Sorry for the lack of posts

Once again I have failed to update: I’ve been preoccupied with the end of my summer job, then moving and starting recently, personal issues. I will try to get something up soon, probably the final Hellriders post.

Until I do, Stay Geeky

Published in: on August 24, 2010 at 9:43 pm  Leave a Comment  
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Ion Exchange Study

Alright, here are some of my study notes on Ion Exchange! I didn’t see many books on it since the 1970s, so I don’t know if this is an unpopular area of chem or just that we’ve had our book buying budget crash since then but I thought I’d post my notes in case anyone wants to read along.

Title: Ion-exchange : introduction to theory and practice / [by] R. W. Grimshaw and C. E. Harland. —
By: Harland, C. E.. Grimshaw, Rex W.
Published: London : Chemical Society, 1975.
ISBN: 0851869696

MA(solid) + B(sol) ⇋ MB(Solid) + A(sol) Where A&B share a charge and M is the opposite charge.

The most important features of a good ion exchanger are:

  1. Hydrophilic structure of regular & reproducible form Makes sense, you need to be able to reproduce it to be useful
  2. Controlled & Effective ion exchange Makes sense again, you want it to exchange the ions that you desire and you don’t want to have to wait a long time for it to doso
  3. Rapid rate of exchange Makes sense, you want to just pour water through a pipe or something, not have to loop it through a filter a dozen times
  4. Physical stability in terms of mechanical strength & resistance to attrition Makes it last longer, you don’t want holes forming in your filter
  5. Thermal stability Don’t see why you need this all the time, couldn’t you control the water temp instead?
  6. Consistent particle size, effective surface area compatible with scaling up to a larger plant What if I’m making a Brita water filter or such? Then I don’t need to scale it up

As of the publishing of this book “All ion-exchange resins suffer some breakdown on being subjected to frequent drying and wetting cycles” so they are stored wet.

While increasing the affinity to an ion means the resin will bind to it more easily it makes the resin harder to regenerate once it is used

You don’t have to completely regenerate a resin, and this is in fact impractical, just get it most of the way and it will work.

To remove most dissolved salts from water:

1) Split Stream or Hydrogen-Sodium Blend method:

In 2 streams run water through a resin. In one stream it is in the Na form, the other the H+ form. The water in the first stream will be filled with NaHCO3, whereas in the other it will get HCl. Mix this two streams and NaCl, water, and carbon dioxide is generated, with a net loss of minerals, and the CO2 is easily removed.

2) ‘Starvation’ or ‘dealkalysing’ method: Has on 1975 had mostly replaced the split stream method: Uses carboxylic acid bead resins to remove the minerals: 2RH + Ca(HCO3)2 => R2Ca + 2H2O +2CO2, thus easily removing the minerals.

To remove all the dissolved salts from water:

1) Two-bed systems: Run the water through an acid-resin to replace any cations with H+, then a base-resin to remove any anions and the H+. This can also be done in reverse. Any acids formed in step 1 will be absorbed in step 2.

(April 8th starts here)

2) Mixed-bed systems: One column of mixed strong acid-base resins. This means that as soon as the cation or anion is removed and H+ or OH- is released it will react. This very rapidly removes all the ions from the water. The best of the listed systems for removing ions however the bed is very difficult to regenerate as you must separate the two resins, regenerate each and then remix them.

3) Combination systems: This involves one or more of the above systems used together. For example a weak-base bed to remove most of the ions followed by a harder to regenerate strong base bed to remove the rest.

4) Condensate polishing (Powdex process): A powdered strong acid/base mixture is placed on a candle filter in a thin layer. I don’t understand this one.

Saline Water treatment:

Salt water is contains too much salt to remove on a large scale via traditional methods (as of 1975). Softening it is practical (exchanging harmful ions for less/non-harmful ones) however. As of 1975 there was work being done on this however:

1) Electrodialysis: This uses layered membrians of which half allows only anions to pass and the other half only cations. This creates cells with almost no ions and cells with very large amounts if ions (See diagram, R. W. Grimshaw and C. E. Harland pg 43). The feeds from these cells is kept separate. The low concentration cells can be fed through a mixed-bed filter to remove more of the ions.

