Ch17. Acids and Bases: A Second Look

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Ch17. Acids and Bases A Second Look

• Brady Senese, 4th

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Chapter 17 Acids and Bases

• Acid-base reactions involve proton (hydrogen ion,
  H) transfer
• The generalization of the Arrhenius definition of
  acids and bases is called the Brønsted-Lowry
  definitions
• An acid is a proton donor
• A base is a proton acceptor
• This allows for gas phase acid-base reactions

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• Species that differ by a proton, like H2O and
  H3O, are called conjugate acid-base pairs

The reaction of HCl and H2O. HCl is the acid
because it donates a proton. Water is the base
because it accepts a proton.
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(a) Formic acid transfers a proton to a water
molecule. HCHO2 is the acid and H2O is the base.
(b) When a hydronium ion transfers a proton to
the CHO2- ion, H3O is the acid and formate ion
is the base.
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Brønsted Acid/Base Reactions Transfer H

• Products differ by one H from the reactants to
  form conjugate
• Conjugate acid-base pairs differ by one H.
• HCN(aq) OH-(aq) ? H2O(l) CN-(aq)
• Note that in the conjugate pairs, the acid has
  one more H than its conjugate base

Brønsted base
conjugate acid
Brønsted acid
conjugate base
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Learning Check

• Identify the Conjugate Partner for Each

Cl-
NH4
C2H3O2-
HCN
F-
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• An amphoteric substances can act as either an
  acid or base
• These are also called amphiprotic, and can be
  either molecules or ions
• For example, the hydrogen carbonate ion

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Learning Check

• Write a reaction that shows that H2PO4- is a
  Brønsted acid when reacted with OH-
• H2PO4-(aq) OH-(aq) ?
• Write a reaction that shows that H2PO4- is a
  Brønsted base when reacted with H3O(aq)
• H2PO4-(aq) H3O(aq) ?

H2O(l) HPO42-(aq)
H2O(l) H3PO4(aq)
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• The strength of an acid is a measure of its
  ability to transfer a proton
• Acids that react completely with water (like HCl
  and HNO3) are classified as strong
• Acids that are less than completely ionized are
  called weak acids
• Bases can be classified in a similar fashion
• Strong bases, like the oxide ion, react
  completely
• Weak bases, like NH3, undergo incomplete reactions

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• The strongest acid in water is the hydronium ion
• If a more powerful proton donor is added to
  water, it quantitatively reacts with water to
  produce H3O
• Similarly, the strongest base that can be found
  in water is the hydroxide ion, because more
  powerful proton acceptors react quantitatively
  with water to produce OH-

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Acid/Base Strengths In Aqueous Solution

• Hydronium ion (H3O) is the strongest acid in
  solution stronger acids react completely with
  water to give H3O
• Hydroxide ion (OH-) is the strongest possible
  base in solution stronger bases react completely
  with water to give OH-
• The reaction of all stronger acids and bases in
  water to the same strength is termed leveling
• Acid-base reactions occur in favor of strength
  reduction

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Conjugate Pairs Have Reciprocal Strengths

• The stronger the acid, the weaker its conjugate
  base
• The stronger the base, the weaker its conjugate
  acid
• Strong acids are ionized 100, hence their anions
  are extraordinarily poor bases
• The conjugate bases of most strong acids are
  neutral

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• Acetic acid (HC2H3O2) is a weak acid
• It ionizes only slightly in water
• The hydronium ion is a better proton donor than
  acetic acid (it is a stronger acid)
• The acetate ion is a better proton acceptor than
  water (it is a stronger base)
• The position of an acid-base equilibrium favors
  the weaker acid and base

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• This can be generalized
• Stronger acids and bases tend to react with each
  other to produce their weaker conjugates
• The stronger a Brønsted acid is, the weaker is
  its conjugate base
• The weaker a Brønsted acid is, the stronger is
  its conjugate base
• These ideas can be applied to the binary acids
  (acids made from hydrogen and one other element)

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Identify The Preferred Direction Of The Following

• H3O (aq) CO32-(aq) ? HCO-3(aq) H2O (l)
• NH4(aq) HCO-3(aq) ? NH3(aq) H2CO3(aq)
• Cl- HCN(aq) ? HCl(aq) CN-(aq)

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Periodic Trends Of Binary Acids (HnX )

• As we read left to right in a period, increasing
  electronegativity of X makes the H-X bond more
  polar
• Acid strength increases with increasing polarity
• As we read top to bottom in a group, the acid
  strength increases due to increasing bond length
  of the HX bond due to increased radius of the
  anion, X

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Learning Check

• Which is stronger?
• H2S or H2O
• CH4 or NH3
• HF or HI

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Oxoacids ( A(O)m(OH)n)

