misc.seament

Introduction:

According to the literature that I surveyed, the electrodeposition of sea-
ment is not a direct oxidation reduction reaction akin to electroplating
but is an indirect process where the precipitation out of solution of the
seament is a secondary by product of the local pH change near the
cathode that occurs when electrolyzing sea water. There may also be
secondary post accretion chemistry involved as well as biological
mineralization by marine organisms which proceed simultaneously
with the electroaccretion. To understand the processes there are a few
chemical principles and reactant names which should be defined prior
to the discussion in the sense that they are relevant to seament
electroaccretion.
 

____________________________________________

pH:

This is the measure of the "acidity" of a liquid. Solutions with low pH
numbers like the electrolyte in car batteries are termed acids. A
solution with a high pH number like a mixture of sodium hydroxide
Drano crystals and tap water is termed a base.
Numerically:
The hydrogen ion [H+] concentration = 10 raised to the -pH power.
The hydroxide ion [OH-] concentration = 10 raised to the -(14 - pH)
power.
The [H+] ion concentration and the [OH-] ion concentration are related
such that as the [H+] ion concentration goes up the [OH-] ion
concentration goes down and vice versa. Their product is always a
constant.
The pH value of pure water = 7.
The pH value of sea water is about 8. Thus sea water is a mild base.

____________________________________________

Solubility product:

This concept states essentially that there is a maximum possible
concentration of a dissolved substance in water.
For example, take magnesium hydroxide Mg(OH)₂ one of the main
constituents of seament.
There is a reaction that can go in either direction where magnesium
hydroxide can dissolve into or precipitate out of solution by the
reaction:
Mg++ + 2OH- <=> Mg(OH)₂
dissolved ions <=> solids
[Mg++] = the concentration of magnesium ions in solution.
[OH-] = the concentration of hydroxide ions in solution.
If [Mg++] times [OH-] exceeds a certain value either because the
concentration of [Mg++] ions increases and/or the concentration of
[OH-] ions increases, then the solid compound Mg(OH)₂ will form
due to the inability of the sea water to keep that high a number of
dissolved ions in solution. This is similar to the common observation
that there is only a limited amount of sugar or salt that will dissolve in a
given amount of water. The solubility product gives the direction that
the reaction will tend to go but does not define the rate. Calcium
carbonate for instance will eventually dissolve into solution in pure
water but to "dissolve" a sea shell into solution in pure water may take
years!
 
 

Note: The precipitation and stability of CaCO₃ and Mg(OH)₂ in sea
water is complicated by interactions with other ions and compounds.
The concentration product of [Ca++] and [CO3--] in near surface sea
water is actually several times greater than that needed in a pure solution
to precipitatate solid CaCO₃. Precipitation inhibiting ions and organic
compounds are thought to be the reason why mass precipitation of CaCO₃
does not spontaneously occur. However once solid CaCO₃ does form it is
inhibited from redissolving by this supersaturated state. Mg(OH)₂ however
does not exist in a supersaturated state in sea water and is therefore not
stable against redissolving.

____________________________________________

Electrolysis:

In a general sense this is the opposite of electrochemical spontaneous
discharge i.e. battery operation. That is, electrolysis is what happens
when you force current into electrodes immersed into a solution to
produce a chemical reaction that will not naturally occur. In the
context of seament production it is the immersion of two conductors
into sea water and applying a voltage between them high enough that
the reaction that evolves hydrogen gas from the cathode is initiated.

____________________________________________

Cathode:

Functionally this is the electrode hooked up to the more negative
terminal of your power supply. In terms of electrochemistry it is the
terminal that supplies electrons to ions in solution to force a chemical
reaction to occur. In terms of mineral accretion it is the electrode on
which the seament will form.

Caution!! Within electrochemical literature the terms anode and
cathode have different polarities depending upon whether one is
defining a power source or a plating cell. The definitions above are the
ones used by Wolf Hilbertz in his articles and patents and by Marshall
Savage in TMP.

____________________________________________

Anode:

Functionally this is the electrode that is hooked up to the more positive
terminal in your power supply. In terms of electrochemistry it is the
terminal where electrons are taken from ions in solution in order to
facilitate a chemical reaction. In terms of mineral accretion it is the
electrode that is spaced away from the seament brick that you are
trying to electroaccrete and that will corrode away before your eyes if
your current density is too high.