2) Ion-exchange methods: Three processes are listed as being suitable due to lost cost of regeneration of the resin.

A) Sirotherm Process: Uses heat to regenerate the resin from a mixed weak acid/base bed. It can take brackish water to a drinkable level, but no further. This uses acids & bases that dissociate much more readily at high temperatures and that have flat (highly buffered) titration curves so that a small change in pH will cause a large change in composition. Passing hot (80C and 20C) streams of water through the bed in turn will give a concentrated stream of water and a dilute stream which can be separated.

B) The Desal process: This uses a ‘new’ resin called Amberlite IRA-68 and 3 coloums. It is very cheap to regenerate and works very well. If I find newer references to it I’ll study it more

C) The sul-bi Sul process: A two-column demineralization that uses HSO4-. Another specific example I may return to.

Ok, I’m skipping ahead a bit as this is going way out into chemical engineering, when I’m pretty sure I’m doing more chemist type work, so I’m looking at Ion Exchange Equilibria now

Ion-exchange capacity: The number of equivalents of exchangeable ions per unit weight or volume of the exchanger.
Normally done with the number of equiv or exchangeable counter0ions per dry kg (mili-equiv /g) of the resin in a chosen ionic form. (Normally H+ for for the cation exchanger & Cl form for the anion exchanger)

You will have to differentiate between how much you can theoretically excess and how much can be practically be accessed. For example taking into account the size of a displacing ion in structures where not all the sites are going to be easily accessible. In weak  acid & base resins this will change with each pH due to buffering.

Ion exchange capacity = (Molarity)(vol of base)*100 / ([Mass of resin][100- {% of resin that is water}])

Simplified: Q= mols of base * 100 / (Mass of resin * (100-%water)

Q is the specific ion-exchange capacity in mol/kg or mmol/g of dry protonated form.

The breakthrough capacity is more useful in columns and is always less then the specific ion-exchange capacity. This is the point at which a measurable number of ions from the solute begin to ‘leak’ through the resin.

Inorganic Chemistry Study

Study Notes for Inorganic Chemistry:

Part 1: All the d block elements can be found in nature except for Tc which is highly-radioactive. Of the Lanthanoids only Pm is not naturally occurring. However of the Actinoids Only Thallium (Th) and Uranium ARE naturally occurring.

The noble metals are named such as they don’t react with H+ under normal conditions.

Lanthanoids must be separated via ion exchange chemistry since they are too similar & hard to react to separate via physical methods.

In general the size of atoms goes increases as you go down the periodic table. The size of an atom decreases as you move from left to right. (1st year chem.) This is because of increasing attracting from the nucleus pulling the electrons closer. This causes the orbitals to contract. Ionization energy increases as a result (though there are many spots the don’t follow this trend such as places with full shells or full half-shells). Electronegativity also increases, with F being the most electronegative atom.
p orbitals tend to be comparable in size to s orbitals, though are a touch smaller. This allows for large amounts of sp mixing. d orbitals are much smaller then s orbitals, but are large enough to participate in bonding. f orbitals are much smaller then even d orbitals, and are generally too small to participate in bonding or to mix with other orbitals.

The Lanthanoids decrease in size from La to Lu, this is known as Lanthanoid Contraction.

Note that the AUFBU principle doesn’t work very well for the d & f block elements. In the d-block elements will get a half filled shell one element early by moving a s-electron into the d-block. In the f-block atoms will hold onto a half filled shell one element longer then normal by placing the new electron into a d-orbital rather then an f-orbital.

In complexes all the electrons are in the d-orbitals.

<Skipping over a ton of stuff he will not ask about)

Early transition 3d metals prefer oxidation state 3, late 3d transition metals prefer 2. 4d & 5d transition metals prefer higher oxidation states. f-block elements mostly will take +3, but no trends.

