(aq), pH, pOH.
• Definition pX -log(X)
• pH -logH(aq) -logH3O(aq)
• pOH - logOH-(aq)

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Some relationships

• Kw H3O(aq)OH-(aq)
• pKw - logKw -logH3O(aq)OH-(aq)
• -log(1.0x10-14) -log H3O(aq)-logOH-(aq)
• 14 pH pOH pH 14 - pOH pOH 14 - pH
• H3O(aq) Kw/OH-(aq) (10-14)/OH-(aq)
• OH-(aq) Kw/H3O(aq) (10-14)/H3O(aq)

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Sample Problem

• The pH of a solution is 9.40. Calculate the
  H(aq) H3O(aq) OH-(aq) and pOH.
• pH -logH(aq) logH(aq) -pH
• H(aq) antilog(-pH) 10-pH 10-9.40
• 4.0 x 10-10 M H3O(aq)
• OH-(aq) (10-14)/ H3O(aq) 10-14/ 4.0 x
  10-10
• 2.5 x 10-5 M
• pOH 14 - pH 14 - 9.40 4.60
• pOH - logOH-(aq) - log(2.5 x 10-5 M) 4.60

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Acidic, Basic, and Neutral Solutions

• Given pH, determine whether a solution is acidic,
  neutral or basic (alkaline).
• If pH 7, H(aq) OH-(aq) and the solution
  is neutral.
• If pH lt 7, H(aq) gt OH-(aq) and the solution
  is acidic.
• If pH gt 7, H(aq) gt OH-(aq) and the solution
  is basic or alkaline.

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For Example

• The pH of a solution is 9.29. The solution is
  basic or alkaline.
• The pH of another solution is 2.13. The solution
  is acidic.
• The pH of yet another solution is 7.00 at 25oC.
  The solution is neutral.

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pH Chart
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pH Properties
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Measurement of pH

• pH indicator
• A species whose color is different in acid and in
  base, which is used to monitor the pH of a
  solution.
• An indicator is a weak acid or base which has a
  different color for its acid form or its
  conjugate base form.
• More accurate pH measurements are made with
  electronic devices called pH meters.

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Hydrolysis Salts in Water

• When salts dissolve in water acid, basic, or
  neutral solutions may be formed. The process is
  called hydrolysis
• The pH of the resulting solution depends on the
  combination of acid and base used in forming the
  salt

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Solution Formed From Salt
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Example Problem

• Does NaC2H3O2 form an acidic, basic, or neutral
  solution when dissolved in water?
• NaC2H3O2, sodium acetate is a salt of a weak acid
  (acetic acid) and a strong base (NaOH).
• HC2H3O2 NaOH NaC2H3O2 H2O
• When the salt dissolves, the acetate ion reacts
  with water to form OH- ion
• NaC2H3O2 Na C2H3O2-
• C2H3O2- H2O HC2H3O2 OH-

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Buffers

• Buffers are solutions which keep the pH of a
  solution almost constant when a strong acid or
  base are added.
• Buffers are prepared by mixing a weak acid and
  the salt of the acid or by mixing a weak base and
  the salt of the base.

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How the acetic acid/sodium acetate buffer works.

• The solution contains HC2H3O2 (which acts as an
  acid) and C2H3O2- (which acts as a base)
• When base (OH-) is added, it is neutralized by
  HC2H3O2
• OH- HC2H3O2 C2H3O2- H2O
• When acid (H) is added, it is neutralized by
  C2H3O2-
• H C2H3O2- HC2H3O2

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Acid-Base Titration

• The concentration or amount of acid or base in a
  sample may be determined by a neutralization
  reaction
• When determining the concentration of a an acid,
  a carefully measured volume of the acid is added
  to a flask.
• A few drops of acid-base indicator are added to
  the flask.
• A standard base, of known concentration is added
  slowly to the flask from a buret until just one
  drop of the added base changes the color of the
  indicator.
• This point is called the end point of the
  titration.

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Buret Used in Acid-Base Titrations
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Titration

• At the end point an equivalent amount of base has
  been added. This is sometimes called the
  equivalence point.
• At the equivalence point, the moles of H in the
  acid is equal to the moles of OH- added.

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Example Problem

• Calculate the concentration of an HCl solution
  determined by a titration in which 50.0 mL of the
  solution is added to a flask. Indicator is
  added, and 29.6 mL of 0.967 M NaOH are required
  to reach the end point.
• HCl NaOH NaCl H2O

mol HCl mol NaOH (0.967 mol/L)(29.6mL)(.001L/m
L)
0.0274 mol
M(NaOH) mol/L (0.0274 mol)/(50.0mL)(0.001L/mL
)
0.548 mol/L 0.548M