importance of secondry coordination sphere around metal centre in the active site?
476658
secondary coordination sphere plays a vital role in two ways; (1) it may give extra stability to the ligand through resonance after conjugation with the metal site through resonance , thus increasing the rate of conjugation with the metal. (2) it may also decrease the rate of conjugation with metal by causing steric inhibition.
1013102
Temporary electrostatic complexes where each ion retains their coordinated waters, but they are still close enough to effectively form a complex between the metal and the ligand.  \nUsually form beween Hard cations and anions \nAnother term is outer sphere complex.  These complexes are difficult to analyze directly, and are normally estimated from coulombic attraction and Bejerrum's ion association model; \nThe log Kstab of ion pair complexes range from ~0 to 4, depending on charge density \nThere can be ion pair or coordinative complexes for the same pair of ions, with an equilibrium established between them.  The major ion pairs in water are: \nCaHCO3+, ; CaSO4°; CaOH+; MgCO3°; MgSO4°\n\nThe degree of interaction is only slightly greater than the electrostatic interactions accounted for by the DeBye-Huckel theory, but this is enough to change activities. \nIon pairs become more important in saline waters and brines, where the interactions must be specifically accounted for using the Pitzer Equations. \n\n\n\nget to this link and you find the whole :)
644051
See if this article has a diagram\n\nhttp://jchemed.chem.wisc.edu/Journal/Issues/1999/Apr/abs542.html
900859
If the metal not in the compound is higher on the activity series than the metal in the compound then a reaction will occur and the metals will switch spots. If the metal not in the compound isn't higher on the activity series than the metal in the compound then no reaction will take place. You got to be joking if you think I'm going to be doing 20 problems for you. The only way to learn these is to do them on your own. Just so you know in #6 a reaction will not occur.\n\nHere is an activity series:\nhttp://en.wikipedia.org/wiki/Reactivity_series\n\nREACH doesn't know how to do these, don't listen to his answer.
744262
2, since it's in group 2 of the periodic table
363082
2Ag+ + CO_33-2 -> Ag_2CO_3
961493
You need to know more than just the electron configuration.  There is contraction from left to right on the periodic table, plus additional periodic trends that must be considered.   \n\nEven more importantly, you need to define whether you are looking for the largest element IN ITS NATURAL  STATE (i.e. Cl2, Hg metal, etc) or AS THE ION.   Atomic and ionic radii differ.\n\nMore clarity is needed here.
752247
Remember that the halide, which is going to eventually serve as the leaving group, is pulling electrons away from the carbon center, inducing a dipole where most of the electron density is around the halide (leaving group) and less of it around the carbon center.  The incoming nucleophile is thus even MORE drawn toward that carbon center when there is more induction of the dipole.  Therefore, the more electronegative the leaving group, the stronger the induced dipole.  The stronger the induced dipole, the more reactive is the carbon center towards SN mechanisms.
711569
If the ligands lie in the x-y plane, then the d-orbitals on the central atom that have all their electron density *out* of this plane will have the lowest energies.  These are the d-xz and d-yz orbitals. \n\nThe d-z^2 orbital has its density mostly out of the x-y plane, but there is a ring of electron density in the x-y plane, so this orbital will have a somewhat higher energy than the d-xz and d-yz orbitals.\n\nThe d-xy and d-(x^2-y^2) orbitals have their electron density concentrated in the x-y plane so their energies will be the highest.  They have the same energy.\n\nSummary, highest energy at the top:\n\nd-xy, x-(x^2-y^2)\nd-z^2\nd-xz, d-yz
1100165
Mg is neseccary for the catalytic reaction. A lot of enzymes have an ion in their catalyitic center which usually co-ordinate the reacting groups.
753266
Al ligands have available electrons which can be donated to the central metal ion.  These electrons are mostly as lone pairs in the ligand.  They enter the d orbitals to form dative bonds.
865719
its not that the co bonds with the heamoglobin it forms a complex i.e it replaces the nitrogen atoms in heamoglobin to form the complex...CARBOXY HEAMOGLOBIN.. so we dont have the case of ...C reacting..
953082
Okay, it's been a looooooong time since I had to use my chemistry skills, so lets see what I have:\n\nCopper is more magnetic with electrons than silver is because it's electrons (with fewer of them) are close to the nucleus (or whatever you call it) of the element.  Silver is more apt to lose its electrons with them being so far away.  \n\nCopper is more often found in a +2 state.  So I'd say that's a safe bet to go with.  I'd only assume a +1 state if your teacher said so.\n\nSo therefore:  2AgNO3 + Cu -> Cu(NO3)2 + 2 Ag would probably be your best bet.\n\nBut once again, it's been a long time since I used this so I don't know how right I might be.
