CHAPTERS 8/9 REVIEW – BONDING
DEFINITION:
A chemical bond is an INTRAmolecular (NOT INTERmolecular) force of attraction between two particles – such as
between two ions (ionic bond) or two nonmetal atoms (covalent bond).
Why do bonds forms?
All bonds form because what results is lower in _______________ and more stable than the individual particles were.
H atom + H atom  H – H + _______________
IONIC BONDS
An ionic bond results when a metal reacts with a nonmetal. An ionic bond comes about when one atom “transfers” an
electron (the metal) to another atom (the nonmetal). Ionic bonds are an electrostatic force of attraction between two
oppositely charged particles. The energy involved with this bond can be calculated using a form of Coulomb’s Law.
Na+(g) + Cl-(g)  NaCl(s) + lattice energy. Lattice energy is the amount of energy ______________when a gaseous
cation and a gaseous anion combine to form an ionic compound according to the below form of Coulomb’s Law.
The more negative (or EXOTHERMIC) the lattice energy – the greater the stability of the ionic compound.
(Q1)(Q2)
Lattice energy = k ------------r
In the above equation, k is a constant, Q1 and Q2 represent the charge on the cation and anion, r represents the distance
between the two centers of the ions bonding.
EXAMPLE: In each pair, which ionic compound has the most exothermic lattice energy? Justify your answer
(a) CaO or Na2O
(b) NaF or NaCl
Electron configurations/sizes of ions:
EXAMPLE: Write the shorthand electron configuration for Zn2+
Isoelectronic – ions containing the same number of _________________________.
For isoelectronic ions, the higher the nuclear charge (Z), the smaller the ion.
EXAMPLE
Order the following ions from largest to smallest: O2-, Na+, Mg2+, FCOVALENT BONDS
A covalent bond results when two nonmetal atoms interact with each other. There is a ___________ of electrons between
the two atoms.
Covalent bonds can be of two types.
#1. NONPOLAR COVALENT or PURE COVALENT.
This type of covalent bond occurs between two identical atoms such as we have with the diatomic molecules.
EXAMPLE: H – H
All of your diatomic molecules exhibit pure covalent bonding.
In each of the above, the two electrons are EQUALLY shared between the two atoms. No one atom has a great ability to
attract a bonding electron to itself over the other.
#2. POLAR COVALENT BOND
This type of covalent bond occurs between atoms of differing electronegativity. In this type of bond, one atom has a greater
ability to attract BONDING electrons to itself than the other. As a result a POLAR bond is formed.
EXAMPLE: H-Cl
The above molecule is said to have a dipole moment - a center of positive charge and a center of negative charge.
When drawing an arrow, point the arrow towards the more ___________________________ atom.
ELECTRONEGATIVITY:
Electronegativity is the ability of an atom to _______________ shared/bonding electrons to it. The most electronegative
atom is __________________. This means that electronegativities ______________as you go left to right across a row of
elements and ______________________as you go down a column of elements. (exceptions: 1st 3 noble gases)
Electronegativities can be used to determine the CHARACTER of a bond. That is whether the bond is MOSTLY ionic –
POLAR covalent or NONPOLAR covalent.
If the electronegativity difference is ZERO or VERY small – the bond is NONPOLAR COVALENT. If the
electronegativity difference is intermediate – the bond is POLAR covalent. If the electronegativity difference is large – the
bond is MOSTLY IONIC.
EXAMPLE: Which of the following would has the most polar bond?
(a) N2
(b) F2
(c) HF
(d) HCl
(e) PCl3
BOND ENERGIES
Bond energy is the energy needed to _________ a bond.
Breaking an IONIC bond requires a LOT of energy. Bond BREAKING is always ________________ and bond
FORMING is always _________________________. Ionic bonds are very stable.
Breaking a COVALENT bond is also endothermic – but it may NOT require as much energy as the breaking of an IONIC
bond requires.
Another VERY important aspect of bond energies is that a single bond is weaker – and longer – than a double bond –
which is weaker – and longer than a triple bond.
BOND ENERGY AND ENTHALPY
Bond energies can be used to find the ∆H for a reaction. HOW?
Using bond energies: ∆H = ∑ BEreactants (bonds broken) - ∑ BEproducts (bonds formed)
There is an example of this in the THERMODYNAMICS review.
LEWIS STRUCTURES
Lewis structures are used to show how the VALENCE electrons are arranged in molecules and ions
Refer to handout given when we covered this chapter in the fall if you do not remember how to draw a Lewis Structure.
Most atoms obey the OCTET rule and surround themselves with 8 electrons. HOWEVER, some do not obey it. Hydrogen
and Beryllium will ALWAYS be “octet deficient” – Boron is also octet deficient. Some atoms have the ability to
EXPAND their octet. Any element in ROW 3 or higher can EXPAND their octet and have more than 8 electrons around
them when they are central atoms. HOW IS THIS ACCOMPLISHED? The “extra” electrons are put into “d” orbitals.
