Chem 1310: Introduction to physical chemistry Part 4: Acids and bases in water

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Chem 1310 Introduction to physical chemistry
Part 4 Acids and bases in water

• Peter H.M. Budzelaar

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About acids and bases

• Acids and bases play a key role in chemistry, in
  particular in water (including in living things).
• Acid-base reactions are fairly simple examples of
  chemical equilibria. Because they are so
  important, they still merit special attention.
• Also, you will see a relation between chemical
  structure and reactivity for the first time here.

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What is an acid?

• Brønsted-Lowry
• An acid is an H donor
• A base is an H acceptor
• Lewis
• An acid is an electron-pair acceptor
• A base is an electron-pair donor
• The Lewis definition is the more general one. In
  water, the two are nearly equivalent, and talking
  about H is easier than figuring out where the
  electrons go, so we mostly use Brønsted-Lowry.

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What is an acid? (2)

• In water, we never have free H but rather H3O.
  This is what you will see in all equations.
  However, acid-base reactions still involve
  transfer of a proton.
• A "free proton" is so electron-poor and reactive
  it will attach itself to anything it encounters.

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Acids and bases in water

• An acid donates a proton to water, forming H3O.
• H2S H2O ? HS- H3O
• The acid can donate a proton.
• Its conjugate base can accept a proton.

H2S is a neutral acidH2O is a neutral base.HS-
is an anionic base.H3O is a cationic acid.
acid
conjugate base
acid
conjugate base
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Acids and bases in water (2)

• NH4 H2O ? NH3 H3O
• HSO4- H2O ? SO42- H3O

acid
conjugate base
NH4 a cationic acidHSO4- is an anionic acid.
acid
conjugate base
acid
conjugate base
acid
conjugate base
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Acids and bases in water (3)

• The reaction doesn't have to be with water
• HSO4- NH3 ? SO42- NH4
• But this is just a combination of the equations
  on the previous slide.

acid
conjugate base
acid
conjugate base
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Acids, bases and equilibria

• For acid-base reactions (in water), we don't
  worry about kinetics.
• Equilibria are established instantly.

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What about bases?

• A base abstracts a proton from water, forming
  OH-.
• H2O NH3 ? OH- NH4
• It doesn't matter whether you are talking about
  an acid and its conjugated base, or a base and
  its conjugated acid...

conjugate acid
base
conjugate acid
base
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What about bases?

• H2O H2O ? OH- H3O
• Water is an acid and a base!

conjugate acid
base
conjugate acid
base
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Acid and base strength

• In water, you cannot have acids stronger than
  H3O
• They simply protonate water, forming H3O
  quantitatively.
• We call them "strong acids".
• Nor can you have bases stronger than OH-.
• They simply deprotonate water, forming OH-
  quantitatively.
• We call them "strong bases".

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Water is a convenient reference

• There will be a scale for acids
• and a similar one for bases

strong
weak
very weak
H3O
H2O
strong
weak
very weak
H2O
OH-
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Water as a reference(H3O/H2O/OH-)

• H2SO4 H2O HSO4- H3O
• H2SO4 is a very strong acid (gtH3O)
• HSO4- is a very weak base, weaker than H2O
• H2CO3 H2O HCO3- H3O
• H2CO3 OH- HCO3- H2O
• H2CO3 is a weaker acid than H3O
• HCO3- is a weaker base than OH-
• CH4 OH- CH3- H2O
• CH4 is a very weak acid, CH3- a very strong base.

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Water as a reference(H3O/H2O/OH-)

• H2O H2O ? OH- H3O
• KW depends on temperature(factor of 10 over
  30).
• Always valid, also in presenceof added acids and
  bases.
• If we know H3O, we also know OH-
  KW/H3O.No need to calculate separately.

H2O is the"almost pure"solvent
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Water as a reference(H3O/H2O/OH-)

• In pure water, H3O OH-, so
• x H3O OH-, x2 10-14, x 10-7 mol/L.
• We call a solution
• neutral if H3O OH- 10-7 mol/L.
• acidic if H3O gt OH- (H3O gt 10-7, OH- lt
  10-7)
• basic if H3O lt OH- (H3O lt 10-7, OH- gt
  10-7)

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pH and acidity

• pH is defined as
• pH - log H3O ( H3O 10-pH )
• base 10 logarithm, don't confuse with "natural
  logarithm" ln!
• Defined in this way to have a convenient
  range,normally 0-14
• pH 7 H3O OH- 10-7 mol/L, neutral
• pH 0 H3O 100 1 mol/L, very acidic (1M
  acid!)
• pH 14 H3O 10-14 mol/L, OH- 1
  mol/L,very basic (1M base!)

