Week 2 Class Notes

Oxygen in Water

Solubility (Equilibrium Concentration )

At 25^oC in equilibrium with air, oxygen solubility is 8.32 mg/l. (How much gas is this??)
1 mole of O₂ occupies 22.4 liters at 0 ^oC and 1 atmosphere pressure, and weighs 32 g.

Actual concentration depends on kinetics as well as equilibrium.
(What is rate of dissolution as compared with rate of consumption in the water?)
** Rate of dissolution depends strongly on contact and amount of surface between air and water.
** Rate of consumption depends on reactions in water which use O₂.
{CH₂O} + O₂ ® CO₂ + H₂O
(degradation of organic material)

How many grams of O₂ are used in consuming 1 gram of organic matter? (1 "mol" CH₂O = 30 g)

What is the qualitative effect of raising temperature on O₂ concentration in water?

Acidity--Capability to neutralize OH^-

Alkalinity--Capability to neutralize H⁺

pH -- Expresses concentration of H⁺
(and also OH^- from K_w)

Strong acids (free mineral acids)

Hydrochloric HCl

Nitric HNO₃

Sulfuric H₂SO₄All totally ionize in water--

Weak acids

Acetic and other carboxylic acids ---COOH group

Carbonic acid H₂CO₃

Acidity of some hydrated metals:

Al(H₂O)₆³⁺ Û Al(H₂O)₅OH²⁺ + H⁺

Buffers

Definition of a buffer

Composed usually of weak acid and its salt or weak base and its salt

Most common buffer in natural waters involves carbonate/bicarbonate/carbonic acid/carbon dioxide system

When CO₂ in air at 350 PPM equilibrates with water, the concentration [CO₂]_(aq) at 25 ^oC is 1.146 x 10^-5M. From the equations we can calculate the pH of this water.

CO₂ + H₂O Û HCO₃^- + H⁺

Can use K_a1, since we suspect that the pH will be below pK_a1

Alkalinity vs Basicity

A solution above pH 7can be called BASIC or ALKALINE.
These are not exactly the same thing--A solution of weak base like bicarbonate has a much lower pH than an equal concentration of a strong base such as NaOH, but the two will neutralize nearly the same amount of acid. The first solution is much more basic, but the two have the same alkalinity.
What species contribute to alkalinity??

[alk] = [HCO₃^-] + 2[CO₃^-2] + [OH^-]

Each of these can neutralize H⁺ ions.

Depending on the pH, these species will make more or less of a contribution to the alkalinity.
For instance, if water is at pH 7, the OH^- is very low as we can calculate from
K_w = [H⁺] [ OH^-] = 1 x 10^-14and
also K_a1and K_a2expressions must be used to determine the distribution of species.

At a higher pH, such as 10, the [OH^-] becomes significant, and the [HCO₃^-] is not.

Note that at high pH, water has less dissolved inorganic carbon than water at pH 7 with the same alkalinity. Dissolved carbon is used to produce biomass
CO₂ + H₂O+ hn ® {CH₂O} + O₂ and HCO₃^- + H₂O+ hn ® {CH₂O} + OH^- + O₂

So alkalinity is a measure of fertility
Note that production of biomass produces OH^- also, making solution more basic.

This, in turn, allows more CO₂ to dissolve.

Solubility of CO₂ in Water

CO₂ is much more soluble in water at higher alkalinity than in pure water
In pure water its solubility is 1.146 x 10^-5

In water containing OH^- what happens to the equilibrium ?
CO₂ + H₂O Û H₂CO₃H₂CO₃ Û H⁺ + HCO₃^-HCO₃^{- Û}H⁺ + CO₃^-As long as substantial amounts of OH^- are present, these equilibria shift because of the reaction H⁺ + OH^- ® H₂O
Metals in Water

Bare ions do not exist in water--hydrated with a shell of water molecules surrounding them.

May also coordinate with other species present

May undergo oxidation or reduction, precipitation, or acid-base reactions.

Fe ( H₂O )₆ ³⁺Û Fe(OH) ( H₂O )₅ ²⁺ + H⁺ (acid-base)

Fe ( H₂O )₆ ^{3+ Û}Fe (OH )₃ _(S) + 3 H₂O + 3 H⁺ (precipitation)
Fe ( H₂O )₆ ^{2+ Û}Fe ( OH )₃ _(S) + 3 H₂O + 3 H⁺ + e^- (redox)
All attempting to reach the most stable, least energetic electronic configuration.
Importance of SPECIATION of metals in water

Hydrated metal ions
Hydroxy species
Metal complexes -- organic or inorganic ligands (reversibly attached to one or more molecules)
Organometallics (bonded to carbon in a compound)

The biological and chemical activity, solubility, mobility etc. of these species is very different.

