How do lone pairs of electrons actually change shape of the molecule of covalent bonds?
491417
This is best explained with an example.\n\nTake an oxygen atom. Six electrons in the outer shell, so it would to have two more. In a water molecule the two extra electrons are supplied by two hydrogen atoms, which share their electrons with the oxygen. \n\nAt first sight one would expect a shape like this for the molecule:  H-O-H, with the positively charged protons as far away from each other as possible. In the actual water molecule the angle between the two O-H bonds is about 105 degrees. \n\nWhat happens is that the four lone electrons like to be (on average) away from the bonding electrons. To make this energetically possible the water molecule is bend.
525685
You can approach this is one of two ways:\n\n1)  By analogy with H2O (because O and S are both in the same group)\n\n2) From first principles ie octet rule, Lewis diagram, hybridisation, shape. \n\nNow S has 6 electrons in outer shell, while H has 1 each, therefore, S will share 1 e- from each H to complete its octet and each H will share an e- with S so H (which is too small to have 8e-s in outer shell) will have a full outer shell. So S is now covalently bonded to to H atoms, accounting for 4 e-s, with 2 non-bonding pairs. That makes 4 pairs of electrons around S all together, so hybridisation is sp3. Shape is tetrahedral minus two vertices ie bent (just like H2O).\n\nI hope this helps.
80472
I'm trying to work with this info:\n\nOctet Rule - atoms want to have eight electrons in their valence shell (two for H) (most will form octets, some will not)\n\nMultiple Bonds\nSingle bond - a covalent bond in which a single pair of electrons is shared by two atoms\ndouble bond - two pairs of electrons are shared by two atoms\ntriple bond  - three pairs of electrons are shared by two atoms\n-generally form with C, N, O, S\n-------\nOkay, oxygen might attach easily to one of the Cls, leaving no spaces, but I'm not sure if there is a doule bond or something weird like that. Same with C.\nIf I remember correctly, CO2 looks like this:\nO=C=O\n------\nWait, check this out- I found the following (and no further info) somewhere on the net.\n\nE) COCl2\n      (C is the \n       central atom)\n\n----------\nWait. No.\n_______\nCl- C -Cl\n___O___\nI'm going to try that...\nI don't know. It lacks the symmetry that you normally see and it's kinda weird looking. What do you think of what I got?\n__________\nI found a pic: http://www.chem.ox.ac.uk/inorganicchemistry3/C/COCl2.html\nLooks kinda like yours.\n______\nHoly Mother of God - http://www.chem.ox.ac.uk/inorganicchemistry3/C/COCl2.html\nI found the name of the compound- phosgene- and did an image search. *exclamation of glee*\n___\nLol, better be right. Later.
423134
CH4 does not have any non-bonding electron pairs.  It has 4 sp3 hybrid orbitals which are all involved in bonding.  SF2 has 4 electron domains according to it's Lewis structure.  This makes the hybridization sp3.  two of these hybrid orbitals are involved with bonding while the other two contain the non-bonding pairs of electrons.  To find the non-bonding pairs you must draw the correct Lewis structure.  a single atom in the molecule will only have a single hybridization.  S in SF2 can not be sp2 and sp3 hybrid.  it is only sp3.  Therefore the non-bonding electrons exist in the sp3 hybrid orbital.
792627
True - symmetry elements cancel dipoles. Take for example CCl4, a tetrahedral molecule. Each C-Cl bond has a dipole associated with it, the negative part is the Cl (Electronegativity (EN) of C= 2.5, EN Cl =3.5). However the symmetry due to the shape of the molecule cancels out all dipole so NET dipole of molecule is zero.\n\nTrue- ex, SO4 or NO3\n\na) 180°, b) 120°, c) less than 109° (eg 104.5° for H-O-H), d) 109' 28" (or 109.5°). The detailed whys of these are due to hybridization of the atoms involved. linear will have sp hybridization and no lone pairs. trigonal planar will have sp2 with no lone pair. Bent can be many combinations, eg sp3 hybridazation with 2 lone pairs of electrons (eg water), tetrahedral will be sp3.
296542
when a bond is formed, some electrons don't participate in that because of hybrid levels, thats why there are some electrons left known as lone pair, they are never single, they are a pair of opposite spin.
892596
Covalent bonding is a form of chemical bonding characterized by the sharing of one or more pairs of electrons by two atoms in order to produce a mutual attraction; atoms tend to share electrons so as to fill their outer electron shells . Such bonds are always stronger than the intermolecular hydrogen bond and similar in strength or stronger than the ionic bond . Commonly covalent bond implies the sharing of just a single pair of electrons. The sharing of two pairs is called a double bond and three pairs is called a triple bond. Aromatic rings of atoms and other resonant structures are held together by covalent bonds that are intermediate between single and double. The triple bond is relatively rare in nature and two atoms are not observed to bond more than triply. \n\nCovalent bonding most frequently occurs between atoms with similar electronegativities where neither atom can provide sufficient energy to completely remove an electron from the other atom. Covalent bonds are more common between non-metals whereas ionic bonding is more common between two metal atoms or a metal and a non-metal atom.
