The ion pair is held together by strong electrostatic attractions. Non-covalent bonds and other weak forces Linus Pauling, Chemical reactivity of molecules- tendency to break and form chemical bonds. Biology of molecules- size and shape of molecules, and the nature of weak interactions with other molecules.
Non-covalent bonds and other weak forces are important in biological structures. In water, these bonds are very weak. Hydrogen bonds-result from electrostatic attraction between an electronegative atom O or N and a hydrogen atom that is bonded covalently to a second electronegative atom.
Caused by slight charge displacements. Hydrophobic attractions-cause non-polar groups such as hydrocarbon chains to associate with each other in an aqueous environment. You can see this more readily using the electrons-in-boxes notation. Only the 2-level electrons are shown. The 1s 2 electrons are too deep inside the atom to be involved in bonding. The only electrons directly available for sharing are the 2p electrons.
Why then isn't methane CH 2? When bonds are formed, energy is released and the system becomes more stable. If carbon forms 4 bonds rather than 2, twice as much energy is released and so the resulting molecule becomes even more stable.
There is only a small energy gap between the 2s and 2p orbitals, and so it pays the carbon to provide a small amount of energy to promote an electron from the 2s to the empty 2p to give 4 unpaired electrons. The extra energy released when the bonds form more than compensates for the initial input. Now that we've got 4 unpaired electrons ready for bonding, another problem arises. In methane all the carbon-hydrogen bonds are identical, but our electrons are in two different kinds of orbitals.
You aren't going to get four identical bonds unless you start from four identical orbitals. The electrons rearrange themselves again in a process called hybridization. This reorganizes the electrons into four identical hybrid orbitals called sp 3 hybrids because they are made from one s orbital and three p orbitals. You should read "sp 3 " as "s p three" - not as "s p cubed". You can picture the nucleus as being at the centre of a tetrahedron a triangularly based pyramid with the orbitals pointing to the corners.
For clarity, the nucleus is drawn far larger than it really is. When a covalent bond is formed, the atomic orbitals the orbitals in the individual atoms merge to produce a new molecular orbital which contains the electron pair which creates the bond. Four molecular orbitals are formed, looking rather like the original sp 3 hybrids, but with a hydrogen nucleus embedded in each lobe.
Note that the inner shell of carbon's electrons are not shown above, only the outer shell of carbon's electrons are involved in the covalent bonding. The molecule can be shown as displayed formula with four carbon — hydrogen single covalent bonds A level note: its called a tetrahedral shape , the H—C—H bond angle is o. SiH 4 will be similar because silicon 2. This displayed formula does indicate the shape of the methane molecule as well as how the four single C-H covalent bonds are arranged, but no relative size of atoms or electronic detail of bond formation by electron sharing.
All the bonds in the above examples are single covalent bonds. Below are three examples 7—9, where there is a double bond in the molecule, in order that the atoms have stable Noble Gas outer electron arrangements around each atom. Carbon and silicon have a valency of 4. More complex examples can be worked out e. In each case link in the atoms so that there are 2 around a H electronically like He , or 8 around the C or O electronically like Ne.
On the left are full 'dot and cross' electronic Lewis diagram for the covalent bonding in the methane molecule.
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