____________________________________________

Voltage:

Measured in volts and also referred to in electrochemistry literature as
EMF for "electro-motive force" that is a measure of the amount of
reaction enabling force that you have put into a system. The greater the
voltage difference in the vicinity of an immersed electrode, the more
reactions are possible, many of which you don't want to occur such as
the oxidation or corrosion of your anode. The voltage difference in
combination with the electrode material determines what reactions are
possible.

The assumption that this reaction potential is equal to the voltage across
the electrodes generally holds true for laboratory test cell conditions (
near zero current and uniform reagent concentrations). A complication
arises when attempting to predict under non-laboratory real world
conditions what oxidation-reduction reactions are possible at a given
electrical operating point.

For example, in a high current electrolysis cell with a resistive
electrolytic media between the electrodes, it is no longer necessarily
true that the voltage applied across the electrodes is equal to the
electrochemical potential seen by the ions local to the electrodes. By
ohm's law (voltage = current x resistance) a very large amount of this
externally applied voltage must appear across the sea water and sea-
ment. Each local electrochemical half cell sees only its local electric
field gradient and local ion concentration. It has no way of knowing
what is on the other end of its circuit. If a significant portion of the
external applied voltage is dropped across the resistances within the
electrolyte, this internal drop must be subtracted from the total voltage
applied by the power source to obtain the effective reaction potential
seen by the anode and cathode electrodes.

These voltage drops are by no means negligible. At the current
densities used by Wolf Hilbertz in his structural member accretion tests
of from 10 to 150 Amps/meter², the corresponding voltage drop
across the sea water alone (conductivity 4 mho/meter), not counting the
seament is from 2.5 to 37 volts per meter of separation between the
anode and cathode.
 
 

____________________________________________

Current:

Measured in amps this is what defines how many charge carriers such
as electrons or charged ions are participating in a given reaction. If a
certain voltage makes a given electrochemical reaction possible the
amount of current through the circuit will determine how much of the
final reactant such as hydrogen gas in an electrolysis cell is produced.
 

The current though a given object is usually determined by the voltage
across it. For a resistor like object such as a volume of sea water below
electrolytic breakdown or a block of seament, the current though it is
linearly proportional to the voltage across it.
 

For the electrodes immersed in sea water the dependence on voltage of
the current is much more complex. The current increases linearly in a
constant slope resistor like fashion with increases in the voltage until
the voltage reaches the value at which it can energize a chemical
reaction involving the ions in solution. At this point the reaction taking
place absorbs the current with relatively little rise in the voltage across
the cell. In this constant voltage plateau region (which actually has a
slight positive slope) the amount of current supplied to the electrode is
what is limiting the rate of production of your reaction products.

A common example of the phenomenon of a relatively constant voltage
across an electrochemical cell over a wide range of currents is that of
the voltage vs. current discharge and charging curves of a rechargeable
battery.

At these voltage plateaus, if you power the electroaccretion system with
a constant voltage source, the operating point is defined primarily by
the external resistances in the circuit such as the resistance of the sea
water, deliberately inserted external resistors and the resistances due to
the seament accretions. If you further increase the voltage across the
electrodes there is a point where the diffusion rate of reactant ions to
the electrode rather than the externally supplied current limits the rate
at which the initial reaction (reaction 1) can take place. The voltage
then starts again to rise with little increase in current. This continues
until the voltage rises sufficiently that a second reaction involving the
ions in solution and the electrode become possible. Then another
voltage plateau occurs. At this new plateau Reaction 1 is still
proceeding at its maximum rate limited by ion diffusion towards the
electrode. Now reaction 2 is added to what is happening at the
electrode. This process continues on to other reactions as the current
density and the voltage increases.

If the electrode in the above graph was the electroaccretion anode
reaction 1 might be the generation of chlorine gas, reaction 2 might be
the generation of gaseous oxygen and reaction 3 might be an oxidation
reaction that corrodes the anode.

The corrosion reactions that may occur are determined by the material
that the anode is constructed out of. In fact the reaction overpotentials
for the chlorine and oxygen gas evolution, according to Mendia 1982,
can vary with different anode material in a way that the order in which
oxygen and chlorine gas evolves as the voltage is slowly increased can
reverse.