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pH and acidity

• Since H3OOH- 10-14,log H3O log OH-
  -14 -pH log OH-pH 14 log OH- 14
  - pOH

this isnearly alwaysnegative!
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Measuring pH

• Electronic ("pH meter")Just put electrode in
  solution, read out the pH. Needs to be
  calibrated, but is fairly accurate.
• Indicator (solution or paper)Contains an organic
  base that changes colour on protonation (or acid
  that changes colour on deprotonation).Convenient,
  easy to carry, limited accuracy(ca 1 pH unit).

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A scale for acid strength

• We characterize the strength of an acidby its Ka
  value
• HA H2O ? A- H3O
• The strength of its conjugate baseis given by
  its Kb value
• A- H2O ? HA OH-

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A scale for acid strength

• You do not need separate values for Ka and
  Kb!(for the same acid/base pair)
• Whether a Ka for an acid is tabulated, or a Kb
  for its conjugated base, is a matter of
  convention and convenience.
• Often, you will find tables of pKa ( - log Ka)
  and pKb ( - log Kb).

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pH calculationsfor strong acids and bases

• Strong acids and bases are always fully
  dissociated, so we know immediately how much H3O
  or OH- is generated.

0.02 M CsOH solution. OH- is a strong base. OH-
210-2 mol/L,H3O 510-13 mol/L, pH 12.3
0.13 M HClO4 solution. HClO4 is a strong acid.
H3O 0.13 mol/L,pH 0.9
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pH calculationsfor strong acids and bases

• Take care! We were neglecting the contribution of
  water itself, assuming all H3O/OH- came from the
  added acid/base.
• For very low concentrations this is no longer
  true, and we need to use the "x method".

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Low concentrations ofstrong acids and bases

• 610-7 mol/L HNO3. Originally, we hadH3O
  OH- 10-7 mol/L. We add 610-7 mol/L H3O,
  but some will be consumed by reactionwith OH-.

H3O OH-
initial 10-7610-7 10-7
change -x -x
equilibrium 710-7-x 10-7-x
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Low concentrations ofstrong acids and bases

• Calculation
• Final H3O 6.210-7, pH 6.2.

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pH calculationsfor weak acids and bases

• Cannot assume all acid has dissociated.Need to
  use the equilibrium expression for Ka.

0.022 M acetic acid (Ka 1.810-5).
HOAc OAc- H3O
initial 0.022 0 (10-7)
change -x x x
equilibrium 0.022-x x x
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pH calculationsfor weak acids and bases

• Calculation
• Simplified version, according to book (MSJ p791),
• assuming x 0.022 Final pH 3.2
• Check your answer 0.022-6.310-4
  0.021, reasonable (but not ideal) approximation.

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pH calculationsfor weak acids and bases

• Full calculation
• Final pH 3.2

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pH calculationsfor weak acids and bases

• For very weak acids (Ka lt 10-10) or very small
  concentrations of weak acids (HA lt 10-5)we
  would need to take explicitly into account
• partial dissociation of the acid
• auto-ionization of water
• This gets too complicated for solving without a
  computer.

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Types of acids

• Acids have a proton that is easily lost
  (transferred).
• It will be attached to an electronegative atom X
  (typically, at least as electronegative as
  nitrogen).
• Examples
• H-F is acidic
• H-OH is somewhat acidic
• H-OClO3 is very acidic
• Why are some acids weak, others strong?
• There are trends, but there is not a simple,
  single rule.

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H-X acids (X not oxygen)
CH4 NH3 H2O HF
SiH4 PH3 H2S HCl
GeH4 AsH3 H2Se HBr
SnH4 SbH3 H2Te HI
decreasingX-H bond strength(MSJ p352)easier
loss of H
very weak acidsweak acids strong acids
increasingX electronegativity(MSJ p355)more
stable X-
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H-X acids (X not oxygen)

• Second (and third) dissociation is always much
  more difficult than first.
• It is harder to remove H from an anion than from
  a neutral molecule!
• So HS- is a much weaker acid than H2S.

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H-O-X acids

• Stronger if
• X more electronegative
• More oxygen atoms attached to X
• Both stabilize negative charge on anion

H2CO3 HNO3HNO2 H2OH2O2
H4SiO4 H3PO4H3PO3 H2SO4H2SO3 HOClHClO2HClO3HClO4
H3AsO4H3AsO3 H2SeO4H2SeO3 HOBrHBrO2HBrO3
Also transition-metal acids H2CrO4, HMnO4, etc
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Very strong oxo-acids
So what about HNO2, HClO2, H2SO3 ?
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Hydrated metal ions

• Cr(OH2)63 H2O ? Cr(OH2)5(OH)2 H3O
• Cr3 is a Lewis acid, complexeswith Lewis base
  H2O.
• The complexation makes water more acidic.
• Higher charge, smaller radius of metal ionÞ
  stronger Brønsted-acidity of hydrated ion.

