explain the effect of lone pairs on the shape of molecules such as H2S and NH3?
351716
In NH3, there are 3 bond pairs and 1 lone pair; there is repulsion between the lone pair and bond pair and between 2 bond pairs. Hence all the 4 pairs arrange themselves as far apart as possible to minimize the repulsion between them, while still maintaining the stability of the molecule. Hence, NH3 has a trigonal pyramidal shape which achieves the above purpose.\nSimilarly in H2S, there are 2 bond pairs and 2 lone pairs. Here there is repulsion between bond pair and bond pair; lone pair and lone pair and lone pair and bond pair, so the pairs spread out. In this case, a 'bent-V-shape' or a 'non-linear' shape for H2S is suitable as it minimizes repulsion and maximizes stability.
336395
I'm not pretty sure about this but because of the low molecular weight, the attraction between molecules is little and therefore outside forces can easily scatter these molecules troughout space. This results in gas state rather than liquid or solid
851509
There are two possibilities. Follow this link for a picture of the first possibility. http://img175.imageshack.us/img175/8223/ncssh2.jpg\n\nBasically, all 3 of these elements follow the octet rule (they want 8 electrons in their outermost shell), and as you can see, they all have 8. Now, S is carrying the - sign because it's formal charge is -1. Formal charge is calculated by Formal Charge=# of electrons the element has (S brings in 6) - the # of electrons it "owns" within the bond. Sulfur "owns" 7 in this case (3 lone pairs and one in a shared covalent bond--thus (3*2)+1=7.\n\nTherefore Formal Charge of Sulfur (in this case): is -1 because\n\n6-7=-1\n\nI probably made it sound more confusing than it is. Sorry about that.\n\nBut other than that, you know that Carbon should be the central atom here because it's coming in with only 4 electrons in it's valence shell, so it's more needy than Nitrogen with it's 5 and Sulfur with it's 6. So you stick it in the middle and then pretty much trial and error. \n\nThe other possibility is here.\nhttp://img175.imageshack.us/img175/6659/ncs2wf4.jpg\n\nThis time Nitrogen is carrying the -1 formal charge, because it comes into the bond with 5 electrons in it's valence shell, but it "owns" 6 in this case (2 lone pairs and 2 electrons out of the 4 in the N=C double bond..(2*2)+2=6. So...Formal charge of -1 comes from \n\n-1=5-6\n\nEven though these two molecules look different, they are in fact the same. This is a concept called resonance. \n\nI hope this wasn't too confusing.
88971
Force holding atoms in a molecule together as a specific, separate entity (as opposed to, e.g., colloidal aggregates; see bonding).\n\nIn covalent bonds, two atoms share one or more pairs of valence electrons to give each atom the stability found in a noble gas. In single bonds (e.g., H-H in molecular hydrogen), one electron pair is shared; in double bonds (e.g., O=O in molecular oxygen or H2C=CH2 in ethylene), two; in triple bonds (e.g., HCºCH in acetylene), three. In coordinate covalent bonds, additional electron pairs are shared with another atom, usually forming a functional group, such as sulfate (SO4) or phosphate (PO4). The number of bonds and the atoms participating in each (including any additional paired electrons) give molecules their configuration; the slight negative and positive charges at the opposite ends of a covalent bond are the reason most molecules have some polarity (see electrophile; nucleophile). Carbon in organic compounds can have as many as four single bonds, each pointing to one vertex of a tetrahedron; as a result, certain molecules exist in mirror-image forms (see optical activity). Double bonds are rigid, leading to the possibility of geometric isomers (see isomerism). Some types of bonds, such as the amide linkages that join the amino acids in peptides and proteins (peptide bonds), are apparently single but have some double-bond characteristics because of the electronic structure of the participating atoms. The configurations of enzymes and their substrates, determined by their covalent bonds (particularly the peptide bonds) and hydrogen bonds, are crucial to the reactions they participate in, which are fundamental to all life
964289
The reason or this dipole moment is the presence of a lone pair electrons on the sulfur atom.  Sulfur has 6 valence electron.  4 participates in bonding with 4 F atoms and 2 lone pair gives a dipole moment for the molecule.  I am not sure how to calculate the dipole moment but I think you use 4 bonding electrons and 2 non bonding electrns in the formula.\nhope this helps - Stella
673911
If you could shrink 'way down and look at a single water molecule, you\nwould find that it is made of two hydrogen atoms and one oxygen atom.\nThey are arranged in a V-shape, with the oxygen at the point.  The result\nof this shape is that there is more negative charge at one end of the\nmolecule than the other (oxygen likes to have negative charge).  We call \nthis a POLAR molecule.\n\nOther polar molecules like salt and sugar and Kool-aid will dissolve in\nwater because the water molecule will be attracted to their charged ends.\nThe water molecule likes their charges so much that it will actually\nsurround the salt or sugar molecules, so that each one looks like a ball\nof water molecules.  We call this SOLVATION.  Since the salt molecule\nis surrounded now by water, it can't stick to other salt molecules, so\nit just sort of floats away.