Due to steric effects the highest oxidation states are achieved with oxygen.

Part 2: Coordination Chemistry

A complex is a chemical entity consisting entirely of metal atoms or a metal center surrounded by ligands.

A ligand is an ion or molecule that binds to a metal.

A coordination complex, also known as a Werner complex is a complex that does not contain carbon-metal bonds

An orgometallic complex is one that does contain metal-carbon bonds

The Coordination sphere is the arragment of ligands around the metal center

The coordination number is the number of ligands that are directly connected to the metal center.

In VB theory the ligand uses its lone pairs to form bonds to the metal center. The metal center provides empty orbitals.

A chelating ligand is one that can make several bonds to one metal center. (From the greek word for claw)

Monodentate: 1 attachment, Bidentate =2, etc.

Larger metals allow higher coordination numbers, as do smaller ligands.

CN2: Always linear (in transition metals), almost exclusively observed in d10 complexs

CN3 is very rare. d10 with strong steric effects will form trigonal planer (most common), d0 will form trig pyramidal w/ strict sterics.

CN4 is very common, usually tetrahedral. D8 metals will sometimes for square planer.

CN5 is also common. Trigonal bypirimidal is the most common, but is very close to square pyramidal, so much so that you will sometimes see both at once.

CN6 is also very common. Octahedral is the most common form of it (and the one we deal with the most in this class), trigonal pyramidal is also possible but quite rare.

CN7 is relatively rare. Can be observed w/ very small ligands or early 3d & 4d transition metals (due to high acidity). Much more common with Lanthanoids however & Actinoids however. Forms 3 strange shapes.

CN8 makes a bunch of odd shapes.

CN9 is relatively rare, try d0, Sc, Y, & f-block elements.

CN10-12 are all f-block elements and we don’t really care much about them.

Summary: CN4:Tetrahedral, CN5 trigbyipyramidal or square pyramidal, CN6: Octahedral.

UNLESS: d8: often square planer, d10 often linear, d0: Strange stuff.


Ionization Isomers are when an anionic ligand from outside the coordination sphere swaps with an anion from the coordination sphere.

Hydration isomers happen when water swaps with a ligand outside the coordination sphere.

Coordination isomers happen when you have a complex ligand with multiple metal centers, and ligands swap between them.

Linkage isomers occure when a ligand has more then one point by which it can attach to the metal center. The isomers occure when you swap between these points. Ex: CN, NO2, SCN.

Diasteriomers: Same as organic. Cis/Trans, Mer/Fac.

Enaintomers: Same as organic but Δ and Λ instead of S/R.

Kf is the formation constant. Formed stepwise, Kf = [complex]/[metal][ligand].

β  is formed by multiplying each Kf together.

Chelate effect: Less entropy loss from chains binding to metal centers then free ligands. Therefore binding chains is favoured.

Macrocycle effect: Same as chelate but comparing cyles & chains.

Part 3:

Crystal Field Theory:

Ligands are modeled as point charges.

Purely electrostatic forces: Therefore ligand:Metal interactions are repulsive.

Therefore orbitals that point at ligands are higher in energy then ones that don’t.

Δo values are the difference between the upper and lower d orbitals in an octahedral feild. No clear trend across the table, higher oxidation states increase it, as does moving down the periodic table.

Crystal Field Stabilization Energy: (0.4* eg electrons – 0.5* tg electrons)Δo.

If CFSE is lower then the pairing energy then ‘high spin’ state generated and electrons moved to the tg orbitals.

If you have the right set up in the tg orbitals you can get Jahn-Teller distortion. Moves xy orbital up, xz, yz down (this doesn’t matter), moves x2 orbital down, x2-y2 up: This can stabilize the complex.

D orbitals of tetrahedral complex’s are opposite of octahedral. (3 up, 2 down). Δt is 4/9Δo due to inefficient splitting.

