CHAPTER 8 - BONDING: GENERAL CONCEPTS
8.1 Types of Chemical Bonds
Chemical bond – forces that cause a ____________ of atoms to behave as a _________.
Why do chemical bonds occur?
Bonds result from the tendency of a system to seeks its _______________ energy state ( _______ stable).
The distance where energy is minimal is called ___________ _________________
Ionic Bonding –
An ionic compound forms when a ___________ reacts with a ________________.
Covalent Bonding –
The 2 types of covalent bonds are ____________ covalent and ________________ covalent.
___________________ covalent bonds involve equal sharing of electrons.
___________________ covalent bonds involve unequal sharing of electrons.
Metallic Bonding – a chemical bond characteristic of metals, in which ____________valence electrons are shared
among atoms in a usually stable crystalline structure.
8.2 Electronegativity
Electronegativity – the ability of an atom in a molecule to _________________ electrons to itself.
Electronegativity values are given in Figure 8.3 on p. 354 in textbook.
Periodic Trend:
Why?
Group Trend:
Why?
The Relationship between Electronegativity and Bond Type
Electronegativity Difference
in the Bonding Atoms
Bond Type
Zero -------------- Intermediate ------------ Large
Nonpolar Covalent ------ Polar Covalent ---------- Ionic
Ionic Character INCREASES --------------------------------
Covalent Character DECREASES ---------------------------
Example
Order the following bonds according to increasing polarity: H-H, O-H, Cl-H, S-H
8.3 Bond Polarity and Dipole Moments
Dipolar - a molecule that has a center of ________________ charge and a center of __________________ charge.
This molecule is said to have a dipole moment.
Example: HCl molecule
The electronegativity of __________________ is greater than that of
__________________. Thus the _______will be partially negative and the
_________________ will be partially positive. The HCl has a dipole
moment.
Does the Cl2 molecule have a dipole moment?
We will discuss the polarity of molecules with more than 1 bond later in the chapter.
8.4 Ions: Electron Configurations and Sizes
Ionic Electron Configurations
Write the electron configuration and determine the charge on each of the following atoms when it forms its most stable
ion (noble gas electron configuration).
(a) Mg
(b) P
(c) Br
(d) Rb
d block example:
Sizes of Ions
Cations are always _________________ than their neutral
atoms.
Anions are always _________________ than their neutral
atoms.
Examples:
Place the following in order of increasing size:
(a) Cu, Cu+, Cu2+
(b) Pd2+, Ni2+. Pt2+
(c) O, O-, O2-
Isoelectronic – ions containing the same number of _________________.
Example: the sulfide ion and the calcium ion
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+, F-
8.5 Formation of Binary Ionic Compounds
Lattice energy – energy _______________ when an ionic solid forms from its gaseous ions.
Therefore, lattice energy has a ________________ sign.
General Equation:
To calculate lattice energy, a modified form of Coulomb’s law is used:
k = proportionality constant (depends on the structure of the ionic solid)
Q1 and Q2 are the charges on the ions
r = distance between centers of the cation and anion
You do not need to know how to calculate lattice energy , but do know the following :
Lattice energy increases (becomes more exothermic) as the ion charges _______________ and the distance between the centers of
the cations and anions ________________________(size of cation/anion decrease).
Example: Which compound in each of the following pairs of ionic substances has the most exothermic lattice energy?
Explain why.
(a) NaCl or KCl
(b) LiF or LiCl
(c) Mg(OH)2 or MgO
(d) Fe(OH)2 or Fe(OH)3
(e) MgO or BaO
*Note – the energy ____________ to break an ionic bond is equal to its lattice energy but opposite in sign.
8.6 Partial Ionic Character of Covalent Bonds
Calculating Percent Ionic Character: Percent Ionic Character = Observed dipole moment
x
100
Calculated dipole moment
No bond reaches 100% ionic character, thus no individual bonds are completely ionic.
Ionic vs. Covalent:
Ionic compounds generally have greater than ______% ionic character.
Ionic compounds generally have electronegativity difference greater than _________.
A compound that conducts an ____________ _____________when melted will be classified as ionic.
8.7 The Covalent Chemical Bond: A Model
When a chemical bond forms, energy is __________________________. When a chemical bond breaks, energy is
__________________________.
Example: Approximately 1652 kJ of energy is required to break a mole of methane, ________, into separate C and H atoms.
Therefore, 1652 kJ is ______________________when 1 mole of methane is formed from gaseous C and H atoms.
Bond energy –
Example: Methane has four identical ___________ bonds. We can calculate the average bond energy for __________.
8.8 Covalent Bond Energies and Chemical Reactions
single bond-
double bond triple bond –
Table 8.4 (p 373) gives average bond energies
Table 8.5 ( p 374) gives bond lengths for selected bonds
A relationship exists between the number of shared electron pairs and the bond length. In general, as the number of shared
electrons_____________________, the bond length __________________, and bond energy _________________.
Bond Energy and Enthalpy
Enthalpy Change –
Bond energy can be used to calculate ∆ H for a reaction.
