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Covalent Bonding Honors Chemistry – Semester 1
Ionic vs Covalent Bonding Ionic bonding = exchange electrons between metal / non-metal
Covalent bonding = sharing electrons between non-metals
Cl
Cl-
Simple Covalent Molecules Simplest example = diatomic molecules
Repulsion between nuclei and between electron clouds
Attraction between nuclei and electron clouds
Covalent bond forms when forces balance
Shared electrons located within molecular orbitals
Energy & Stability Most individual atoms have low stability
atoms form compounds!
Unbonded atoms = high potential energy
Bonded atoms = low potential energy
Bond Length Bond length = marks point of lowest potential energy
Measured in Ångstrom (Å) or picometer (pm)
1Å = 100 pm = 1 x 10-10 meters
Nuclei vibrate, distance changes constantly
average distance = bond length
Bond Energy = energy required to break bonds in 1 mol of a
chemical compound
Energy released when atoms bond
Example:
energy released when
I mol of H2 bonds
= -436 kJ/mol
436 kJ/mol must be
supplied to break bonds
Bond Energy High bond energy = strong bonds
Strong bonds = short bond lengths
Bond energy predicts reactivity of compound
low bond energy reacts more easily
Review Questions Why is H2 more stable than individual atoms of H?
H2 has lower potential energy than H
In what type of orbital are shared electrons in a bond
located?
Molecular orbital
What happens to the potential energy of two atoms as
they approach each other to form a covalent bond?
Potential energy decreases
What happens to potential energy if the atoms get too
close to each other?
Potential energy increases
Review Questions What do you call the the distance between two atoms in
a covalent bond at which potential energy is lowest? Bond length
Which of these is likely to have covalent bonds? Recall that covalent bonds tend to form between non-metals.
H2
MgO
O2
CO2
H2O
K2SO4
CH4 (methane)
C6H12O6 (sugar)
Yes
No
Yes
Yes
Yes
No
Yes
Yes
Electronegativity &
Covalent Bonding
Helps predict type of bond between elements
Generally:
Metals less electronegative than nonmetals
Smaller atoms have greater electronegativity
Nonpolar vs Polar
Covalent Bonds
Shared electrons closer to atom with higher electronegativity!
Non-polar covalent: electron pair(s)
shared equally between nuclei of
similar electronegativity
Polar covalent: electron pair(s)
shared unequally between nuclei with
different electronegativity
Cl Cl
Cl H
Example: Polar bond HF H, F = nonmetals Covalent bond
Electronegativity values:
F = 4.0; H = 2.1
Most shared electrons with F
Partial negative charge (δ-) in F, partial positive charge in
H (δ+)
Creates dipole
Dipole = a molecule or part of a molecule that
contains both positive and negative regions
Electronegativity, Bond
Polarity & Bond Strength
The greater the difference in electronegativity:
the greater the polarity of the bond
the stronger the bonds
Electronegativity & Bond Type
This scale is only a guide!!
Some compounds may deviate!!
Practice:
Predicting Bond Types Na – F
Electronegativity difference: 3.1 ionic
C - Cl
0.5 polar covalent
Ca – O
2.5 ionic
Al - Cl
1.5 polar covalent