1146263
Molecular bonding.\n\nThis is a function of the atomic weight (protons in the nuclei) and the distrobution of the electron shell.  Give me three years and I will explain it fully.
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.
724737
Electron cloud of N-F bonds will pull closer to N as you reasoned... However -- this will reduce the overall repulsion between the N-F bonds!  Thus, the repulsion from the lone pair-bond e will (in essence) be greater)and will tend to push N-F bonds closer togehter -- thereby reducing the bond angle for NF3.
697038
This is kinda hard to draw on here...\n......_\nK--Br|\n......_\n\nThe lines around Br represent 2 electrons, and the -- represents a bond...ignore the periods
901447
Note that we're not taking energy differences into account.  In cases where the difference in orbital energy is greater than 15eV you get negligible overlap.  Now onto the symmetry-allowed overlaps.\n\n1s and 2s are both spherical and overlap to form a sigma bond.\n\n1s and 2p(x) form a sigma bond.  You can tell because if you draw the 2p(x) lobe you'll see it overlap with the corresponding 1s.\n\nThe 2p(x) and the 2p(y) do not overlap.  There is no possible bond.  Draw the 2p(x) extending toward the other atom as positive or shaded (remember: positive does not denote charge).  The other atom has its 2p(y) pointing up and down (perpendicular) and there is equal positive and negative overlap, thus canceling both of them out.  No net overlap.\n\n3p(y)-3p(y) form a pi bond.  Once again, draw these out and see it for yourself.  Draw the 3p(y) orbitals as rather fat so you can see how they overlap.  Since they are not cylindrically symmetric they are not a sigma bond.  Pretend you were to rotate one of the atoms about the x-axis... you would change orientation and there'd be less overlap.  For the sigma bonds above, rotation does not affect the overlap.\n\n2p(x)-2p(x) have positive overlap.  Draw them and see why.  They point toward one another.  Moreover, if you rotate one of the 2p(x) orbitals about the x-axis you keep the same structure/orientation since the 2p(x) is along the x-axis.  Thus, you can tell that this gives a sigma bond.
696797
One way is by growing a big, pure crystal of the material. Since all the molecules are the same, and the crystal is pure, then all the molecules must fit together a certain way to form the shape of the 'ideal' crystal. \n\nX-ray crystallography, which the previous answerer referred to, is when they shine a thin beam of x-rays through the crystal, to see how they are diffracted. This shows how the crystal is constructed inside.\n\nIn addition to this, if the molecules are large enough, an electron microscope can sometimes see the layers of molecules on the outside of a crystal, or the inner layers if it is split open. \n\nBy using techniques like this, scientist can build up a mental concept of what shape and size the basic units of the crystal must be, the molecules. Then they can make little stick-and-ball shapes that represent the molecules and build up large structures (representing very tiny crystals.) If their predictions are correct, the large shape they build up should resemble the actual crystals, with the same angles and the same shape planes, etc.\n\n19 NOV 06, 1756 hrs, GMT.
537794
Okay, to begin with; the molecular orbital theory is based on the "octet rule."  This means that, with exceptions of hydrogen and helium of course, an atom is most stable with 8 electrons (the octet) in its outer shell.  \n\nA diatomic molecule is a pair of identical atoms forming a small molecule.  The most common diatomic molecules are in the halogens; Cl, F, I, Br.  These atoms each lack one electron from a full octet, and will covalently bond in pairs to each gain a full octet.  Oxygen gas (O2) will bond in a double pair, sharing four electrons to gain an octet on each atom.  \n\nBeryllium contains 4 electrons in total.  The electronic configuration is 2:2 or if you use the other method is 1s2 2s2.  A full octet would be 2:8 or 2:0 to gain the most stability, and the Be2+ is actually in use in labs (2:0).  \n\nBecause it would need to gain electrons from or donate them to some other element with a greater number or greater need, beryllium can't form a diatomic molecule or a pair with another of itself.  \n\nIt is difficult to explain without pictures, but there are a few shell pictures available at Chemical Elements which will let you compare these elements.
627227
The carbonyl double bond is between atoms of different electronegativity, resulting in an electron deficite at the weaker atom (C) and a surplus at the stronger atom (O) along the bond.\n\nStill, the carbon-carbon double bond also does create a dipol (or rather quadrupol) moment:\n\nThe carbon and all other bonded atoms lie in a plane. The double bond electrons are distributed: 2 to the direct bond between the atoms, 2 to the overlapping P-orbitals. This creates a slight electron deficite in the plane of the bond and a slight electron surplus above and below, which is what makes double bonds more reactive than single bonds.
1075073
well it depends, but generally they don't stay that way for long. They react with other atoms that are charged opposite of whatever they are. This is how chemical bonds are formed. It can either be a covalent bond in which they atoms are "sharing" electrons in order for both to be neutrally charged or they form an Ionic bond.
218524
In Ionic bond, elements looses or gains electron(s) and become ions and the electrostatic potential of these ions bind them\n\nIn covalent, the participating elements are bound by sharing electrons between them.