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Organic acidscarboxylic acids

• The CO groupis electron-withdrawingand allows
  resonancestabilization. Without it,the OH group
  is hardly acidic.

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Common organic acids
Formic acid(the stuff that hurtswhen an ant
bites)
Acetic acid(vinegar,wine gone bad)
Butyric acid("unwashed" smell)
Stearic acid(the Na salt ishousehold soap)
Benzoic acid(common foodpreservative)
Acetylsalicylic acid(Aspirin)
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The strength of an organic acid(see MSJ p782)

• Formic (our reference)
• Acetic weaker(CH3 donating)
• Fluoroacetic stronger(F withdrawing)
• Trifluoroacetic strong!

Benzoic stronger (resonance) Pyruvic
stronger(withdrawing)
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Curiosities

• HCN Hydrocyanic acid, Prussic acid.Negative
  charge stabilized by N atom.Toxic! (not because
  of acidity)
• HN3 Hydrazoic acid. Nitrogens areless effective
  than oxygens at stabilizing negative charge, but
  hey can do so.Toxic (like HCN).Salts used in
  explosives and airbags.

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Curiosities (2)

• Ascorbic acid (vitamin C).
• Essential food componentand preservative.
• The anion has adelocalized negativecharge and
  severalelectron-withdrawingoxygens.

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Types of bases

• Anions of weak acids.
• Weaker acid Þ stronger conjugate base (KaKb
  Kw!).
• Hydroxides
• Ammonia, amines
• Phosphines are much weaker bases
• Bases are often used for cleaning purposes. They
  break down many organic compounds ("liquid
  plummr"). Don't get strong base on your skin.
• Amines cause the "fishy" smell of fish.

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A special caseamino-acids, zwitterions

• Amino-acids combine an acidic and a basic group
  in the same molecule. In the solid state and in
  neutral polar solvents, they exist as
  zwitterions
• This is not resonance but tautomerism!

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Amino acids

• In acidic solution, they exist in the protonated
  cationic form
• In basic solution, they exist in the deprotonated
  anionic form

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Peptides and proteins

• Peptides (including proteins) are formed by
  coupling amino acids head-to-tail, via peptide
  linkages.
• These linkages are neither acidic nor basic.
• Only the head, tail and some side groups (see MSJ
  p592) are acidic or basic.

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Strength of bases

• Bases must be able to accept a proton, so they
  must have an available electron pair

H2O is a weak base.
NH3 is a stronger base. N less electronegative
Þ electron pair more available
OH- is a strong base, O2- is even
stronger (negative charge!).
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Strength of bases

• In general, provided an atom has an available
  electron pair, basicity goes up
• going up in the periodic table (NH3gtPH3)
• going left in the periodic table (NH3gtH2O)
• with increasing negative charge(PO43-gtHPO42-gtH2PO
  4-)
• with more electron-donating substituents
  (CH3OHgtH2O)

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Solutions of salts

• A salt consists of an anion (conjugate base of an
  acid) and a cation (often an acid itself, or
  conjugate acid of a base). It is formed by
  neutralization of the acid with the base (or vice
  versa).
• If both the original acid and the original base
  are strong, the solution will be neutral (the
  conjugate base and acid are very weak).
• NaCl salt of HCl (strong) and NaOH (strong)
  neutral.

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Solutions of salts

• If the original acid is strong but the base is
  weak, the solution is acidic.
• NH4Cl from HCl NH3. NH4 is (weakly) acidic.
• If the original acid is weak but the base is
  strong, the solution is basic.
• NaOAc from HOAc NaOH. OAc- is (weakly) basic.
• If both acid and base were weak
• If Ka(acid) gt Kb(base), solution is acidic.
• If Ka(acid) lt Kb(base), solution is basic.

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pH calculations for salts

• For a salt of strong acid and weak base need
  only consider conjugate acid of weak base. See
  calculations for weak acids.
• For a salt of weak acid and strong base need
  only consider conjugate base of weak acid. See
  calculations for weak bases.
• For salts of weak acid and weak base need to
  include both, calculation can be complicated.

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Lewis acidity that is notBrønsted-Lowry acidity

• Metal ions are strong Lewis acids. They bind H2O,
  but also other Lewis bases like NH3 or anions of
  various acids.

deep dark blue
light blue
red,insoluble
colourless solution
white,insoluble
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Lewis acidity that is notBrønsted-Lowry acidity

• Other elements can act like Lewis acids
• (but generally the electron pairs move around
  then)