An increase in the voltage across a pair of electrode operating within
these plateau regions does not necessarily result in a corresponding
increase in the local electrode electrochemical potential. This is
because of the electrical principle than a change in voltage applied to a
group of series connected components will divide proportionally to the
effective resistances of the components at that particular operating
point. The effective resistance of the electrode in a plateau region
where there is little local change in voltage for large changes in current,
is approximately zero by definition. The larger effective resistances in
the circuit such as such as the sea water and the seament will have
most of the applied external voltage increase appear across them. This
would hold true until the current density increases to the point that the
plateau region starts to slope upwards.

The charging of a battery using an external voltage source in series
with a current limiting resistor is a very common place example of this
non symmetric dropping of changes in externally supplied external voltages
when dealing with electrochemical cells and linear resistances in series.
As the charging voltage is increased the voltage across the battery being
charged stays relatively constant while the increase in voltage appears
across the series limiting resistor.

Note that the determinant of the electrode's voltage operating point in
these curves is the electrode current density. This is because the
transitions to different reactions occur when the ions available for
lower voltage reactions become exhausted and no longer use up the
current supplied by the power supply at that lower voltage. The
electrode voltage can then rise with little additional increase in current.
Intuitively, the larger the area of the electrode that is exposed to the
water, the larger the volume of water and hence reactant ions in the
electrode sea water boundary layer immediately available for
electrochemical reactions. If you have twice the anode area for
instance you will require twice the total electrode current to use up the
ions immediately next to the anode at the rate that they can diffuse in
from the surrounding ocean. This means that for a given amount of
current needed at the cathode, you can keep away from the current
density where anode corrosion rises due to secondary reactions by
simply having adequate anode surface area.
 

____________________________________________

Le Chatelier's Principle:

This is the chemical law that states that if you externally change the
products or reactants in a reversible chemical reaction the reaction will
alter its initial equilibrium state to try to undo the change you just
made.

____________________________________________

Some important reactants and products:

From Hilbertz' patent 5,543,034:

" Apart from oxygen and hydrogen, one cubic mile of seawater
contains:

chlorine-89 500 000 t
sodium-49 500 000 t
magnesium-6 125 000 t
sulfur-1 880 000 t
calcium-1 790 000 t
potassium-1 609 000 t
bromine-306 000 t
carbon-132 000 t

and 51 other minerals and elements."
 

The specific roles of some of the more important ions and compounds:
 

H₂O
aka water. This is where the action takes place and this is what gets
electrolyzed to gaseous hydrogen and oxygen. This is the source of the
H+ (hydrogen ions) that get electrolyzed to form hydrogen gas. The
water is also the source of the OH- (hydroxide) ion that when combined
with dissolved magnesium ions forms the mineral brucite Mg(OH)₂
that makes up the bulk of seament accretions at high current densities.

Cl-
Chloride ions. The negative ion partner to NaCl aka sodium chloride,
table salt. This is the positive-negative ion pair that makes up most of
the dissolved solids in sea water. The sodium (Na+) ion does not take
part in the electrodeposition process. The Chloride ion gets converted
to gaseous Cl₂ at the anode during the electrolysis process.

Ca++
Calcium ion. This is the positive ion in (calcium carbonate)
that makes up the mineral argonite. This is the hard
high strength mineral that we want to form on our cathode.
It is insoluble in near surface ocean water due to the supersaturated
concentrations of its component ions.

Mg++
Magnesium ion. This is the positive ion in Mg(OH)₂ (magnesium
hydroxide) which makes up the mineral brucite. This is the softer
easier to dissolve accreted solid that makes up the majority of the sea-
ment that forms at high current densities. We want to minimize the
proportion of this in our final seament structures. Mg(OH)₂ unlike
CaCO₃ is soluble in near surface ocean water.

CO₃--
Carbonate ion. This is the negative ion in CaCO₃, the "good" mineral
in seament that we want a lot of. The ultimate source of the increased
CO₃-- ion concentrations in seament reactions is from dissolved
carbonic acid ( H₂CO₃).
 

____________________________________________
 

The Electrodeposition Process:
 

The seament electrodeposition reaction is initiated by placing two
electrodes in sea water and applying a voltage across the electrodes
sufficient for hydrogen gas to be evolved at the cathode. For now we
will ignore what is happening at the anode and concentrate on the
reactions that are occurring at the cathode where the seament is being
deposited.
 

(1) As the voltage across the electrode rises there will be a point where
the cathode will become electronegative enough to attract the hydrogen
ions in solution in water, donate electrons to them and take them out of
solution by converting them into hydrogen gas which then bubbles up
to the surface of the ocean.

