Freshwater <=> Clean Energy
Electrolysis is used to make industrial levels of hydrogen for fuel cells. The most efficient electrolysis of hydrogen requires de-ionized (pure) water to make a KOH (potassium hydroxide) solution. The electrolysis process extracts hydrogen gas ( H2 ) and oxygen gas ( O2 ) from the water, and the KOH is recycled.
Anode oxidation: 2H2O + 4KOH ==> O2 + 4H+ + 4OH- + 4K+ + 4e- ==> O2 + 4H2O + 4K+ + 4e-
( note that acid, H+produced from the oxidation, reacts with hydroxide anion, OH- dissociated from KOH, to reform water)
Cathode reduction: 4H2O + 4K+ + 4e- ==> 2H2 + 4OH- + 4K+ ==>2H2 + 4KOH
( note that hydroxide, OH- produced from the reduction, re-associates with the potassium cation, K+, to recycle KOH )
Hydrogen is an energy carrier that can be used to store renewable (continuously sourced) energies, such as solar or wind, for use when the energy is unavailable. This energy can thus be used continuously--rather than continually--to power equipment on site for other applications, such as water treatment and carbon sequestration discussed below.
algal mats--in the 2.1 billion year old Kona
Formation, dolomitic carbonates near
Marquette,Michigan. Stromatolites are some of the
oldest fossils on Earth. Blue-green algae
(procaryotic cynanobacteria) produced Earth's
first atmospheric oxygen, and the associated
carbonate sediments removed the CO2
from Earth's original atmosphere.
In the strictest (correct) sense, “alkalinity” refers to total alkalinity (ANC, see number #2). In the normal pH range of shallow groundwater, as well as human blood, the predominant component of total alkalinity is bicarbonate. Unfortunately “alkalinity” has thus become synonymous with “bicarbonate,” which often leads to confusion in communication when discussing the conservation of alkalinity (see #3).
In the strictest (correct) sense, “alkaline” refers to a solution that has a positive total alkalinity. In pure water equilibrated to the atmosphere (i.e. precipitation) at standard temperature and pressure (STP), alkaline solutions range from pH 5.6 and higher. Unfortunately “alkaline” has incorrectly become a term referring to any pH greater than 7, where pH 7 is the neutral point of pure water without any dissolved carbon at STP. Unfortunately “alkaline” has also incorrectly become a term referring to the opposite of acidity, i.e. “basicity”. These uses should be avoided. Though it is a continuously increasing function of pH, TOTAL ALKALINITY IS NOT A LINEAR FUNCTION of pH AND DOES NOT START AT pH 7.
Total alkalinity quantifies the amount of acid that can be neutralized in solution, thus total alkalinity is also known as the acid neutralizing capacity (ANC). Total Alkalinity is defined as "the sum of bases in equivalents (or milli-equivalents) titratable with strong acid." Total alkalinity can be expressed as an intensive property of a solution as equivalents per liter (or milli-equivalents per liter, mEq/L).
Total Alkalinity has several components, including excess hydroxide base expressed as (OH- - H+), the direct acid neutralizing capacity; and the buffers, bicarbonate (HCO3-) and carbonate (CO32-), the buffering capacity.
Total Alkalinity [ANC] = [HCO3-] + 2[CO3--] + [OH-] - [H+]
where brackets [ ] are concentrations. In some natural waters, other bases that act as buffers may be present in significant enough quantities that they must be included in the calculation of total alkalinity. These include dissolved silicates, borates, ammonia, organic bases, sulfides, and phosphates.
Inorganic carbon, initially in the form of dissolved CO2, forms the weak acid H2CO3 (carbonic acid), which gives rise by dissociation to its conjugate bases, bicarbonate and carbonate, that act as buffers. Unlike with the direct acid neutralizing capacity, the buffering components neutralize incoming (added) acid or base without changing the pH, although total alkalinity does change.
The video above shows the alkalinity components in water in equilibrium with a constant total carbon (Ct) concentration of 1 x 10-3 moles / liter, the same concentration as H+ at a pH of 3 and of OH- at a pH of 11, the endpoints of the graph. Throughout the pH domain of the graph, equilibrium concentrations of the different species of dissolved carbon [carbonic acid, bicarbonate, and carbonate] as well as H+ and OH- are plotted. The equilibrated level of the water with zero [negligible] alkalinity is where HCO3- = H+, at a pH of 4.7. For comparison, a water in equilibrium with 1981 levels of atmospheric CO2 (315 ppm) has a lower total dissolved carbon (Ct) concentration of approximately 1.2x10-5 moles/liter, with a resulting pH of 5.6 at the point of negligible alkalinity, the pH of pure water (with no other acid) precipitation. Most commonly, as bases weathered from rocks are added to the water in the soil, the pH of the water quickly rises above 7, and alkalinity increases with rising levels of first bicarbonate, and then carbonate, in equilibrium with the rising pH.
The video shows the changing levels of carbon species from an alkalimetric titration, the steady addition of milliliters (mL) of hydroxide base into a liter of water containing this constant amount of inorganic carbon (0.001 moles / liter). Carbonic acid, bicarbonate, and carbonate exchange places as their relative equilibrium levels change with the increasing pH.