• Increase in acid strength as the
  electronegativity of the central atom, A,
  increases
• Increase in acid strength as the number of oxygen
  atoms on (hence the oxidation state of) the
  central atom, A, increases
• Electrical induction through the central atom
  weakens strength of the bond to H

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Learning Check

• Which is stronger?
• H2SO4 or H3PO4
• HNO3 or H3PO4
• H2SO4 or H2SO3
• HNO3 or HNO2

• H2SO4
• HNO3
• H2SO4
• HNO3

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• The strength of an acid can be analyzed in terms
  the the basicity of the anion formed during the
  ionization
• The basicity is the willingness of the anion to
  accept a proton from the hydronium ion
• Consider H2SO4 and H3PO4

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Anions Of Oxoacids Are Basic

• Oxygen atoms are electron withdrawing, thus the
  charge on an anion is located on the lone oxygens
• The more oxygen atoms there are that share the
  same charge, the less basic is the anion
• The stronger the base behavior of the anion, the
  greater the strength of the conjugate acid

-
-
2 O share the (-) charge
3 O share the (-) charge
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• In terms of the percentage of molecules that are
  ionized, sulfuric acid is a stronger acid than
  phosphoric acid
• There is a third definition for acid and bases
• It is a further generalization, or broadening, of
  the species that can be classified as either an
  acid or base
• The definitions are based on electron pairs and
  are called Lewis acids and bases

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Summary

• The strengths of the binary acids increases from
  left to right within the same period
• For example, HCl is stronger acid than H2S which
  is a stronger acid than PH3
• The strengths of the binary acids increase from
  top to bottom within a group
• For example, HI is a stronger acid than HBr which
  is a stronger acid than HCl
• Trends are also present in the oxoacids (acids of
  hydrogen, oxygen, and one other element)

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Summary

• When the central atom holds the same number of
  oxygen atoms, the acid strength increases from
  the bottom to top within a group and from left to
  right within a period
• Acid strength HClO4 gt HBrO4 gt HIO4
• Acid strength HClO4 gt H2SO4 gt H3PO4
• For a given central atom, the acid strength of an
  oxoacid increases with the number of oxygens held
  by the central atom
• Acid strength H2SO4 gt H2SO3

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Lewis Acid/Base Reactions

• Lewis acids accept an electron pair to form
  coordinate covalent bonds
• Lewis bases donate lone pairs of electron to form
  coordinate covalent bonds
• Neutralization is the formation of a coordinate
  covalent bond between the donor and acceptor

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Lewis Acids and Bases

• Lewis acids
• molecules ions with incomplete valence shells
• molecules ions with multiple bonds that can be
  shifted to accept electrons
• molecules or ions with central atoms that can
  accommodate additional electrons
• Lewis bases
• molecules ions that have complete valence
  shells with unshared electrons

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Learning Check

• Identify the Lewis acid and base in the following
• NH3 H ?NH4
• F- BF3 ? BF4-

-
-

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Carbon dioxide (a Lewis acid) reacts with
hydroxide ion (a Lewis base) in solution to form
the bicarbonate ion. The electrons in the
coordinate covalent bond come from the oxygen
atom in the hydroxide ion.
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• Lewis acids
• Molecules or ions with incomplete valence shells
  (for example BF3 or H)
• Molecules or ions with complete valence shells,
  but with multiple bonds that can be shifted to
  make room for more electrons (for example CO2)
• Molecules or ions that have central atoms capable
  of holding additional electrons (usually, atoms
  of elements in Period 3 and below, for example
  SO2)

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• Lewis bases
• Molecules or ions that have unshared pairs of
  electrons and that have complete shells (for
  example O2- or NH3)
• All Brønsted acids and bases are Lewis acids and
  bases, just like all Arrhenius acids and bases
  are Brønsted acids and bases
• Consider a proton transfer from the Lewis
  perspective

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• For example, the proton transfer between the
  hydronium ion and ammonia

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• In general, most metal oxides react with water to
  form bases, and nonmetal oxides react with water
  to form acids
• In Section 5.5 metal oxides were called base
  anhydrides and nonmetal oxides were called acid
  anhydrides
• When cations dissolve in water, they form species
  called hydrated ions
• Hydrated metal ions tend to be Brønsted acids

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• For the monohydrate of the metal ion Mn the
  equilibrium can be represented as

The metal ion makes the hydrogens on the water
more acidic.
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Hydrated Metal Ions Can Act as Weak Acids

• Electron deficiency of metal cations causes them
  to induce electrons towards it from the water of
  hydration
• M(H2O)m n H2O?M(H2O)m-1OH(n-1) H3O
• The higher the charge density, the more acidic
  the metal.
• Acidity increases left to right in a period.
• Acidity decreases top to bottom in a group.

