COVALENT BONDING AND POLARITY

- POLARITY

- INTERMOLECULAR BONDING

nonpolar covalent bonding

-the electron pair is shared equally, ex. H2(g)

-"tug of war" with equal partners

polar covalent bonding

-when an electron pair is not shared

equally, there is a localized negative charge around one atom, represented by the symbol δ-, while the other atom is more positively charged, δ+

- a polar covalent bond has a slightly negative end and a slightly positive end, ex. H2O, HCl (g)

- polar molecules exhibit some ionic character!

- whether a bond is polar covalent depends on the difference between the electronegativities of the bonded atoms

Electronegativity- the measure of an atom's ability to attract the pair of

electrons it shares with another atom within a covalent bond

- increases up and right on periodic table

-metals have a lower electronegativity then nonmetals

IncreasingElectronegativity

Polar and Nonpolar Moleculespolar molecules

- molecules that have a positively charged end and a negatively charged end

- the slight difference in charge within a covalent molecule is called a DIPOLE

nonpolar molecules

-do not have charged ends, ex. H2

Polarity of a molecule depends on:1. the presence of polar covalent bonds

2. the three-dimensional shape (geometry) of the molecule

ex. ammonia, NH3(g): polar molecule because it contains polar covalent bonds and a pyramidal shape

ex. methane, CH4(g) nonpolar polar covalent bonds are all arranged

symmetrically about the central carbon: symmetrical tetrahedral shape

HOW DO WE KNOW WHAT TYPE OF BOND WILL FORM?- a bond difference of 0.4 or less is considered to be pure covalent

[equal sharing of electrons]

- a bond difference between 0.4 to 1.7 is considered to be polar covalent [unequal sharing of electrons]

- a bond difference of 1.7 or greater is considered to be ionic [electrons were transferred]

Increasing Electronegativity Difference

100% Slightly Polar Very 100

%Covalent Polar Covalent Polar IonicBonding

In order to truly know if a substance is ionic or covalent, experimental data is needed to verify that the properties do apply!!!

Polarity is closely related to intermolecular forces!!!

Intermolecular Forces

- weaker than covalent bonds, but can be stronger than ionic bonds

- the temperature at which a liquid boils reflects the kinetic energy needed to overcome the attractive intermolecular forces (likewise, the temperature at which a solid melts)

- intermolecular bonds are broken when a molecular compound melts and boils

Thus, the strength of the intermolecular forces determines the physical properties of the substance

1. Hydrogen bonding

-water is a highly polar molecule:the difference between the electronegativities of the oxygen and hydrogen atoms is 1.4, and the molecule is bent

-since oxygen has the higher electronegativity, the electrons spend more time around oxygen than they do around hydrogen

the oxygen end of a water molecule has a slightly negative charge

the hydrogen end of the molecule is slightly positive

-the arrangement of atoms and bonds forms a bent shape

- the dipoles created by O-H in water are attracted to opposite charges creating hydrogen bonds: the strongest of all intermolecular forces

- water is capable of dissociating ionic compounds

2) van der Waals forces

-weak forces of attraction between molecules, such as the dipole–dipole force and the London dispersion force

a) dipole–dipole force (DDF)

-occurs between polar molecules, such as hydrogen chloride, HCl

-the slightly positive end of one hydrogen chloride molecule is attracted to the slightly negative end of a neighbouring hydrogen chloride molecule

b) London dispersion force (LDF)

- an intermolecular force of attraction that forms between atoms of neighbouring molecules as a result of a temporary imbalance in the position of the atoms’ electrons

-forms between all molecules, polar and nonpolar

- the side of the atoms with more electrons develops a temporary negative charge, and the side with fewer electrons develops a temporary positive charge; if same happens to neighbouring molecule they attract each other

- since electrons move quickly, the dipole lasts for only a fraction of a second