509181
H2 because it will be electronically more like Helium. It will have an electron with 1/2 spin up and 1/2 spin down. It will therefore have a happy orbital. You can also say that the bonding orbitals are occupied, and the antibonding orbitals aren't.
78348
The answer has to do with symmetry of the molecule and hydrogen bonding.  That should get you going in the right direction
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.
238505
I'm not sure to what extend should I answer this. But just what I have in mind and I think might be related. \n-The molecules within the ball are closely packed together (they interact with one another) ; and this what gives solid their firm/fixed shape.\n-The energy of the molecules is low compared to energy of molecules of liquid or gas.The molecules basically stay fixed on their position and vibrate with minimum amplitude.\n\nThat's mostly about physics for ball. I'll add more if I remember more. Cheers
753982
You have some errors there:\n\nSO4 with the -2 charge has 4 bonded pairs and no lone pairs (unbonded pairs) on the sulfur (but three on each of the four oxygens, which I assume you are not counting).\n\nH2 is correct with 1,0\n\nH2O2 is correct with 3, 4 (two lone pair on each oxygen)\n\nCH2O has 4 bonded pairs (four between the carbon and the oxygen for a double bond) and 0 non bonded pair.\n\nCO2 has 4, 0, with double bonds between each oxygen and the central carbon\n\nHF, is 1,3\n\nCl2 is 1, 6, if you count the lone pairs on the two chlorine like you are doing with F in HF but not the O in SO4 with the -2 charge.\n\nSO2 has a resonance structure where three bonded pair are smeared over two bonds between the S and O.  There are six unbonded pair (3 per oxygen) on the two oxygen.\n\nCH4O, (methyl alcohol?, CH3OH?) has four bonded pair between the carbon and three hydrogen and the one oxygen.  The oxygen has another bonded pair for its hydrogen and two unbonded pair.  This is 5, 2 as you have it.
1124895
The NH3 molecule has an unshared pair of electrons and the H+ ion has lost the one electron in its outermost shell. This enables NH3 to readily donate its unshared pair of electrons and H+ to readily accept that pair. That way, they form a dative covalent bond, in which both H+ and NH3 act simultaneously on each other.  The NH4+ ion is formed, in which the charge is distributed over the whole structure.  It is not localised on any atom. The overall charge is +1 originating from the H+ ion.
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.
878681
Molecules exist in different shapes. Many of the physical & chemical prpoerties of molecules arise due to the different shapes.\n\nSome of the common geometrical shapes are trigonal, planar, tetra hedral etc.The VSEPR theory provides a simple treatment for predicting the shapes of polyatomic molecules. The theory was originally proposed by Sigdwick and Powell in 1940. \n\nIn case of molecules with central atom having one or more lone pairs VSEPR treatment is as follows: In these type of molecules, both lone and bonded pairs of electrons are present. The lone pairs are localised on the central atom, & bonded pairs are shared between the two atoms. Consequently, the lone pair of electrons occupy more space as compared to the bonding pair. This causes great repulsion between lone pair of electrons.\n\nThese repulsion effectscause deviations in the predicted bond angles in molecules. \n\nThe descending order of repulsion interaction is lp - lp > lp - bp > bp - bp\n\nlp - lone pair\nbp - bonded pair
758383
Since the central atom, C, does not have any lone pair electrons on it, and it is bonded to 3 oxygen atoms, the general shape around the central atom is trigonal planar.  The hydrogen atom is bonded to one of the oxygen atoms.\n\nTrigonal planar indictates that all 4 atoms, C and three O, lie in a plane and that the O-C-O bond angle is 120o for all bonds.
883178
A) Valence Electrons\n\nI believe b, c and d show the results of the grouping not the reason they're grouped that way to start out with...\n \nb... No... The atom size increases instead of decreases\nc... No...\nc... No...  \nversa...
826844
H\n |    \nC=O:  (You need 2 lone pairs on Oxygen)\n |   \nH\n\nIt doesnt get much better than that :)