“d” orbitals show up for the first time in the 3 rd row.
EXAMPLES: Draw Lewis Structures for the following molecules/ions:
(a) HCN
(b) BF3
(c) I3-
RESONANCE
Resonance occurs when more than one valid Lewis Structure can be written for a particular molecule.
EXAMPLE: Draw all resonance forms for the nitrate ion
EQUIVALENT resonance is shown in the above. How does resonance come about? Would experimentation show the
above molecule to have 1 double and 2 single bonds? NO!!!!! What would be shown in experimentation is that all bonds
are equivalent to each other and that they are all of the same average length. Bond order – based on the number of electron
pairs being shared in a bond. What is the bond order of the N-O bond in the nirate ion?
Refer back to Ch 8 notes and be familiar with FORMAL CHARGE.
MOLECULAR GEOMETRY
Molecular geometry or structure is the 3 dimensional arrangement of atoms in a molecule. This arrangement is determined
by using VSEPR theory. (Valence Shell Electron Pair Repulsion)
What does this mean? It simply says that electron pairs will position themselves as far from each other as possible so that
the most stable arrangement of atoms (lowest in energy) will result.
IT IS VERY IMPORTANT FOR YOU TO KNOW THE TABLE OF GEOMETRIES THAT YOU WERE GIVEN.
FROM A LEWIS STRUCTURE YOU SHOULD BE ABLE TO ANSWER THE FOLLOWING.
EXAMPLES: For the following molecules: identify the molecular geometry, bond angles, electronic geometry, polarity
(a) ClF3
(b) XeF4
(c) CF4
Remember a molecule is polar if the molecular geometry is asymmetrical or the molecular geometry is symmetrical, but the
atoms bonded to the center atom are different. A molecule is nonpolar if the molecular geometry is symmetrical and the
atoms bonded to the center atom are the same.
HYBRIDIZATION
Hybridization is the mixing of atomic orbitals (such as “s”, “p”, and “d” orbitals) to form new orbitals of equal energy.
WHY DO WE NEED HYBRIDIZATION?
The reason hybridization is needed is that the VALENCE ELECTRON MODEL of Lewis Structures is not adequate to
explain the bonding in all particles – such as in methane. 4 equivalent bonds are needed to correctly describe the structure
of methane and only hybridization allows for 4 equivalent bonds in this molecule.
TYPES OF HYBRIDIZATION:
The types of hybridization are sp3, sp2, sp, dsp3, and d2sp3.
To determine type of hybridization, draw the Lewis Structure count the number of effective electron pairs around an atom:
(remember: 1 lone pair – 1 effective e- pair, 1 bond (doesn’t matter whether is a single, double, or triple) = 1 effective e pair)
2 effective electron pairs = sp hybridization
3
“
= sp2 “
4
“
= sp3 “
5
“
= sp3d “
6 “
= sp3d2 “
(On the AP, this hybridization is often written as dsp 3)
(On the AP, this hybridization is often written as d2sp3)
sp3d and sp3d2 only occurs with atoms with an expanded octet surrounding them.
Be able to determine number of sigma and pi bonds in a molecule/ion:
Single bond =
Double bond =
Triple bond =
BONDING REVIEW HOMEWORK
1. Which of the following bonds is expected to be most polar?
(a) C-Si
(b) C-N
(c) C-O
(d) C-F
(e) C-H
2. Ranking the ions S2-, Ca2+, K+, Cl-, from smallest to largest gives us the order as
(a) S2-, Cl-, K+, Ca2+
(b) Ca2+, K+, Cl-, S2(c) K+, Ca2+, Cl-, S2(d) Cl-, S2-, K+, Ca2+
(e) Ca2+, K+, S2-, Cl3. The total number of lone pairs in PCl3 is
(a) 1
(b) 8
(c) 10
4. Which of the following molecules has a dipole moment?
I. CO2
II. SO3
III. CCl4
(a) I and II only
(b) II and III only
(d) 12
(e) 14
(d) IV only
(e) II, III, and IV
IV. IF5
(c) III only
5. Which of the following ionic compounds has the most exothermic lattice energy?
(a) CsI
(b) LiCl
(c) LiF
(d) CsF
6. What is the molecular geometry of XeF2?
(a) bent
(b) trigonal planar
(c) linear
7. Of the following molecules, which has the largest bond angle?
(a) O3
(b) OF2
(c) HCN
(d) CH2O
8. What is the molecular geometry of ICl4- ? Bond angles?
10. Consider the following molecule: N≡CCH=CHCl
How many sigma and pi bonds are in this molecule?
(d) square planar (e) t-shaped
(e) CCl4
Electronic geometry?
9. What is the hybridization of the central atom in SeF4? In SeF6 ?
Each molecule – polar or nonpolar?
(e) MgO