Electrons will repel each other to maximum separation and minimum repulsion. Greatest angles are between lone pairs of electrons, which cause more repulsion due to a larger electron cloud. Lone pairs of electrons count toward an area of electron density but as they repel to a much greater extent they reduce the bond angle by 2.5°
Linear - 2 bond pairs, 0 lone pairs, 180° (e.g.- BeCl2)
Electronegativity is the ability of an atom's nucleus to attract the bonding pair of electrons within a covalent bond. A dipole is a difference in charge between two atoms caused by a shift in electron density in the bond. If the difference becomes large enough the bond is considered as ionic. Polar bonds can be cancelled out in a symmetrical molecule, so the whole molecule itself is not polar. The stronger the attraction between the atoms, the higher the bond enthalpy and the shorter the bond length.
Intermolecular forces (London forces)
London forces cause all atoms and molecules to be attracted to each other.
1) Electrons in charge clouds are always moving very fast. At any particular moment, electrons in an atom are likely to be more to one side than the other, creating a temporary dipole.
2) This dipole can cause another temporary dipole in the opposite direction on a neighbouring atom. The two dipoles are then attracted to each other.
3) Electrons are constantly moving, the dipoles are being created and destroyed all the time.
Larger molecules have larger electron clouds, meaning stronger London forces.
Molecules with greater surface areas also have stronger London forces as they have a bigger exposed electron cloud.
Intermolecular forces (dipole-dipole)
Polar molecules have permanent dipole-dipole interactions. The charges on polar molecules cause weak electrostatic forces of attraction between molecules.
These permanent dipole-dipole interactions are stronger than instantaneous dipole-induced dipole interactions.
Electrostatically charged rods placed near a jet of a polar liquid, such as water, will attract the liquid molecules towards the rod
Intermolecular forces (hydrogen bonds)
Hydrogen bonding is the strongest intermolecular forces. Hydrogen bond only takes place when hydrogen is bonded to fluorine, nitrogen or oxygen. This is because these elements are very electronegative, so they draw the bonding electrons away from the hydrogen atom. The bond is so polarised and hydrogen has such a high charge density that hydrogen atoms form weak bonds with lone pairs of electrons on the F, N, O atoms of other molecules. Hydrogen bonds increase boiling and melting points. Ice is less dense than liquid water.
For one substance to dissolve into another, bonds in the substance have to break, bonds in the solvent have to break and new bonds have to form between the substance and solvent.
Ionic substances dissolve in polar substances such as water.Ions are pulled away from the ionic lattice by water molecules that they are attracted to. This is hydration. Some ionic substances do not dissolve as the electrostatic force of attraction between ions is too strong.
Alcohols dissolve in water due to polar -OH bond. Hydrogen bonds form between them. The carbon chain is not attracted to water so the more carbon atoms there are, the less soluble the alcohol will be.
Halogen alkanes contain polar bonds but dipoles are not strong enough to form hydrogen bonds with water. The hydrogen bonds between water molecules is stronger than the bonds that would be formed with halogen alkanes, so halogen alkanes do not dissolve.
Non-polar substances dissolve best in non-polar solvents.Substances such as ethene have London forces between their molecules and form similar London forces with non-polar solvents such as hexane.
Redox is the transfer of electrons. Oxidation is the gain of oxygen and reduction is the loss of oxygen. OIL RIG.
The oxidation state of any element e.g. Fe, H2, O2, Pa, S8 is zero (0).
The oxidation state of oxygen in its compound is -2, except for peroxides like H2O2, in which the oxidation state for O is -1.
The oxidation state for hydrogen is +1 in its compound, except for metal hydrides, such as NaH, LiH etc., in which the oxidation state for H is -1.
The oxidation states for other element are then assigned to make the algebraic sum of the oxidation states equal to the net charge on the molecule or ion.
The following elements usually have the same oxidation states in their compounds.