Using bond energies:
∆H = ∑ Dreactants (bonds broken) - ∑ Dproducts (bonds formed)
Example 1: Calculate the ∆H, enthalpy, using bond energies for the following reaction: H2(g) + F2 (g)  2HF (g)
Example 2: Calculate the ∆H for the following reaction using bond energies: H-C=C-H(g) + H2 (g)  CH2=CH2
8.9 THE LOCALIZED ELECTRON BONDING MODEL
Localized Electron (LE) Model – assumes a molecule is a compound of atoms that are bound together by
_______________ _______________ of electrons using the_____________ _____________ of the bound atoms.
Lone pairs – pairs of electrons ________________ on an atom.
Bonding pairs – pairs of electrons found in the space _________________ atoms.
8.10 LEWIS STRUCTURES
Lewis Structures – shows how the ___________________ electrons are arranged among the atoms in
the molecule/polyatomic ion.
Every period 1 and 2 element (with the exception of H, He, B, and Be) can form compounds of lowest energy if their
highest energy levels are filled (s2p6). This is called the _______________ rule.
Hydrogen follows a ____________rule, it needs _________ electrons to be stable. (He already has a _________)
We will discuss Be and B later.
RULES FOR DRAWING LEWIS STRUCTURES
1. Choose the center atom. It is usually the least electronegative atom. C is always the center atom, H is never the
center atom.
2. Draw a skeletal structure – symmetrically arrange the other atoms around the center atom.
3. To determine number and types of bonds - one strategy you can use is S = N – A
S= shared electrons (those involved in bonding)
N = needed electrons. This is the total number of electrons needed for an atom to be stable (either 8 or 2 – we
will discuss exceptions to this later)
A = available valence electrons
4. Complete the structure by adding lone pairs to complete octets for all atoms
5. Double check the # of electrons used in the structure – must equal A, number of available valence electrons.
Examples
1. PH3
2. HCN
3. CO2
4. phosphite ion
5. carbonate ion
8.11 Exceptions to the Octet Rule
The second row elements B and Be often have fewer than 8 electrons (electron deficient) around them in compounds.
Usually, Be needs (N=) _____ electrons, and B needs (N=) __________ electrons.
Examples:
1. BeCl2
2. BH3
Odd number of electrons: In a few molecules and polyatomic ions, the number of valence electrons is odd and an octet
around each atom cannot be achieved.
Ex. NO
Elements in the third period of the periodic table and beyond often satisfy the octet rule but can exceed the octet rule by
using their empty valence d orbitals (expanded octet)
When using S=N-A, if S is less than the number needed to bond all atoms to the central atom, then an expanded octet is
needed around the center atom.
When writing the Lewis structure for a molecule, bond all atoms to the central atom and satisfy the octet rule for the
atoms first. If electrons remain after the octet rule has been satisfied then place them on the element having available d
orbitals (elements in period 3 or beyond) – this will involve the center atom.
Examples:
1. PCl5
3. RnCl2
2. ClF3
8.12 RESONANCE
Resonance –
Example: nitrate ion
The arrows do not mean the structure “flips” from one resonance structure to another. They simply show the structure is
the __________________ of the resonance structures. Measurements in bond length suggest all 3 N-O bonds lengths are
____________________________.
Bond order – The number of electron pairs involved in a covalent bond. A single bond has a BO of ______, a double
bond ________, and a triple ______. For resonance structures, the BO is the average of the bonds.
Example: For the nitrate ion, the bond order for nitrogen-oxygen bond is
Formal Charge
Resonance usually involves equivalent Lewis structures – contain same number of single and multiple bonds.
Nonequivalent Lewis structures contain different numbers of single and multiple bonds.
When we assign oxidation numbers to atoms involved in a covalent bond, we always count both the shared electrons as
belonging to the more ____________________ atom in a bond. It is useful for keeping track of electrons in a redox
reaction, but not a realistic estimate of charges of atoms in individual molecules. Another definition of charge on an atom
in a molecule, the formal charge, can be used to evaluate Lewis structures and select the most stable structures when
nonequivalent structures are present.
Formal charge – the hypothetical charge on an atom in a molecule or polyatomic ion. It is the difference between the
number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule.
To calculate the formal charge on an atom:
1) Determine the number of valence electrons on the free, neutral atom.
2) Take the sum of the lone pairs electrons and one-half the shared electrons. This is the number of valence
electrons assigned to the atom in the molecule. Subtract this number from the number of valence e- on the
free atom.
The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species.
If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with formal charges
closest to zero and with any negative formal charges on the most electronegative atoms are considered to best describe the
bonding in the molecule or ion.
Example: Draw all resonance structures and select the most stable one for the thiocyanate ion, SCN-.
8.13 MOLECULAR STRUCTURE: THE VSEPR THEORY
Molecular structure – the 3D arrangement of the atoms in a molecule
VSEPR (valence shell electron-pair repulsion ) model is useful is predicting the geometries of molecules
formed from nonmetals. The main postulate is that the structure around a given atom is determined
principally be minimizing the electron-pair repulsions – bonding and nonbonding pairs around a given
atom will be positioned as far apart as possible.
