Week 5 - Chemical Speciation Flashcards
(23 cards)
What is a chemical species in solution?
The actual form an element, molecule or ion is present in a solution, in its most simplest form-solid or dissolved.
This could be as a free ion (e.g., Ca²⁺), a molecule (e.g., CO₂), or part of a complex (e.g., [Cu(NH₃)₄]²⁺). In water, species can be dissolved, bound, complexed, or solid.
What is chemical speciation?
Chemical speciation refers to identifying and quantifying all the different forms (species) an element takes in a system.
For example, carbon can exist in water as H₂CO₃, HCO₃⁻, or CO₃²⁻, depending on pH.
What are dissolved, colloidal, and particulate forms of chemical species?
Dissolved
-Definition: Individual ions or molecules that are fully integrated into the water.
-Size: Smaller than ~0.001 µm (1 nanometer).
-Example: Na⁺, Cl⁻, HCO₃⁻, Ca²⁺ – you can’t see them, and they don’t scatter light.
-Chemistry: These species are “solvated” (explained below) and are freely mobile and reactive.
-Test: They pass through filters with a pore size of 0.45 µm.
Colloidal
Definition: Tiny particles suspended in solution—not true solutions, but not large enough to settle out quickly.
-Size: ~1 nanometer to ~1 micrometer (0.001–1 µm).
-Example: Iron (oxy)hydroxide particles, clays, organic aggregates.
-Chemistry: Chemically reactive due to large surface area. They can adsorb ions and affect transport of metals and nutrients.
-Test: These don’t pass through very fine filters, but they also don’t settle easily—must use ultrafiltration or centrifugation to isolate.
Particulate form:
-Definition: Larger solid matter suspended in water.
-Size: >1 µm (often considered >10 µm operationally).
-Example: Sand grains, bits of decaying plant material, large clumps of minerals.
-Chemistry: These may be less reactive per unit mass but can release or adsorb species as they dissolve or settle.
-Test: Removed by normal filtration (e.g., 0.45 µm filter).
(Distinction b/w colloid and particles is somewhat arbitrary)
How are different forms of species separated or studied?
Using filters (e.g., 0.45 µm is a common cutoff), ultrafiltration, or centrifugation. These are operational distinctions—meaning there’s no hard chemical line between, say, dissolved and colloidal; it’s based on size and method used.
What is ligand binding or complex formation?
This occurs when neutral molecules (ligands) or anions like NH3 bind to metal cations to form complexes such as [Cu(NH₃)₄]²⁺. These complexes can alter metals’ solubility, mobility, and reactivity in solution.
What is sorption (adsorption, absorption, ion exchange)?
Adsorption: Binding of ions/molecules to a surface (e.g., Pb²⁺ adsorbing onto clay).
Absorption: Uptake into the internal structure of a solid.
Ion exchange: Swapping of ions between a solid and the solution (e.g., Ca²⁺ replacing Na⁺ on a resin in a water softener).
What is chemical activity? How is it measured?
Chemical activity reflects how “effective” a species is in participating in reactions. It’s lower than concentration in real (non-ideal) solutions due to interactions between ions.
Formula:
Activity (a) = activity coefficient (γ) × molality (m)
What is an activity coefficient (y)?
It adjusts concentration to reflect non-ideal behavior. In dilute solutions, γ ≈ 1. In salty water (high ionic strength), γ < 1, meaning species are less chemically “available.”
Calculated using Davies equation.
What is ionic strength (I) and how does it affect activity?
Ionic strength (I) measures how many charged particles are in solution.
In concentrated solutions, spacing between dissolved ions is low enough to allow nearby anions and cations to get electrostatically attracted to eachother
These attractions diminish the ability of ions to participate in chemical reactions (e.g mineral growth, ion exchange)
Because their effective concentration is lower than actual concentration in solution, its mulitplied by activity coefficient
Formula:
I= 1/2 ( ∑mi x zi^2)
where
m = molal concentration of species
z = charge on species (squared)
High ionic strength reduces activity by increasing electrostatic shielding between ions.
Once you calculate ionic strength (I) using formula, you plug that into Davies euqation to get activity coefficient (y).
You can then calculate ion activity using:
ai = γi x mi
What is the Davies equation?
The Davies equation is used to calculate the activity coefficient (γ) of an ion in solutions with moderate ionic strength (up to ~0.5 M).
It is an extension of Debye-Hückel theory and accounts for non-ideal behavior in more concentrated solutions like freshwater and brackish water. This theory states that ionic strength controls activity.
Formula:
logyi = -Azi^2 [(√I / 1 + √I) - 0.3I]
where:
-A is a temperature-dependent constant (≈ 0.5085 at 25°C in water)
zi = charge of the ion
I = ionic strength
What factors reduce the chemical reactivity of individual ions in aqueous solution?
- Electrostatic shielding:
Other ions in solution surround and partially neutralise charged ions, decreasing their effective interactions. - Ion pair formation (aqueous complexes):
Oppositely charged ions are close enough in solution (e.g., Ca²⁺ and SO₄²⁻) to form a transient, weakly bound complex (e.g., CaSO₄⁰). These are not as chemically active as free ions.
Extent of ion pairing heavily dependent on temperature and ionic strength (salinity).
Ca^2+ + SO4^2− ⇌ CaSO4^0
This reduces the concentration of “free” Ca²⁺ and SO₄²⁻, lowering their reactivity.
Why it matters:
Ion pairing alters speciation and must be accounted for in modelling natural waters, especially in high ionic strength environments.
What is the Saturation Index (SI)?