Calcium in Water

Comes from solution of various minerals, is a principal cause of water hardness

Temporary hardness caused by calcium bicarbonate can be removed by boiling as:
Ca⁺²+ 2 HCO₃^- Û CaCO₃(s) + CO₂ + H₂O

Dissolution and precipitation of Calcium minerals is closely tied to the amount of dissolved CO₂

CaCO₃(s) + CO₂ + H₂O Û Ca⁺² + 2 HCO₃^-
The high level of Ca⁺² in groundwater is due to CO₂ produced in biomass degradation by bacteria--not from atmospheric CO₂

In natural water the concentration of HCO₃^-and Ca⁺² are about 1 x 10 ^-3 M, which is just about what one would expect from water in equilibrium with limestone (CaCO₃) and atmospheric CO₂ (Can calculate from Solubility constant for CaCO₃ (calcite) and the K’s for the CO₂ dissociations.
Results of the calculation:

[CO₂] = 1.146 x 10^-5M	[CO₃^-2] = 8.96 x 10^-6	[H⁺] = 5.17 x 10^-9
[HCO₃^-] = 9.98 x 10^-4	[Ca⁺²] = 4.99 x 10^-4	pH = 8.29

Alkalinity is determined by titration with standard acid to a phenolphthalein endpoint with the results expressed as mg/liter of CaCO_3.
If we titrate 100 ml of an water sample with 0.001 M HCl and it takes 2.05 ml to reach the endpoint, what is the alkalinity? (phenolphthalein changes color at a pH 8.3)

If we titrated using methyl orange indicator instead (changes at pH 4.3), about how much acid would be required for the titration?

Equilibria involved in groundwater in contact with calcite, sediment and CO₂ from the air

CO₂ (air) Û CO₂ (water)

CO₂ (sediment) Û CO₂ (water)

CaCO_{3
(limestone) Û}Ca²⁺ (aq) + CO₃^- (aq)

CaCO₃(s) + CO₂(aq)+ H₂O Û Ca²⁺ + 2HCO₃^-

All these are governed by equilibrium constants :
K_a1
K_a2
K_sp
And the reaction of calcite with dissolved CO₂ has an eq const of
K = [Ca²⁺] [HCO₃^-]²/[CO₂]
K = Ksp Ka1/Ka2 = 4.2 x 10^-5

Complexation in water
Unidentate ligand, such as CN^- , NH₃
Chelating agents bond in more than one site simultaneously
Some common chelating or complexing agents

Effects of complexation: May enhance or reduce solubility. May reduce the tendency of ions to sorb onto surfaces and therefore increase mobility in groundwater. May cause oxidation or reduction or stabilize certain oxidation states. In biology, complex molecules containing Mg or Fe are important
(chlorophyll and hemoglobin)
Complexation:

Complex is the group with the metal ion at the center surrounded by ligands
May have positive, negative or neutral charge
Coordination number is the number of ligand groups bonded around it.
Metal ions have empty orbitals which are filled by the electron pairs from the ligands.
Some ligands may be rather specific for certain metals, and the pH of the solution may change the selectivity if the ligand is a weak acid, like EDTA

In this case, the ligand is much more available when the pH is high, as it will exist in the ionized form. As pH becomes more acidic the ligand becomes protonated, and H⁺ competed with the metal ion for the sites.
.
Nitrilotriacetic Acid (NTA)
Used in detergents to substitute for phosphate
Ionizes in three steps and can coordinate through the three COO^- groups and through the extra pair of electrons on the N
Ka1 = 2.18 x 10^-2
K_a2 =1.12 x 10^-3
K_a3= 5.25 x 10^-11
What is the effect of NTA and other chelating and complexing agents in the environment?
Pb(OH)₂, Fe(OH)₃, Al₂O₃, FeS, Fe₂S_{3 :}many metal oxides, hydroxides, sulfides are quite insoluble, have very low K_sp
If one removes the product of dissolution: increases the solubility.
Solubilization of metals in solids by NTA

Pb(OH)₂ + HT^2- Û PbT^- + OH^- + H₂O
need the K_eq for this reaction:
Use:
Pb(OH)₂ Û Pb²⁺ + 2 OH^-Ksp = 1.61 x 10^-20
HT^2- Û H⁺ + T^3- Ka₃ = 5.25 x 10^-11
Pb²⁺ + T^3- Û PbT^- Kf = 2.45 x 10¹¹
H₂O Û H⁺ + OH^- Kw = 1 x 10^-14
Add up: to get desired eqn.
Remember when an equation is reversed, K is inverted
When equations are added, K's are multiplied.

Complexation calculation:
Assume that a water sample contains 25 mg/l of Na₃NTA (MW 257). Total concentration of NTA (complexed or not) is:

The system is in contact with solid Pb(OH)^2. at pH 8. In which form does the NA exist predominately? PbT^- or HT^2- ?
Phosphates are also important in water
These hydrolyze depending on the pH
e.g.

H₄P₂O₇ + H₂O ® 2H₃PO₄
So, while the complex phosphate chains are good complexing agents, they mostly don’t survive in water and are much of a problem as far as solubilizing and mobilizing metal ions.
Humic Substances

Decay-resistant parts of vegetable matter

Divided into;

Humin : Non extractable with strong base very inoluble matter exists in water as suspended particles or solid masses ahd has ion exchange properties.
Humic acid: Is extracted by strong base but precipitates when extract is neutralized.
Fulvic Acid: Is extracted in strong base and remains soluble even in acidified extract. These materials exist in solution in water.
Properties: Usually high molecular weight, complex structures, contain carboxyl and phenolic hydroxyl groups which can chelate with metal ions, esp strongly bind Al and Fe ions.
Fulvic acids can keep important metal ions in solution, esp. Fe³⁺
Complexation and corrosion
Corrosion of metal is an oxidation process
4Al + 3O₂® 2Al₂O₃
but this requires O₂ to reach the surface, and usually water or H⁺ to catalyze the reaction. Al₂O₃forms a dense film on the surface which inhibits further reaction.
A chelating or complexing agent can shift the equilibrium:
Al₂O₃ ® 2Al^-3 + 3 O^-2
aiding corrosion.
(also cleaning the surface of oxide films, as necessary in plating processes)

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