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• The charge density of a cation is its charge
  divided by its volume
• The higher the charge density, the better a
  cation is at drawing electron density from a O-H
  bond and the more acidic it is
• Within a given period, the cation size increases,
  and the charge density decreases, from top to
  bottom
• As a result, the most acidic hydrated cations are
  found at the top of a group
• As the cation charge increases, it becomes more
  acidic

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• When the charge (oxidation number) is small, its
  oxide tends to be basic
• When the cation ion charge is 3, the oxide
  begins to become acidic
• An amphoteric species is capable of acting as
  both an acid and base
• Aluminum oxide is an example of an amphoteric
  compound

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• Many nonmetal oxides are acid anhydrides
• For example

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• Water undergoes self-ionization or autoionization
  making it a weak electrolyte
• This equilibrium is described by the ion product
  of water

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• At 25oC in pure water it has been found that
• In any aqueous solution, the product of H and
  OH- equals Kw
• This provides an alternate a way to define the
  acidity or basicity of a solution

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• Neutral solutions H3O OH- or H
  OH-
• Acidic solutions H3O gt OH- or H gt
  OH-
• Basic solutions H3O lt OH- or H lt
  OH-
• To make the comparison of small values of H
  easier, the pH was defined
• In terms of the pH
• Neutral solutions pH 7.00
• Acidic solutions pH lt 7.00
• Basic solutions pH gt 7.00

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Learning Check

• Complete the following with the missing data

11.5
3.110-12
4.310-10
4.60
6.710-13
12.2
3.9210-9
5.593
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The pH of some common solutions. H decreases,
while OH- increases, from top to bottom.
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Indicators Help Us Estimate pH
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• The pH of a solution can be measured with a pH
  meter or estimated using a visual acid-base
  indicator (see Table 17.3, page 762)
• An acid-base indicator is a species that changes
  color based on the pH
• Calculating the pH of a strong acid or base is
  easy because they are 100 dissociated in
  aqueous
• For example, the pH of 0.10 M HCl is 1.00 and the
  pH of 0.10 M NaOH is 13.00

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• In the last example it was assumed that the total
  concentration of H was due to the strong acid
  (HCl) and OH- was due to the strong base (NaOH)
  
• This assumption is valid because the
  autoionization of water is suppressed in strongly
  acidic or strongly basic solutions
• This assumption fails for very dilute solutions
  of acids or bases (less than 10-6 M)

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Strong Acids Ionize 100 in Water

• As the substances are placed into water, they
  form H3O .
• The H3O formed by the acid suppresses waters
  ionization. (if acid gt 10-7 M)
• The pH can be calculated from the concentration
  of H3O produced by the strong acid
• The reaction of strong acids occurs irreversibly,
  so we show the reaction with a ? instead of using
  a double arrow

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Learning Check

• What is the pH of 0.1M HCl

• HCl(aq) H2O(l) ?H3O(aq) Cl-(aq)
• 0.1 N/A 0 0 I
• -0.1 -0.1 0.1 0.1 C
• 0 N/A 0.1 0.1 end
• pH -log(0.1) 1

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Strong Bases Dissociate 100 In Water

• They are strong electrolytes that form OH- when
  dissolved
• pOH can be calculated from the OH- from the
  solution
• Waters contribution is negligible if the base is
  sufficiently concentrated OH-gt10-7M

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Learning Check

• What is the pH of 0.5M Ca(OH)2?

• Ca(OH)2(aq) ? Ca2(aq) 2OH-(aq)
• 0.5 0 0 I
• -0.5 0.5 0.52 C
• 0 0.5 1.0 end
• pOH -log(1.0) 0
• pH 14

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Auto-ionization of Water (Kw)

• Water ionizes to a very small extent (Kw10-14 at
  room temperature) according to the following
  reaction
• H2O(l) H2O(l) ? H3O(aq) OH-(aq)
• Since water is present in all aqueous solutions,
  the Kw equilibrium exists in all aqueous
  solutions.
• KwH3O OH-
• Kw 10-14 at 25C
• When H3OOH-, the solution is neutral.

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pH and Kw

• pH is defined for aqueous solutions only, and is
  temperature dependent, because Kw is
• pH-logH3O
• It derives from the auto ionization of water.
• KwH3OOH-
• log(Kw)logH3O logOH-
• -log(Kw)-logH3O - logOH-
• pKwpHpOH
• pHgt7 is basic pH7 is neutral pHlt7 is acidic

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Using pH

• pH-logH3O
• To find H3O
• -pH logH3O
• 10-pH H3O
• When the pH of a solution is 12.2,
• 10-12.2 H3O
• 6.310-13 M
• note that we lose a significant figure when we
  take the antilog

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Learning Check

• Complete the following with the missing data

11.67
4.710-3
12.11
7.7610-13