What is the normal distance in a covalent bond?
398514
They all vary, depending on the size of the atoms involved.  Covalent bonds distance is defined as 1/2 the distance between the two atoms radii.  As you can see this is difficult to determine if they are not the same atom.  Most bond lengths are estimated based on similar structures.
957540
You have three types of bonds (essentially)\n\nYou have ionic, non polar covalent, and polar covalent.  Each has a strength in bonding.\n\nIonic is the strongest.  These will have the highest melting points (just try to melt salt for instance).  The non polar covalent bonds tend to be liquids or gases at room temperature (very low melting points).  Examples of these would be Methane , Benzene, ethane etc.\n\nPolar covalent bonds have a stronger pull on the electrons and therefore tend to be higher in melting point than non polar and lower than ionic.
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.
838159
Sigma corresponds with "s" orbitals.  Pi corresponds with "p" orbitals.  We know that 2 electrons fill up the s orbitals.  We know that it takes 6 electrons to fill up the p orbitals.  Carbon has four bonding sites, so it could theoretically have four bonds.  That is 1 sigma bond and 3 pi bonds.  Bottom line is to remember that it corresponds with orbitals.  Hope this helps!
97857
Yeah, it's the greatest electronegativity difference between the two atoms that are forming that bond.  In other words, it's the difference between how far away the two elements are from fluorine, the most electronegative element.  Diatomic molecules like Cl2 are very nonpolar because both Cl atoms are equidistant from fluorine, and HBr is more polar because Br is closer to fluorine than hydrogen.
12669
A covalent bond\nThe other is an Ionic Bond which is when one atom gives an electron to another atom.
1027607
Covalent bonds are where two (or more) atoms "share" outer layer electrons in order to achieve stability. Polar molecules are where the electronegativity of the atoms do not cancel each other out, and hence the molecule forms a charge.\n\nA carbon and a hydrogen forms a covalent bond. The carbon has 12 protons while the hydrogen has only one. Hence, the carbon will draw the shared electron from hydrogen towards itself more than that of the hydrogen drawing the electron shared by the carbon, and hence the carbon becames negatively charged because the electron from hydrogen spends more of its time around carbon while the hydrogen, left with its single proton, becomes positively charged.\n\nAn example of a polar covalent bond is water. \n- O---H +\n    /\n    H\n   +
13313
"Normal" bonding occurs when a ligand donates electrons to a metal. When a metal ion donates electrons back to the ligand, this is called back-bonding. The combination of normal bonding and back-bonding creates a strong bond between the ligand and the metal. The reason that phosphines, carbon monoxide, and cyanide are so poisonous is because they bond readily to iron in biological systems and cannot be displaced by the ligands (such as oxygen) which should bond to iron in normal metabolic processes.
952034
polar refers to bonds and molecules where the electrons tend to be closer to one atom than another because of the "strength" with which the atoms "pull" the electrons.  This "strength" is usually measured by a number called "electronegativity".  There should be a chart or list of these numbers in your text book.\n\n"Covalent" is a type of chemical bond where the electrons are shared between atoms, rather than transferred (ionic bond).  Usually covalent bonds form between two non-metals.\n\nFrom your question it looks like the first two terms (polar and nonpolar) are referring to individual chemical bonds.  Find the two atoms attached to either end of the bond, look up their electronegativities and see if one atom "pulls" on the electrons stronger than the other.  This would produce a polar BOND.\n\nThe last two (nonpolar covalent; polar covalent) could refer to both a bond or a molecule.  For the molecule you need polar bonds,  and then you have to know the shape of the molecule since it is possible for the individual bond "pulls" to cancel each other - kind of like a tug of war.  Generally if a molecule is symmetrical (triangle, tetrahedron or linear triatomic) the "pulls" cancel.  Although the BONDS may be polar covalent, the entire MOLECULE would be nonpolar covalent.
541274
Bond order of CO is 3, bond order of CO+ is 2.5.  CO+ is paramagnetic.\n\nCO is (sigma)^2 (pi)^4, CO+ is (sigma)^2 (pi)^3
837515
You need to look at the central atom and determine the number of lone pairs and bonding pairs of electrons (multiple bonds in this case will not count, as the orbitals used in those bonds will be orthogonal to the bonding axis and thus won't be hybridized).  2 is sp, 3 is sp2, 4 is sp3, 5 is sp3d, 6 is sp3d2.
214929
www.visionlearning.com/library/module_viewer.php?mid=55 - 66k \n\nwww.biology.arizona.edu/biochemistry/ tutorials/chemistry/page2.html - 7k - \n\nhttp://en.wikipedia.org/wiki/Chemical_bond
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
840267
The word molecule tells you that it is a covalent compound (if it were ionic, it would be called a unit cell). Covalent compounds are USUALLY between nonmetals but in some cases can be between a metal and a non-metal. It all boils down to electronegativity difference (how much harder one atom is pulling on the electrons in question than the other atom). I hope this helps!
963224
The space between the molecules
754819
no too sure since its been a while since iv done this, but its a simple enough molecule, C2Br4  2Br to each carbon and a double bond between the carbons, its a linear moleule, not sure about the electronegativity of the elements as my notes r all packed away but the Br atoms should sit slightly of veritcal from each other toward the outside of the molecule. and im sure that the two halves sit at 90 degrees to each other, not sure if u call it this but if u do then 2 Br attached to one carbon will sit along the x- axis of the molecule and the two Br attached to the other carbon will sit along the Z-axis, ask ur tutor or teacher if this is correct as im not 100% sure\n\ntetrahedral only applies to CBr4, thats because the 4 Br are comin off the same carbon, because there are 2 carbons it changes it all together\n   (ignore the full stops, needed them to position the atoms decently)\n\nBr&#92; ........ /Br\n.......C=C\nBr/ ........ &#92;Br\n\nThis is not perfect but imagine the top Br on the left aiming toward you at 45 degrees and the top one on the right aiming away at 45 degrees, hope this helps
809957
The most polar bond will be between the atoms with the greatest difference in electronegativity.\n\nF2 has bond that is not polar since the electrons are equally shared between the two identical atoms. The electronegativity series is As<S<Br<F so the biggest difference is between As and F and the most polar bond is in AsF3.\n\nAs belongs to the same family as N. Thus in AsF3 it forms 3 single bonds and also has 1 lone pair. So the electron-group arrangement is tetrahedral but since there are only 3 identical bonds there will be a net dipole moment.\n\nI already explained why F2 has no dipole moment\n\nBrF3 has trigonal bipyramid electron group arrangment (3 bonds+2 lone pairs) but not corners are occupied by identical atoms (2 have lone pairs of Br) so the dipole moments do not cancel out and the molecule has a net dipole moment\n\nSF4 has again trigonal bipyramid electron group arrangement(4 bonds + 1 lone pair) with one lone pair at one of the corners so it has a net dipole moment\n\nSF6 has octahedral electron group arrangement (6 single bonds)\nand since all corners have F atoms the dipole moments of the bonds cancel each other and the molecule has no net dipole moment.