Covalent bond 2 Flashcards
VSEPR Theory use
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.
electron domains
both bonding pairs and lone pairs of electrons around the central atom
electron (domain) geometry
arrangement of all the electron domain regions
around the central atom.
molecular geometry
arrangement of only bonding regions
around the central atom
→ The lone pairs on the central atom are only shown in diagrams.
VSEPR THEORY
● 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
shape of molecules
2 electron domains
bonding regions: 2
lone pairs: 0
shape and bond angle: linear, 180
shape of molecules
3 electron domains
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
shape of molecules
4 electron domains
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
determining shape of a molecule
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.
BCl4-
● 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°
h2o2
● 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°
definition of electronegativity
Electronegativity: the ability of an atom to attract the shared pair of electrons in a covalent bond.
formation of electronegativity
● 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.
polar covalent bonds
● 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.
what determines the polarity of covalent bonds
● 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).
Electronegativity difference and ionic bond
● 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
Do all molecules with polar bonds become polar molecules
● 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.
intermolecular forces
● 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.
types of intermolecular forces
● 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
London dispersion force
● 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.
what determines London dispersion force
● 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.
dipole-dipole attraction
● The electrostatic attraction between polar molecules.
● The more polar a molecule, the stronger this attraction.
● Stronger intermolecular force than London dispersion force.
Determining the polarity of molecules
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
Hydrogen bonding
● 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
strength of hydrogen bonding
● Hydrogen bonds are the strongest form of intermolecular force.
● Molecules with hydrogen bonding have boiling points significantly higher than that
would be predicted from their molar mass
