Aromaticity - huckel's rule & whether or not something is planar?
969728
Huckel's rule has many exceptions (even flat systems such as pyrene can violate Huckel's rule)  There is also reference to determining aromaticity in fullerenes which are not flat.
306971
Cr has 6 electrons\nEach Benzene contributes 6, so that's \n\n6 + 6(2) = 18. CrPh2 does\n\nFor Co(CO)4-\n\nCo provides 9 (uncharged)\nNegative charge provides 1\nCO (carbonyl) provides 2 each, so that's 8\n\n9 + 1 + 8 = 18 electrons\n\nBoth organometallics obey the 18 electron rule.
304043
Because algebra skills are very lacking in most people. It's just to make sure that they know what the rearranged formula's look like. But you are right, those 2 laws are probably all you need to remember. P=IE is always true, but remember than Ohm's Law is only an approximation for materials that have a linear E vs R graph, such as metals at room temperature. Its not true for semiconductors and gas discharges.
786820
Flourines electron configuration is [He]2s2 2p5.  In the p-orbitals there are three subshells, each subshell is capable of taking only two electrons a piece.  This leaves you with one electron that is unpaired.  Because in the p-orbitals you can only have a total of 6 electrons.  So, 6-5=1.
1090457
a the number of outer electron that will bond with other atoms to fill thei valence shells
44394
1. The Aufbau principle: Orbitals are occupied in order of increasing energy. 1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<4f<5d<6p<7s<5f<6d\n\n2. The Pauli exclusion principle: Only two electrons can ocoupy an orbital and they must have opposite spin.\n\n3. Hund's rule: When there are degenerate orbitals, one electron is added to each with their spins parallel until all are half full. Then a second electron with opposite spin is added to each degenerate orbital. (See below)\n\n(Basically every space/box/circle must have an arrow pointing up before any space/box circle can have one pointing down.)\n(Everyone gets one before anyone gets two.)
644051
See if this article has a diagram\n\nhttp://jchemed.chem.wisc.edu/Journal/Issues/1999/Apr/abs542.html
900522
Unstable is not the best word.  Being a 5 membered ring in organic chemistry is highly desirable so there isn't a problem here.  The compound however is extremely flammable.
706929
T-shaped molecules cannot be from trigonal planar because the only angles possible in a trigonal planar arrangement are 120 degrees, and T-shaped have a 180/90 angle difference.  This is possible for trigonal bipyramidal molecules because you would have two atoms at the axial positions and one equatorial, leaving you with two lonepairs in the other equatorial positions and a 180/90 angle difference between atoms around the central atom.
793636
shawty's NOT good at stoichiometry...has made a fundamental error in the final line so end result is rubbish...\nhow can you end up with MORE Cu than you started with when it's 2 moles : 2moles? \n\n(8/159)x 127 = 6.4 g
852794
dobereiner's like that mendeleev guy right? if my memory is right, all  of the elements above have similar chemical properties, that's why he placed them in his "triads". This means that all of them act similarly. Remember that it is the VALENCE electrons that causes the specific reaction of an element. So, as all three elemnts have the same number of valence electrons, they act similarly.
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.
784377
This is true.  The Px, Py, and Pz orbitals are degenerate in all atoms provided the atom is not subjected to a magnetic field. \n\nIn hydrogen and other single electron systems (He+ etc) only the n quantum number matters to determine energy.  So you can actually say that the 2s, 2px, 2py and 2pz are all degenerate.  \n\nThe s, p etc orbitals lose this degeneracy when there is more than one electron due to electrostatic repulsion between the electrons, so then you need both the n quantum number and the l quantum number to determine energy.
731498
The orbitals themselves have no real "shape".  The graphical representations you are looking at are actually density maps of where an electron in that orbital is most likely to be found.\n\nIf you take a look at an "f" orbital, for example, you'll see that the "shape" is not contiguous.  It looks like a hourglass going through a donut.  Electrons being the happy things they are can appear the various different permitted energy levels around a nucleus and can even "tunnel" outside the atom itself.  Just remember when you're looking at those pictures that you're not really looking at shapes but density graphs of where an electron in that type of orbital can most probably be found.
925332
Are not integers. Example: C-C bonds in benzene are bond order 1.5
743269
Yes.\nVSEPR was developed in 1957, while in 1940 Sidgwick and Powell reviewed the structures of molecules then known. They suggested that for molecules and ions that only contain single bonds, the approximate shape can be predicted from the number of the electron pairs in the outer or valence shell of the central atom. The outer shell contains one or more bond pairs of electrons but it may also contain unshared pairs of electrons (lone pairs). Bond pairs and lone pairs were taken as equivalent, since all electron pairs take up some space, and since all electron pairs repel each other. Repulsion is minimized if the electron pairs are orientated in space as far apart as possible.\n1. If there are two pairs of electrons in the valence shell of the central atom, hte orbital containing them will be oriented at 180<sup>o</sup> to each other. It follows that if these orbitals overlap with orbitals from other atoms to form bonds, then the molecule formed will be linear.\n2. If there are three electron pairs on the central atom, they will be at 120<sup>o</sup>  to each other, giving a plane triangular structure.\n3. For four electron pairs the angle is 109<sup>o</sup> 28', and the shape is tetrahedral.\n4. For five pairs, the shape is a trigonal bipyramid\n5. For six pairs, the angles are 90<sup>o</sup>  and the shape is octahedral.\n\n\nVSEPR was developed extensively by Gillespie and  Nyholm who improved the Sidgwick-Powell theory to predict and explain molecular shapes and bond angles more exactly. Gillespie summarized his theory in 4 points:\n1. The shape of the molecule is determined by repulsions between all of the electron pairs present in the valence shell. (This is the same as the Sidgwick-Powell theory).\n2. A lone pair of electrons takes up more space round the central atom then a bond pair, since the lone pair is attracted to one nucleus whilst the bond pair is shared by two nuclei. It follows that repulsion between two lone pairs is greater than repulsion between a lone pair and a bond pair, which in turn is greater than the repulsion between two bond pairs. Thus the presence of lon pairs on the central atom causes slight distortion of the bond angles from the ideal shape. If the angle between a lone pair, the central atom and a bond pair is increased, it follows that the actual bond angles between the atom must be decreased.\n3. The magnitude of repulsions between bonding pairs of electrons depends on the electronegativity difference between the central atom and the other atoms.\n4. Double bonds cause more repulsion than single bonds, and triple bonds cause more repulsion than a double bond.
115249
Try this URL... it shows some diagrams.  Although it does not have the full set of configurations for d and f orbitals.