188677
Figure 21: Heteropolymolybdate structures. (A) The structure of [Mo< ...  \n  \nchemical compound\n\nCrystal field, ligand field, and molecular orbital theories \n\nConsiderable success in understanding certain coordination compounds also has been achieved by treating them as examples of simple ionic or electrostatic bonding. The German theoretical physicist Walther Kossel's ionic model of 1916 was revitalized and developed by the American physicists Hans Bethe and John H. Van Vleck into the crystal field theory (CFT) of coordination, used by physicists as early as the 1930s but not generally accepted by chemists until the 1950s. This view attributes the bonding in coordination compounds to electrostatic forces between the positively charged metal ions and negatively charged ligands--or, in the case of neutral ligands (e.g., water and ammonia), to charge separations (dipoles) that appear within the molecules. Although this approach meets with considerable success for complexes of metal ions with small electronegative ligands, such as fluoride or chloride ions or water molecules, it breaks down for ligands of low polarity (charge separation), such as carbon monoxide. It also requires modification to explain why the spectral (light-absorption) and magnetic properties of coordinated metal ions generally differ from those of the free ions and why, for a given metal ion, these properties depend on the nature of the ligands. Thus, the complex ion hexafluoroferrate(III) (FeF63-) has magnetic properties to be expected from a substance with five unpaired electrons, as does the free iron(III) ion (Fe3+), whereas the magnetic properties of the closely related hexacyanoferrate(III) ([Fe(CN)6]3-) correspond to only one unpaired electron.\n\nSince 1950 it has been apparent that a more complete theory, which incorporates contributions from both ionic and covalent bonding, is necessary to give an adequate account of the properties of coordination compounds. Such a theory is the so-called ligand field theory (LFT), which has its origin in the more general, but more complicated, theory of chemical bonding called the molecular orbital (MO) theory. (Molecular orbitals describe the spatial distributions of electrons in molecules, just as atomic orbitals describe the distributions in atoms.) This theory accounts with remarkable success for most properties of coordination compounds.\n\nAn important conclusion from LFT is that two types of bonds, called sigma () bonds and pi () bonds, occur in coordination compounds just as they do in ordinary covalent (organic) compounds. The more usual of the two are  bonds, and they are symmetrical about the axis of the bond;  bonds, which are less common, are unsymmetrical with regard to the bond axis. In coordination compounds,  bonding may result from donation of electrons from ligands, such as fluorine or oxygen atoms, to empty d orbitals of the metal atoms (the designation d merely being a way of indicating the particular orbitals involved). An example of this type of bonding is the chromate ion, (CrO4)2-, in which the oxygen atoms donate electrons to the central chromium ion (Cr6+). Alternatively, electrons from d orbitals of the metal atom may be donated to empty orbitals of the ligand. This is the case in the compound tetracarbonylnickel, Ni(CO)4, in which empty  orbitals in the carbon monoxide molecules accept d-orbital electrons from the nickel atom.
1008643
I actually asked an inorganic chemist the very same question, and this is what he told me:\nI seperated out the section that I thought was the best reason he gave, and I think it's important to read the last paragraph.  His other reasoning seems good, too.  I'm not trying to say it doesn't make sense.  I just think the part I seperate is probably what they had in mind when designating the groups.\n\n When you say they are put into the same "group" I assume you mean the group B numbering within the transition metals columns, which would make the Fe, Co and Ni triads all part of group VIIIB.  Now with the main group, also termed group A, the significance of the roman numeral to signify the standard oxidation state of the element is obvious.  The oxidation states get more complex in the transition metal series and frequently multiple oxidation states are observed.  The Cu triad (Cu, Ag, Au) is termed IB due to stable +1 states that are common especially silver which is only found as elemental or +1, where copper may alos be found as a +2 and Au may be found as +3.  The Zn triad (Zn, Cd, Hg) are exclusively +2 and are thus labled IIB group.  You then skip back to the begining of the transition metals and number across due to stable +3 and +4 for Sc triad and Ti triad respectively.  There are a few reasons why in the classical B grouping Fe, Co, and Ni are lumped into a "group" in the middle of the transition metal series.  First as you go across the transition metals the +2 state is becomeing continuously more stable and prevalent.  For Fe there are several oxidations states that are common, however +2 and +3 are most stable and common.  When you move over to Co the +2 becomes the dominant oxidation state with +3 also commmon and very rare examples of higher oxidation states.  When you move over to Ni it is almost exclusively +2 or +3 with rare exceptions.  The second point that I would consider in grouping them together would be their adherence to the 18 elecron rule.  These elements in the middle of the transition metall series follow the rule almost always where earlier elements frequently violate the rule with lower electron counts or even odd number counts and the later elemnts are frequently content with 16 or even 14 electron counts, the later condition is common for group IIB.  \n\n\nAnother consideration to give this group the lable VIIIB is the d8 configuration that is common (Fe(0), Co(I), and Ni(II)) are all d8 systems and frequently analogous compounds can be made with the different metals.  \n\nI hope this shed some light on your question.  I also should note that the system that groups these elements together is an older group numbering system and the new and current grouping system avoids the grouping by groups A and B and simply numbers across the table from 1 to 18, with groups 3-12 being the transition metals.
922867
Check out that link.  It has a picture that shows how it looks like: the grey Al atom is in the center, and 3 green Cl atoms are around it in a triangular arrangement.
574078
This might help :)\nhttp://www.chemicalforums.com/index.php?board=8;action=display;threadid=3263