Part 4: Molecular Orbital Theory

Molecular Orbitals are formed from atomic orbitals through a linear combination of atomic orbitals (LCAO)

The number of molecular orbitals formed is the same as the original number of atomic orbitals.

Mixing s & p orbitals is known as hybridization, increases bonding & antibonding character. Can only mix if they have the same symmetry.

Electronegativity lowers the relative energy of the bond.

Bonds become more ionic as the difference between the orbitals it is made of increases.

This is just a general example, CO disproves.

In octahedral feilds:

  • there is no spd mixing.
  • t2g is nonbonding (unless π donating/accepting is involved)
  • 6 lowest orbitals are mostly on the ligands


  • t2 orbitals interact with ligand orbitals of the same symmetry.
  • e orbitals are nonbonding as there are no ligand orbitals with the same symmetry.
  • d orbital order is reversed.

The amount of splitting at the frountier orbital depends on the different between the electronegativities.

π-bonding in octahedral complexs: Δo is increased by π-donation, decreased by π-acceptation.

MO theory can explain the spectrochemical series, with good π donors having weak splitting, σ donors having strong splitting and π acceptors having very strong splitting.

Metals also contribute to splitting: higher oxidation state splits more. Splitting increases down a group. However no trends within a given oxidation state among transition metals.

d0 prefers cis, d2 prefers trans.

Octahedral complexs prefer 18 electrons, tetrahedral prefer 16 electrons.

Part 6. Metal-Metal Bonding and Transition Metal Clusters

Atoms are considered bonded when they are closer then the atomic radius or the metallic radius.

A polynuclear or polymetallic complex contains more then one metal center. 3 types exist: Chains, Clusters (Closed structures with metal-metal bonds), and Cage Compounds (Closed structures with bridging ligands between the metal centers)

To form these metal-metal complexes the metal atoms must have partially empty subshells (so they can accept electrons), have a low oxidation state (To prevent repulsion between cations), and posses the ability to form strong metal-metal bonds (No DUH).

In metal-metal bonds the σ bond is the strongest, followed by the π bonds, then of course the δ bonds.

Metal-metal double bonds exist, though are surprisingly weak: many non-metal bonds are stronger then them. This is due to small overlap, especially with the δ-orbitals, and heavy inter-ligand repulsion.

A Quintuple bond has been formed. That is just wrong. Moving along….

Now a bunch of examples of 1-4 bonded complexes. Not sure what to learn from that. Wasn’t feeling so hot that day, didn’t pay much attention. *sigh* At the start you just pull off the δ electrons, but after they are gone you have to add antibonding electrons.

d4 metals can make double bonds. That is Re(III), Cr(II), Mo(II), W(II), Ru(IV), Nb(I), Despite the fact it is d4 no known quadruple bonds of Mn(III) exist.

Multiple metal-metal bonds tend to be reaction centers.

Alright, moving on to more useful (& thus testable things): Structures of clusters with metal-metal bonds:

We name ligands that bridge atoms μX, as in μ1. The subscript is the number of metal atoms the ligand bridges.

A closo cluster is completely closed, convex, and single shelled. The atoms in it will form a polyhedron. If the polyheadron only had triangular faces it is known as a deltahedron.

There are 4 general bonding types in closo clusters:

  1. Electron precise clusters have exactly 1 electron pair per polyhedron edge.

For a transition metal the 18 electron rule should be followed. μ1, halogens, hydrogen atoms, and groups such as SiR3 all supply 1 electron. NH3, PR3 and CO each supply 2 electrons. Bridging and such will change all of this.

To get the number of metal-metal bonds simply divide unpaired the electrons by 2. As a formula:
M-M bonds = (18*metal atoms –the total electrons)/2
However this only works for metal-metal single bonds.