Things to note in the video above:
a) Total alkalinity (in red) is a continuously increasing function of pH, BUT NOT A LINEAR FUNCTION of pH.
b) As the carbonate components exchange places, they neutralize base (in an alkalimetric titration ... in an acidimetric titration they neutralize acid), thus "resist the pH change." This is represented by steep ramps on the titration curve (in blue).
c) Once a buffering component is consumed, there is a rapid pH change until the next buffering component takes over. This is represented by gradual flats on the titration curve (in blue).
d) Without carbonate buffers, 1 mL of 1 molar OH- added to a liter of pure water would result in an OH- concentration of 0.001 corresponding to a pH of 11. By the end of the titration, pH 11 is only reached after about 3 mL of 1 molar OH- has been added. Thus the buffering capacity (as HCO3- + CO3--) of the total alkalinity is twice the direct neutralization capacity.
No, total alkalinity [ANC] does NOT change with temperature, pressure, or Pco2. In other words, total alkalinity is conservative. But individual components of the sum may / will change, maintaining the sum, including bicarbonate, which confusingly is often called “alkalinity.”
Total Alkalinity [aka Acid Neutralizing Capacity, ANC], is defined as:
Total Alkalinity [ANC] = [HCO3-] + 2[CO3--] + [OH-] - [H+]
where brackets [ ] are concentrations, and in words as, "the sum of bases, in equivalents (or milliequivalents), titratable with strong acid,"
ANC is a conservative property, which means it does not change with temperature, pressure, or Pco2. However, its individual components can change while maintaining the constant sum.
Total alkalinity can be changed under conditions of fixed pH by changing the amount of total inorganic carbon (abbreviated TIC or Ct) dissolved in solution.
Total alkalinity will only change with a chemical reaction involving an acid or base, or the addition of on acid or base to the solution.
Total alkalinity is conserved, even with boiling, so long as the solution is not excessively evaporated, but not necessarily bicarbonate. As you raise temperature, the equilibrium constants will change according to the Van't Hoff equation, thus the balanced concentrations in solution will change, thus the pH will change, but the ANC (in milliequivalents) will not. It takes a chemical reaction with an acid or base, or an addition of an acid or base, to change ANC (total alkalinity).
Boiling will exsolve and liberate CO2 from solution, which will cause effective Pco2 to drop (if unconfined). This in turn will cause bicarbonate alkalinity, HCO3-, to drop, but as HCO3- drops by an "equivalent" amount, H+ (acid) will drop by the same "equivalent" amount and pH will rise. Thus TOTAL alkalinity (the sum) will be unchanged (conserved).
Yes, but it requires acid neutralization in the process. You can raise levels of total alkalinity even under conditions of fixed pH, including blood pH homeostasis.
Remember, total alkalinity is NOT "basicity," i.e. it is NOT defined as a pH > 7 nor pOH. It is, in words, "the equivalent sum of bases titratable with strong acid, as measured in equivalents," and it is a function of pH and Total Inorganic Carbon [abbreviated TIC or Ct]. It is also known as the acid neutralizing capacity (ANC), of which the carbonate components provide a buffering capacity.
As the graph indicates, a higher alkalinity at a given fixed pH has a higher TIC, but raising TIC alone does NOT increase alkalinity, and it drops pH.
Two components must be changed together to raise total alkalinity at a given fixed pH: 1) TIC must be increased (the body does this by increasing dissolved CO2 with respiration), and 2) pH must be fixed by either simultaneously consuming acid (incorporation of acid into hemoglobin may do this), or adding base OH- (Kangen water has hydroxyide alkalinity OH-). Added base OH- converts increased CO2 (increased TIC) to HCO3- by this reaction: CO2 + OH- => HCO3- which both raises TIC and maintains the pH.
The graph is a partial reconstruction of the graph by geochemist Kenneth Deffeyes (1965), but using pH's of interest to the electro-reduced water (ERW) community, including pH settings of Enagic's SD 501 for making Kangen water. The Deffeyes reference is:
Kenneth S. Deffeyes; Carbonate equilibria: A graphic and algebraic approach. Limnology and Oceanography, 10 (1965), pp. 412–426.
pH values of interest on the graph include:
pH 2.5 Enagic strong acidic water (compliment to strong Kangen water); reported pH of Coca Cola due to presence of phosphoric acid (not due to carbonation alone).
pH 3.5 example acidic pH in range of graph
pH 5.0 acidic pH just below the CO2 equilibrated level
pH 5.7 pH of solution equilibrated to atmospheric CO2 (pH of rain)
pH 6.0 Enagic beauty water
pH 7.0 neutral water
pH 7.5 approximate upper limit of blood (pH = 7.45)
pH 8.1 pH of an equilibrated baking soda (NaHCO3) mixture (not plotted)
pH 8.2 pH of seawater (not plotted)
pH 8.5 Enagic Kangen drinking water 8.5
pH 9.0 Enagic Kangen drinking water 9.0
pH 9.5 Enagic Kangen drinking water 9.5
pH 11.5 Enagic strong Kangen water (compliment to strong acidic water)
Note the total alkalinity of "strong Kangen water" (pH 11.5) is almost entirely from OH- when TIC is 0. The other pH levels go to ~0 when TIC goes to zero and are thus manageable by the human body through respiration, for example pH = 9.5, the highest drinking setting of the SD 501
Note: TIC is proportional to Pco2. The graph assumes an open system where CO2 can be increased by dissolution of CO2 gas.
Note: Most additional total alkalinity is converted to bicarbonate alkalinity, HCO3-. In addition to blood pH homeostasis there is blood bicarbonate (HCO3-) homeostasis, so any boost to blood alkalinity will eventually be moderated to normal blood levels. For arterial blood gas, normal bicarbonate levels range from 22 to 26 mEq/L