CHAPTER 9 – BONDING AND MOLECULAR STRUCTURE:
FUNDAMENTAL CONCEPTS
_______________ refers to the way atoms are arranged in space and
______________ describes the forces that hold adjacent atoms together.
Valence electrons
The electrons in an atom can be divided into the ___________ electrons and the
___________ electrons.
The valence electrons determine the ___________ ___________ of the atom
because chemical reactions result in the loss, gain or rearrangement of these
electrons.
For main group elements, the valance electrons are the s and p electrons in the
outermost shell. For the main group elements, the number of valence electrons is
equal to the group number in A/B system.
Valence electrons for the transition elements include the electrons in the _____
and _________ orbitals.
Lewis structure for atoms
The element’s symbol represents the atomic nucleus together with the core
electrons. Up to four valence electrons, represented by dots, are placed one at a
time around the element’s symbol; then, if any valence electrons remain, they are
placed to ones already there. Lewis structure for carbon is
Draw Lewis structure for
Li
Ne
A maximum of 8 valence electrons can be arranged around an atom and is referred
to as _________ of electrons. An atom having an octet of electrons is very stable.
All noble gases, except He, have eight valence electrons.
Atoms (main group) can achieve the stable octet configuration in one of the
following ways
1. An atom can lose a valence electron and form a __________ :
2. An atom can gain a valence electron and form an __________ :
Ionic bond is formed when one or more valence electrons are transferred from one
atom to another creating negative and positive ion.
3. Atom can __________ valence electrons with another atom and form covalent
bonds.
In most cases, the bonds are somewhere in between ionic and covalent bonding.
Net atomic charge = # of protons - # of electrons
For elements in groups 1, 2 and 3 the ionic charge is normally the ________
__________
For elements in groups 15, 16 and 17 the ionic charge is normally the
__________________ - ____
Ionic Compounds
When anion and cation are brought together, a force of attraction occurs between
them. This force is called _____________ force and is given by Coulomb’s law
where k is the Coulomb constant (8.988 × 109 N · m2 · C-2), n is the charge of the
ion (in units of e-), e is the charge of an electron (1.6022 × 10-19 C), and d is the
nucleus-to-nucleus distance.
The energy released when a cation and anion form an ionic bond can be calculated
using a related formula:
This formula calculates the bond energy for one gaseous molecule.
Calculate the bond energy for a sodium-chloride bond (bond length = 279 pm).
An ionic solid consists of millions upon millions of ions arranged in an extended
3D network called a ___________ ___________.
Since each ion is surrounded by oppositely
charged neighbours, it is held tightly in its
location. At room temperature each ion can move
just a bit around its average position.
Considerable energy is required to move an ion
fast enough and far enough to escape the
attraction of its neighbour. Hence ionic solids
have high __________ __________ indicating
that ionic bonds are strong.
Give the molecular formula and the charges associated with each ion
1. Sodium fluoride
2. Magnesium chloride
3. Calcium oxide
4. Aluminum oxide
Which has higher melting point MgO or NaCl and why?
Most ionic compounds are hard
solids. The force of attraction
between the oppositely charged
neighbours makes the solid rigid.
A blow with a hammer causes the
lattice to cleave cleanly along a
sharp boundary. Why is it so?
The energy associated with the formation of ion pair can be calculated from the
following equation.
__________ ____________ is the energy of formation of one mole of a solid
crystalline ionic compound when ions in the gas phase combine. The equation for
lattice energy is similar to coulomb’s law, except the constant is different (For
NaCl, k = 3.525 × 10-19 J • nm, CsCl, k = 3.82 × 10-19 J • nm)
What is the lattice energy of NaCl (radius of Na+ = 0.095 nm and Cl- = 0.181 nm)?
What is the lattice energy of LiBr?
The ionization energy (IE) of lithium is 520 kJ/mol
Energies of vaporization of lithium and Br2(l) are 134.7 kJ/mol and 15.46 kJ/mol resp.
Energy of formation of gaseous bromine (Br2
2 Br) is 111.7 kJ/mol
Electron affinity of bromine is –324 kJ/mol
Formation energy of LiBr (Li(s) + 1/2 Br2(l)
LiBr(s)) is –351.2 kJ/mol.
Covalent compounds
In molecules or ions made up of entirely _____________ atoms, the atoms are
attached by covalent bonds.
Energy is _______________ when covalent bond forms
∆E = _______ kJ/mol
The _______________ energy of reaction means that the product (H2) is more
stable than the reactants (2 × H).
By convention, chemists actually write the reverse reaction equation and list the
bond dissociation energy (i.e. how much energy is required to break the bond and
move the atoms far enough away from each other that they don’t interact):
∆Edissociation = 435 kJ/mol
Consider the structure of F2: F has ____________ valence electrons. The Lewis
structure shows that F has one unpaired electron and 3 electron pairs. In F2, the
unpaired electrons, one on each F, pair up in the covalent bond.
The pair of electrons in the F – F bond is called ________ _________ and the
other 6 pairs are called __________ __________.
In CO2, the carbon atom shares two pairs of electrons with each oxygen and so is
linked to each O atom by a ___________ bond. The formation of a ____________
bond is necessary for the carbon and both oxygen atoms in CO2 to achieve the
octet configuration.
Systematic approach to constructing Lewis structure of molecules and ions
Let’s draw the structure of formaldehyde, CH2O
1. Identify the central atom. It is normally the one with least electron affinity (C,
N, P and S often appear as central atoms). The rest are terminal atoms.
2. Determine the total number of valence electrons
3. Place one pair of electron between each pair of bonded atoms to form a single
bond.
4. Use any remaining pairs as lone pairs around each terminal atom (except H)
such that each terminal atom is surrounded by eight electrons
5. Place any excess electrons around the central atom.
6. If the central atom has fewer than eight electrons, move one or more of the lone
pairs on the terminal atoms into position intermediate between the center and the
terminal atoms to form multiple bonds.
7. Calculate the formal charge for each atom and indicate any which is not zero.
Draw the Lewis dot structure of SO2.
Draw the Lewis dot structure of NO3-
In NO3-, which atom has the charge?
formal charge = # valence e- in neutral atom – [# lone pair e- + ½ # bond e-]
¾ N has ___ lone pair electrons + ___ bond electrons.
Neutral N has ___ valence electrons.
Therefore, formal charge of N = ___ - [___ + ____] = ____
¾ O(a) has ___ nonbonding electrons + ___ bond(s).
Neutral O has ___ valence electrons.
Therefore, formal charge of O = ___ - [___ + ____] = ___
¾ O(b) has ___ nonbonding electrons + ___ bond(s).
Neutral O has ___ valence electrons.
Therefore, formal charge of O = ___ - [___ + ____] = ___
Formal charge (Qf) is the charge on an atom assuming that every bond is
completely covalent.
As a general rule, we want to keep the formal charge on each atom as close to 0 as
possible (without giving any atom more than a complete octet).
Draw the best Lewis dot structure for BF3.
Exceptions to the octet rule
Compounds with an atom having fewer than 8 electrons
Boron has ____________ valence electrons and so is expected to form
__________ covalent bonds with other nonmetals. This results in a valence shell
for boron with only _______ electrons. Hence it can accept _______ pair of
electron from another atom. Molecules with lone pair can fulfill this role. When
both the bonding pair electrons originate from one of the bonded atoms, the bond
is called a _______________ ______________ bond.
Compounds with atom having more than 8 valence electrons
Elements in the third and higher periods often form compounds and ions in which
the central atom has more than four valence electron pairs. (e.g. SF6, PF5)