THE HYDRONIUM ION

• The proton does not actually exist in aqueous solution as a bare H+ ion.

• The proton exists as the hydronium ion (H3O+).

• Consider the acid-base reaction:

HCO3- + H2O H3O+ + CO3

2-

Here water acts as a base, producing the hydronium ion as its conjugate acid. For simplicity, we often just write this reaction as:

HCO3- H+ + CO3

2-

Conjugate Acid-Base pairs

• Generalized acid-base reaction:HA + B A + HB

• A is the conjugate base of HA, and HB is the conjugate acid of B.

• More simply, HA A- + H+

HA is the conjugate acid, A- is the conjugate base

• H2CO3 HCO3- + H+

AMPHOTERIC SUBSTANCE• Now consider the acid-base reaction:

NH3 + H2O NH4+ + OH-

In this case, water acts as an acid, with OH- its conjugate base. Substances that can act as either acids or bases are called amphoteric.

• Bicarbonate (HCO3-) is also an amphoteric

substance:

Acid: HCO3- + H2O H3O+ + CO3

2-

Base: HCO3- + H3O+ H2O + H2CO3

0

Strong Acids/ Bases

• Strong Acids more readily release H+ into water, they more fully dissociate– H2SO4 2 H+ + SO4

2-

• Strong Bases more readily release OH- into water, they more fully dissociate– NaOH Na+ + OH-

Strength DOES NOT EQUAL Concentration!

Acid-base Dissociation• For any acid, describe it’s reaction in water:

– HxA + H2O x H+ + A- + H2O

– Describe this as an equilibrium expression, K (often denotes KA or KB for acids or bases…)

• Strength of an acid or base is then related to the dissociation constant Big K, strong acid/base!

• pK = -log K as before, lower pK=stronger acid/base!

][

]][[

AH

HAK

x

x

• LOTS of reactions are acid-base rxns in the environment!!

• HUGE effect on solubility due to this, most other processes

Geochemical Relevance?

Organic acids in natural waters• Humic/nonhumic – designations for organic

fractions, – Humics= refractory, acidic, dark, aromatic, large –

generally meaning an unspecified mix of organics– Nonhumics – Carbohydrates, proteins, peptides,

amino acids, etc.

• Aquatic humics include humic and fulvic acids (pKa>3.6) and humin which is more insoluble

• Soil fulvic acids also strongly complex metals and can be an important control on metal mobility

pH• Commonly represented as a range between

0 and 14, and most natural waters are between pH 4 and 9

• Remember that pH = - log [H+]– Can pH be negative?– Of course! pH -3 [H+]=103 = 1000 molal?– But what’s ?? Turns out to be quite small

0.002 or so…– How would you determine this??

pH

• pH electrodes are membrane ion-specific electrodes

• Membrane is a silicate or chalcogenide glass

• Monovalant cations in the glass lattice interact with H+ in solution via an ion-exchange reaction:

H+ + Na+Gl- = Na+ + H+Gl-

The glass• Corning 015 is 22% Na2O, 6% CaO, 72%

SiO2

• Glass must be hygroscopic – hydration of the glass is critical for pH function

• The glass surface is predominantly H+Gl- (H+ on the glass) and the internal charge is carried by Na+

glass

H+Gl-

H+Gl-

H+Gl-

H+Gl-

H+Gl-

H+Gl-

H+Gl-H+Gl-

Na+Gl-

Na+Gl-

E1 E2

Analyte solution Reference solution

pH = - log {H+}; glass membrane electrode

pH electrode has different H+ activity than the solution

SCE // {H+}= a1 / glass membrane/ {H+}= a2, [Cl-] = 0.1 M, AgCl (sat’d) / Ag

ref#1 // external analyte solution / Eb=E1-E2 / ref#2

E1 E2

H+ gradient across the glass; Na+ is the charge carrier at the internal dry part of the membrane

soln glass soln glass

H+ + Na+Gl- Na+ + H+Gl-

Values of NIST primary-standard pH solutions from 0 to 60 oC

pH = - log {H+}

K = reference and junction potentials

pKx?

• Why were there more than one pK for those acids and bases??