  1. Clusters with one 2 electron, 3 center (2e3c) bond per triangular face
    If a deltrahedra with no more then 4 edges/faces meeting at any   one vertex does not have enough electrons it forms 2e3c bonds at        the faces.
    The number of 2e3c bonds is: (18*metal atoms – total electrons)/4
  2. Clusters that follow the Wade Rules
    If you have even less electrons then above then you use Wade’s               rules.
    Wade’s Rules: You need 2n+2 skeleton electrons to form a stable   closo compound. n = # of metal atoms. Does not work for     tetrahedral compounds, they only need 8 electrons instead of 10.
  3. And of course: Special Unique Snowflakes, I mean none of the above.

Polyoxometallates: Oxo complexes have no metal-metal bonds, but use an oxo (O2-) ligand as a bridge. They are a result of acid/base chemistry. In 3d cations the coordination number must be 4 and only 1 oxygen can be shared between any 2 atoms. In 4d & 5d metals the larger sizes open up coordination number 6 and allow multiple oxygen per 2 metal atoms, allowing for highly complex structures.

Part 7: Orometallic chem. of the transition metals

Hapicity is the number of atoms in a single ligand bonded to one metal center. Denoted by superscript ηx. If no superscript then the maximum # of atoms is being used. Note this may be different then the number of bonds.

CO bonds w/ carbon, not oxygen. 1 bond w/o back donation, 2 with. σ donor, π acceptor.

Phosphines are not officially orgometallics but we will treat them as such. σ donor, π acceptor. Electrons transferred from metal in σ* bond.

Hydride directly bonds to the metal. Can serve as bridges.

H2 can bond to metals. The stronger the π backdonation the weaker the H-H bond until you have 2 hydrogen lignands instead of one H2 ligand.

η1 hydrocarbons are all σ no π.

η2 alkenes & alkynes will change to alkanes if there is enough backdonation.

η4 alkenes that are nonconjuated act as 2 independent double bonds.

Benzene on the other hand donates via 1 σ donation and 2 π donations. Dontes electrons back via δ orbitals.

Common orgometallic reactions:

  1. Ligand substitution. Thermally or photochemically activated. Preceeds in steps that each have 16 electrons, problems with solvent competition.
  2. Oxidative addition: Turn XY into MX + MY. Are common for 16 electron species. Ligands usually add cis. Reverse = reductive elimination
  3. 1-1 insertion. Atom one away inserts between ligand & metal.
  4. 1,4 insertion: η2 ligands become η1 ligands w/o oxidation state change on the metal.
  5. Reverse of that, β-hydride elimination.
  6. α-Hydrogen abstraction. Removes hydrogen from C, adds a bond to metal.

Electron counting:

Neutral-Ligand method: all ligands are nutral, you count how many e- then donate to get this way.

Electron count: Metal valance electrons + Donated by ligands + charge of complex

Ligands donate 1 or 2 electrons.

Is fast an easy, but overestimates character of bonds. However it overestimates covalent character, underestimates charge of metal, makes it hard to find oxidation state of the metal.

Donor Pair Method: Ligands donate electrons in pairs. They are either neutral or charged.            Oxidation # of the metal = total charge on complex – charge on ligands.

Number of electrons form metal is group # – oxidation state.

Total electrons is sum of electrons on metal and donated by ligands.

Oxidation numbers easy to get, however overesitmats charge, assumed reactivity might be wrong.

So sorry I don’t have any more gaming posts yet: I’m busy with exams. I might get some more of  The Wyzard’s stuff edited on the weekend. I’ll update this post as I do more study, so I don’t have a stream of ichem study posts. Anyway, off to get food now. Until next time Stay Geeky


Post updated as I head to bed: I’d better get a lot more done tomorrow! Ah well, I think I can be ready.

Updated with another section of notes.

Update: I’m now done this exam, but will leave these notes up for interests sake. Until I get embarrassed by how rushed the notes near the end are, or the fact I never got through all of them.  Anyway, I’m going to see if I can get some RPG posts up latter tonight before Quantum comes and eats me alive.

Published in: on December 8, 2009 at 10:07 pm  Comments (2)  
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