2e- + 2H+ <=> H₂ (gas)

(2) As the hydrogen ions become depleted near the electrode, there is a
chemical reaction involving the carbonic acid in sea water that will try
to reestablish the old equilibrium concentration. Carbonic acid can
disassociate to form bicarbonate ions and hydrogen ions by the
reactions:

H₂CO₃ <=> H+ + HCO₃- <=> 2H+ + CO₃--

As the hydrogen ions near the cathode become depleted this reaction
because of LeChatelier's Principle will move to the right to try to create
more H+ species in solution. This will also increase the concentration
of carbonate ions ( CO₃-- ) in solution. In fact the concentration of
CO₃-- ions can become large enough such that the reaction:

Ca++ + CO₃-- <=> CaCO₃ (solid)

can proceed to the right, precipitating out solid calcium carbonate onto
the cathode when the solubility product of Ca++ and CO₃-- exceeds
that of what can be kept dissolved as ions in solution. This solid
calcium carbonate, also called argonite, is the good stuff in seament:
hard, strong and nearly insoluble.

The precipitation from solution of CaCO₃ is the first accretion related
reaction that occurs as the cathode voltage and current density is
ramped up. As the cathode voltage and current density is increased
further another reaction starts to dominate.

(3) This second reaction is what forms the majority of the seament at
high current densities. As the hydrogen ions near the cathode are
turned into hydrogen gas, the region near the cathode becomes depleted
of H+ ions and by the chemical laws governing acid-base equilibrium
this increases the pH of the region near the cathode making the liquid in
the vicinity of the cathode into a base. This forces the reaction:

H₂O <=> H+ + OH-

To proceed to the right to try to replenish the H+ ions. This makes the
local hydroxide ion concentration go up. As the hydroxide ion ( OH- )
concentration increases the reaction:

Mg++ + 2OH- <=> Mg(OH)₂ (solid)

can take place once the solubility product of the magnesium ions and
hydroxide ions exceed what can be held in solution. The solid that
forms, magnesium hydroxide is also called brucite. This is a softer
more soluble material than Calcium carbonate and we would like as
little of this brucite as possible in our final seament accretion.

Seament composition versus accretion current density.
(LeQue, "Corrosion and Protection of Offshore Drilling Rigs," Corrosion, 1950)
 
 

Using the figures given above by Hilbertz for the mineral content of sea
water and the facts that the atomic weight of calcium = 40, and the
atomic weight of magnesium = 24, the ratio of the concentration of
magnesium ions to calcium ions in sea water can be calculated to be
about 13 to 1. This is actually close to the ratio of brucite to argonite
measured by Hilbertz in his electroaccreted structures. This indicates
that at a cathode current density high enough to consume all local ions
making the accretion process diffusion limited, that the molecular ratio
of Mg(OH)₂ to CaCO₃ will be about 13 to 1. Hilbertz measured a 12
to 1, to 15 to 1 ratio in his tests. This also implies that if a high ratio of
CaCO₃ to Mg(OH)₂ is desired in the initially accreted mass that the
accretion rate must be significantly slower than in Hilbertz high rate
accretion runs.

In a standard electroplating reaction positive ions in solution would be
attracted to the negatively charged cathode, come in contact with it
acquiring one or more electrons to neutralize its charge and plate itself
onto the cathode in the process. The seament electroaccretion process
is driven however by increases in the carbonate and hydroxide ion
concentrations near a hydrogen producing cathode causing the
concentration products of [Mg++][OH-] and [Ca++][CO₃--] to exceed
their maximum value in sea water near the water-electrode interface.
Local precipitation of mineral solids then results. One complication in
this seament accretion process as opposed to classical cathode
electroplating is that there is no guarantee that all of the minerals
produced will necessarily adhere to the cathode. The paper by:
Mendia, "Electrochemical Process for Wastewater Treatment," which
utilized electrolysis in a mixture of sewage and sea water for waste
treatment reported that the Mg(OH)₂ produced in their electrolysis cell
was in the form of free particles. Under what conditions the
precipitated minerals adhere to the cathode or drop away to form free
particles will have to be one of the subjects of our future research.

Since the seament accretion process is primarily driven by local ionic
concentration gradients near the cathode it is theoretically sensitive to
the velocity of the water near the cathode. High water velocity at the
surface of the cathode can dilute these gradients reducing the formation
rate of the seament. This in fact was reported, but not quantified, by
Humble in his paper "Cathodic Protection of Steel in Sea Water with
Magnesium Anodes."
 
