Covalent bond 2

Valence shell electron pair repulsion theory This theory is applied to predict the shape of a molecular species. It states that the bonding atoms in a molecule will achieve a geometry (shape) around the central atom in the valence shell.

both bonding pairs and lone pairs of electrons around the central atom

arrangement of all the electron domain regions around the central atom.

arrangement of only bonding regions around the central atom → The lone pairs on the central atom are only shown in diagrams.

● Electron domains are arranged to achieve: minimum repulsion and maximum separation ● Each multiple bond is treated as just one bond pair/ one bonding region ● The repulsion decreases in the following order: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair

bonding regions: 2 lone pairs: 0 shape and bond angle: linear, 180

bonding regions: 3 lone pairs: 0 shape and bond angle: trigonal planar, 120 bonding regions: 2 lone pairs: 1 shape and bond angle: bent , 117-119

bonding regions: 4 lone pairs: 0 shape and bond angle: tetrahedral,109.5 bonding regions: 3 lone pairs: 1 shape and bond angle: trigonal pyramidal, 107 bonding regions: 2 lone pairs: 2 shape and bond angle: bent, 104.5

1) Draw the Lewis structure. 2) Count the total number of electron domains on the central atom. 3) Determine the electron domain geometry as follows: ● 2 electron domains - linear ● 3 electron domains - triangular planar ● 4 electron domains - tetrahedral 4) Determine the molecular geometry from the number of bonding electron domains. 5) Consider the extra repulsion caused by the lone pairs and adjust the bond angles accordingly.

● 4 electron domains around the central atom B ● With 4 bonding pairs of electrons and no lone pairs ● Electron pairs repel in maximum separation ● tetrahedral shape with bond angle 109.5°

● 4 electron domains around the central atom O with 2 bonding pairs of electrons and 2 lone pairs ● Electron pairs repel in maximum separation ● Lone-pair repulsion is greater than bond-pair repulsion ● H-O-O bond angle = 104.5°

Electronegativity: the ability of an atom to attract the shared pair of electrons in a covalent bond.

● Atoms share their valence electrons to form a covalent bond. ● The sharing of electrons is equal when the atoms of the same element share pair(s) of electrons. ● When atoms of different elements form a covalent bond, the sharing of electrons may be uneven.

● The more electronegative atom will draw the bonding pair of electrons towards itself → it will have a partial negative charge (δ-) ● The less electronegative atom will have a partial positive charge (δ+) ● The pair of partial opposite charges is known as a dipole.

● The greater the electronegativity difference between two elements, the more polar a covalent bond is. ● Dipole moment is used to quantitatively measure the bond polarity or molecular polarity (units: Debye, D).

● If there is a large difference in electronegativity of the two atoms in a molecule, the least electronegative atom’s electron will transfer to the other atom. ● This in turn leads to an ionic bond. equal or less than 0.4 = pure covalent bond Larger than 0.4 to 1.7 = polar covalent Larger than 1.7= ionic bond

● Not every molecule with polar bonds is polar itself. ● Dipoles have directions and they can cancel each other out when they are in opposite directions - no net dipoles. ● As a result, the molecule is non-polar.

● The electrostatic attraction force between molecules or non-metal atoms (noble gases). ● They are much weaker when compared with ionic bond and covalent bond. ● Its existence is independent of the strength of covalent bond holding the bonding atoms together within a molecule.

● Intermolecular forces account for the physical states as well as some physical properties (e.g. melting and boiling points and density) of the molecular substances. ● Categories of intermolecular forces: 1) Van der Waals force a) Dipole-dipole attraction b) London dispersion force 2) Hydrogen bonding

● Also called instantaneous dipole-induced dipole attraction. ● Electrostatic attraction in all molecules (no matter they are polar or non-polar) and non-metal atoms. ● Random movement of electrons causes a temporary dipole. ● The electron distribution of the nearby atom or molecule will be disturbed, causing an induced dipole.

● Strength of London dispersion force depends on the number of electrons (or the molecular mass) of the molecule or the atom. ● Increasing number of electrons, more probable to form an instantaneous dipole with a larger size, stronger London dispersion force.

● The electrostatic attraction between polar molecules. ● The more polar a molecule, the stronger this attraction. ● Stronger intermolecular force than London dispersion force.

For molecules with same chemical bonds - ● Asymmetrical shapes with lone pair(s): polar molecule ● Symmetrical shapes (no lone pairs): non-polar molecule For molecules with different chemical bonds - ○ Having at least one polar bond (ΔEN > 0.4): polar molecule ○ Having no polar bonds: non-polar molecule

● A special type of dipole-dipole attraction that occurs when there is a hydrogen atom covalently bonded to a highly electronegative atom X (X can be N, O, or F) in the molecules. ● In this situation, the hydrogen atom becomes highly electropositive, which strongly attracts to the lone pair of electrons on atom X in the other molecule