• H3PO4 H+ + H2PO4- pK1

• H2PO4- H+ + HPO4

2- pK2

• HPO41- H+ + PO4

3- pK3

BUFFERING

• When the pH is held ‘steady’ because of the presence of a conjugate acid/base pair, the system is said to be buffered

• In the environment, we must think about more than just one conjugate acid/base pairings in solution

• Many different acid/base pairs in solution, minerals, gases, can act as buffers…

Henderson-Hasselbach Equation:

• When acid or base added to buffered system with a pH near pK (remember that when pH=pK HA and A- are equal), the pH will not change much

• When the pH is further from the pK, additions of acid or base will change the pH a lot

][

][log

HA

ApKpH

Buffering example

• Let’s convince ourselves of what buffering can do…

• Take a base-generating reaction:– Albite + 2 H2O = 4 OH- + Na+ + Al3+ + 3 SiO2(aq)

– What happens to the pH of a solution containing 100 mM HCO3- which starts at pH 5??

– pK1 for H2CO3 = 6.35

• Think of albite dissolution as titrating OH- into solution – dissolve 0.05 mol albite = 0.2 mol OH-

• 0.2 mol OH- pOH = 0.7, pH = 13.3 ??

• What about the buffer??– Write the pH changes via the Henderson-Hasselbach

equation

• 0.1 mol H2CO3(aq), as the pH increases, some of this starts turning into HCO3-

• After 12.5 mmoles albite react (50 mmoles OH-):– pH=6.35+log (HCO3-/H2CO3) = 6.35+log(50/50)

• After 20 mmoles albite react (80 mmoles OH-):– pH=6.35+log(80/20) = 6.35 + 0.6 = 6.95

][

][log

HA

ApKpH

Greg Mon Oct 11 2004

0 10 20 30 40 50 60 70 80 90 1005

5.5

6

6.5

7

7.5

8

8.5

Albite reacted (mmoles)

pH

Bjerrum Plots

• 2 D plots of species activity (y axis) and pH (x axis)

• Useful to look at how conjugate acid-base pairs for many different species behave as pH changes

• At pH=pK the activity of the conjugate acid and base are equal

pH0 2 4 6 8 10 12 14

log

ai

-12

-10

-8

-6

-4

-2H2S

0HS-

S2-

H+OH-

7.0 13.0

Bjerrum plot showing the activities of reduced sulfur species as a function of pH for a value of total reduced sulfur of 10-3 mol L-1.

pH0 2 4 6 8 10 12 14

log

ai

-8

-7

-6

-5

-4

-3

-2

6.35 10.33H2CO3* HCO3- CO3

2-

H+

OH-

Common pHrange in nature

Bjerrum plot showing the activities of inorganic carbon species as a function of pH for a value of total inorganic carbon of 10-3 mol L-1.

In most natural waters, bicarbonate is the dominant carbonate species!

Titrations• When we add acid or base to a solution

containing an ion which can by protonated/deprotonated (i.e. it can accept a H+ or OH-), how does that affect the pH?

pH0 2 4 6 8 10 12 14

log

ai

-8

-7

-6

-5

-4

-3

-2

6.35 10.33H2CO3* HCO3- CO3

2-

H+

OH-

Common pHrange in nature

Carbonate System Titration

• From low pH to high pH

Greg Wed Oct 06 2004

0 5 10 15 20 25 30 35 40 45 502

3

4

5

6

7

8

9

10

11

12

NaOH reacted (mmoles)

pH

Greg Wed Oct 06 2004

0 5 10 15 20 25 30 35 40 45 50-16

-14

-12

-10

-8

-6

-4

-2

NaOH reacted (mmoles)

So

me

sp

eci

es

w/

HC

O3- (

log

act

ivit

y) CO2(aq) CO3--HCO3

-

Titrations precipitate

Greg Wed Oct 06 2004

2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7-6

-5.5

-5

-4.5

-4

-3.5

pH

Som

e m

iner

als

(log

mol

es)

Fe(OH)3(ppd)

Boehmite

BJERRUM PLOT - CARBONATE• closed systems with a specified total carbonate

concentration. They plot the log of the concentrations of various species in the system as a function of pH.

• The species in the CO2-H2O system: H2CO3*, HCO3-,

CO32-, H+, and OH-.

• At each pK value, conjugate acid-base pairs have equal concentrations.

• At pH < pK1, H2CO3* is predominant, and accounts for nearly 100% of total carbonate.

• At pK1 < pH < pK2, HCO3- is predominant, and accounts for

nearly 100% of total carbonate.

• At pH > pK2, CO32- is predominant.

pH0 2 4 6 8 10 12 14

log

ai

-8

-7

-6

-5

-4

-3

-2

6.35 10.33H2CO3* HCO3- CO3

2-

H+

OH-

Common pHrange in nature

Bjerrum plot showing the activities of inorganic carbon species as a function of pH for a value of total inorganic carbon of 10-3 mol L-1.

In most natural waters, bicarbonate is the dominant carbonate species!