 

Hartt et al., "Calcareous Deposits on Metal Surfaces in Seawater- A Critical Review,
Corrosion, Vol. 40, No. 11, pp. 609-618, 1984; presented some quantitative data
on this local water velocity effect on the efficiency of seament accretion (above).
 
 
 
 

Hartt et al., "Calcareous Deposits on Metal Surfaces in Seawater - A
Critical Review, Corrosion, Vol. 40, No. 11, pp. 609-618, 1984. The
authors claim that the primary effect of aeration is to produce high local
water flow rates in the vicinity of the accretion cathodes. At equivalent
temperatures, there is a clear increase in the proportion of CaCO₃ in
agitated as compared with stagnant water.
 
 

If the local water velocity and level of turbulence is high enough it can also
directly dislodge freshly deposited seament particles. On the other hand the
high local water velocity near the cathode also increases the diffusion rate of
ions into the region near the cathode's surface where electrodeposition is
occurring. Potentially, this could be used to selectively increase the maximum
rate at which CaCO₃ can form before triggering formation of large amounts
of Mg(OH)₂. Hartt et al. 1984 presented some quantitative data to this effect.
 
 
 
 
 
 

Hartt et al., "Calcareous Deposits on Metal Surfaces in Seawater - A
Critical Review, Corrosion, Vol. 40, No. 11, pp. 609-618, 1984.

Precipitation of CaCO₃ via electroaccretion of sea water is theoretically influenced
by water temperature due to the decrease in solubility of CaCO₃ as water temperature
increases. This is contrary to the behavior of most dissolved solids which increase in
solubility with increases in temperature. CaCO₃ solubility is heavily influenced by the
atmospheric carbon dioxide - dissolved carbonic acid equilibrium. The amount of
dissolved carbon dioxide in sea water decreases as temperature increases.
The carbon dioxide - carbonic acid then reduces the amount of CO₃-- ions that can
be held in solution. At a high enough temperature CaCO₃ in fact will spontaneously
precipitate. The formation of scale in high temperature water pipes is one manifestation
of this increase in the tendency for CaCO₃ to precipitate at elevated temperatures.
Unfortunately the increase in seament electroaccretion efficiency becomes significant
only at temperatures exceeding 40C. Near surface tropical ocean water is about 28C.
 
 
 
 

Alteration of seament composition versus time in ocean water.
(Hilbertz, "Solar-generated Construction Material from Sea Water to Mitigate Global Warming,"
Building Research and Information, Vol. 19, No. 4, 1991)
 
 

There are other reactions that take place in addition to this pH triggered
concentration gradient precipitation process. Hilbertz and other have
reported the spontaneous post electroaccretion conversion of brucite
(Mg(OH)₂) to argonite (CaCO₃) in seament structures immersed in
sea water. This process has been reported to proceed with rate
constants on the order of months to years for significant amounts of
conversion to take place. The exact chemical and/or biological
mechanisms responsible for this conversion is not known or at least not
reported in my limited literature search.

There was much speculation though about the mechanism of this increase in the
measured proportion of CaCO₃ to Mg(OH)₂ after current flow ceased. One
explanation is that the Mg(OH)₂ simply dissolves out of the CaCO₃ Mg(OH)₂
mass. A variation of this is that the dissolution of the Mg(OH)₂ increases the pH
of the local environment triggering the precipitation of CaCO₃. Still another possibility
mentioned in that the mechanism is conversion by biological organisms of
Mg(OH)₂ to CaCO₃ (Hilbertz 1979, Hilbertz 1991).
 
 

An additional mineral accretion process in seament has been reported
by Hilbertz, simultaneous bioaccretion by marine organisms. Hilbertz
reported growth by marine organisms, some of which form calcium
carbonate skeletons as they grow, on the surface of seament structures
while active electroaccretion was in process. He further claimed that
the high carbonate ion concentrations, high pH, and high electron
availability at the surface of electroaccreting structures enhances the
metabolism of crawling and sessile shell forming marine organisms.
This can be both an additional source of solid calcium carbonate
deposits for seament structures as well as a possible mechanism for the
conversion of brucite to argonite to enhance the strength of existing
electroaccreted and electroaccreting structures. Much of these possible
biological processes will need to be confirmed and their effects
quantified by future research before they can be factored into
the engineering design of large scale seament structures.
 
 

____________________________________________
 
 

Eric R. Lee
erlee@stanford.edu