SI = log(IAP/K). It shows how close a system is to equilibrium:
SI = 0: at equilibrium
SI > 0: supersaturated (mineral might precipitate)
SI < 0: undersaturated (mineral might dissolve)
What is the Ion Activity Product (IAP)?
IAP is the product of the activities of dissolved ions. It’s compared to the solubility constant (K) to assess whether a mineral will precipitate or dissolve.
What is the simplest mechanism by which inorganic ions dissolve?
Through congruent dissolution of a mineral — the solid dissolves completely into its constituent ions without forming new phases.
What is the effect of water temperature on mineral solubility?
In most cases, increasing temperature increases solubility due to greater kinetic energy of water molecules.
Exception: If the dissolution is exothermic (releases heat), solubility may decrease with increasing temperature.
Adding heat causes the system to shift in the reverse direction (precipitation) to oppose the added heat — per Le Châtelier’s Principle.
Why does gas solubility decrease in warm water?
Gas dissolution is exothermic, so rising temperature reduces gas solubility. This is described by Henry’s Law:
C=kH x P
Where:
C = concentration of gas in solution
kH = Henry’s constant (temperature dependent)
P = partial pressure of the gas
thus inverse relationship between solubility of gas and temp described by temperature dependence of the constant in Henry’s law
What is the role of residence time in mineral dissolution?
Longer residence time allows more dissolution, though the rate decreases as the solution approaches saturation (parabolic dissolution curve).
It’s a graphical description of how dissolution rate decreases over time as a system approaches saturation.
So, residence time controls how far the system progresses toward equilibrium with the mineral — and the parabolic curve reflects the diminishing rate of that progress.
How does the ionic potential (z/r) influence solubility?
Ionic potential = charge/radius.
Higher z/r (high charge, low radius) (e.g., Al³⁺, C⁴⁺): lower solubility due to strong electrostatic interactions and tendency to form solids.
Lower z/r (e.g., Na⁺, K⁺): more soluble
However, in some cases:
Formation of polyatomic anions can increase solubility by complexing with high-charge ions (e.g., forming Al(OH)₄⁻), reducing their free-ion activity.
Complexing, meaning binding with ligands to create complex ions
How does pH affect solubility?
pH controls speciation of ions and minerals.
Acidic pH:
Often increases solubility of metals by protonating anions:
Protonation is the process by which a hydrogen ion (H⁺) binds to a negatively charged species, such as an anion.
In aqueous systems at low pH, where H⁺ concentrations are high, anions like carbonate (CO₃²⁻), phosphate (PO₄³⁻), or sulfide (S²⁻) can become protonated to form neutral or less negatively charged species (e.g., HCO₃⁻, H₂PO₄⁻, or HS⁻).
This reduces their availability to combine with metal cations to form solid precipitates. As a result, metal–anion minerals become less stable and may dissolve more readily, increasing the solubility of metals under acidic conditions.
Reaction with OH⁻ in metal hydroxide equilibria:
At low pH, the concentration of hydrogen ions (H⁺) in solution is high. For metal hydroxides such as Fe(OH)₃, which dissolve according to the equilibrium:
Fe(OH)3 (s)⇌Fe3+(aq) + 3OH(aq),
the added H⁺ reacts with OH⁻ ions to form water.
This removal of OH⁻ from solution reduces the concentration of a product, which causes the dissolution equilibrium to shift to the right according to Le Chatelier’s Principle. As a result, more of the solid Fe(OH)₃ dissolves to restore equilibrium, leading to increased solubility under acidic conditions.
Alkaline pH:
Can increase solubility of amphoteric metals (e.g., Al) by forming soluble hydroxide complexes.
At very high pH, in addition to increasing OH⁻ concentration, the hydroxide ions can act as ligands and form soluble hydroxo-complexes with metal ions such as Fe³⁺. For example:
Fe3+ +4OH− →[Fe(OH)4] −
This complexation reaction removes free Fe³⁺ from solution. According to Le Chatelier’s Principle, the system responds by shifting the Fe(OH)₃ dissolution equilibrium to the right to produce more Fe³⁺. As a result, even though OH⁻ is abundant, more of the solid dissolves, leading to increased solubility due to complex formation.
How does hydrolysis affect mineral solubility and pH?
Hydrolysis = bonding of H⁺ to an anion in a mineral.
It consumes protons, which can decrease solubility by reducing available H+, making solution more basic.
Fe(OH)3(s) ⇌ Fe3 + (aq)+3OH-(aq)
At high pH, there’s already a lot of OH⁻ in solution.
You’re adding more of a product (OH⁻).
According to Le Chatelier’s Principle, this pushes the reaction to the left.
That means precipitation happens, not dissolution.
So solubility goes down.
How does ligand binding affect solubility?
Ligands (e.g., Cl⁻, OH⁻, CO₃²⁻) bind to metal ions to form soluble complexes, which increase solubility by lowering the activity of the free ion.
How does ionic strength influence solubility?
At high ionic strength, strong ion-to-ion interactions can create electrostatic shielding, which reduces the activity of individual ions in solution.
This reduction in ion activity can increase solubility by suppressing the formation of solid phases (precipitates).
However, it can also inhibit solubility by limiting interactions needed for dissolution. e.g. if dissolution requires ions to interact with water or ligands, and they’re too shielded. So can also mean less driving force for dissolution.
The effect depends on the mineral type, speciation, and saturation state of the system.
How does chemical activity influence solubility?
Solubility depends not just on concentration, but on activity (a = γ × m). At high ionic strength, activity is lower, altering equilibrium and apparent solubility.