In VSEPR notation the molecule is represented by a formula using the letter A for the central atom, X for
terminal atoms and E for lone pairs of electrons. The table below uses this notation and shows the geometry
of various combinations.
Determining the Shape of a Molecule
The best way to determine the architecture of a molecule is to:
1. Determine what the central atom is.
2. Draw the Lewis structure of the molecule.
3. Determine the number of bonding pairs and lone pairs around the central atom.
4. Refer to the following chart and determine the shape of the molecule.
# of
Bonding
Groups on
Central
Atom
# of
Lone
Pairs
on
Central
Atom
VSEPR
notation
Molecular
Geometry
Bond
Angle
2
0
AX2
Linear
180
Linear
2
1
AX2E
Bent
less than
120
Trigonal
Planar
2
2
AX2E2
Bent
Less than
109.5
Tetrahedral
2
3
AX2E3
Linear
180
Trigonal
Bipyramidal
3
0
AX3
Trigonal
Planar
120
Diagram
Electron Domain
(Electronic)
Geometry
Trigonal
Planar
3
1
AX3E
Trigonal
pyramidal
Less than
109.5
2
AX3E2
T-shaped
90
4
0
AX4
Tetrahedral
109.5
4
1
AX4E
See-saw
90 and
120
4
2
AX4E2
Square planar
90
5
0
AX5
Trigonal
bipyramidal
90 and
120
5
1
AX5E
Square
pyramidal
90
6
0
AX6
Octahedral
90
Tetrahedral
Trigonal
Bipyramidal
Tetrahedral
Trigonal
Bipyramidal
Octahedral
Trigonal
Bipyramidal
Octahedral
Octahedral
PRACTICE : MOLECULAR GEOMETRY
1) Draw the Lewis structure and predict the molecular structure (geometry) – including
bond angles for each of the following:
(a) XeCl2
(b) SeO3
(c) TeF4
(d) SCl2
electron domains are considered either bonds on the central atom or unshared pairs of electrons on the
central atom. note: multiple bonds are considered one domain.
electron domain (electronic) geometry: 3-dimensional arrangement of electron domains
molecular geometry - the 3-dimensional arrangement of the atoms
EX: Water and methane
Draw the Lewis structure of each
Molecular geometry of each
Count the number of electron domains surrounding the central atom.
Determine the electron domain geometry.
Moecular Polarity
Dipole moment – molecule has a center of ______________ charge and a center of ____________________ charge.
Some molecules have polar bonds but do not have a dipole moment. This occurs when the individual bond polarities are arrange in
such a way that they ____________________ each other out.
Example: CO2
If the individual bond polarities do not cancel each other out – then a dipole moment exists.
Example: H2S
Summary
If a molecule has nonpolar bonds, then the molecule itself must be ______________________.
If a molecule has polar bonds and an asymmetrical shape, the molecule will be _______________________.
Examples of asymmetrical geometries:
If a molecule has a symmetrical shape and the bonds are polar, the molecule may be _______________ or _____________.
Examples of symmetrical geometries:
It’s nonpolar if all the atoms bonded to the central atom are the same. Ex:
It’s polar if all the atoms bonded to the central atom are not the same. Ex:
HOMEWORK
(All HW Assignments must be done on notebook paper!!)
HOMEWORK 10/17
Without using an electronegativity table, answer the following:
1. Predict the order of increasing electronegativity in each of the following groups of atoms:
(a) Na, Rb, K
(c) F, Cl, Br
(b) B, O, Ga
(d) S, F, O
2. Predict which bond in the following groups will be most polar:
(a) C-F, Si-F, Ge-F
(c) C-O or Si-O
(b) Se-Br, S-Br, O-Br
`
(d) P-Cl
3. Rank the following bonds in order of increasing ionic character:
N-O, Ca-O, C-F, Br-Br
4. Write the shorthand electron configuration for the most stable ion formed by each of the following
elements:
(a) S
(c) Ra
(b) Br
(d) Cs
5. Write the shorthand electron configuration for the following transition metal ions:
(a) Fe2+
(c) Ni2+
3+
(b) Fe
(d) Mn4+
6. Which of the following ions have noble gas configurations?
(a) V3+, V5+, Ni2+
(c) Rb2+, Sr2+, Sn2+
2+
5+
2+
(b) Ta , Ta , Cd
(d) P3-, S3-, Br37. Give 3 ions that are isoelectronic with argon. Place these in order of increasing size.
HOMEWORK 10/18
1. For each of the following groups, place the atoms and/or ions in order of decreasing size:
(a) V, V2+, V3+, V5+
(d) P, P-, P2-, P3+
+
+
+
(b) Na , K , Rb , Cs
(e) O2-, S2-, Se2-, Te22- +
2+
(c) Te , I , Cs , Ba
2. Predict the empirical formulas of the ionic compounds formed from the following pairs of
elements. Name each compound.
(a) Al and Cl
(c) Na and O
(b) Sr and F
(d) Ca and S
3. Which compound in each of the following pairs of ionic substances has the most exothermic
lattic energy? Justify your answers!!
(a) LiF, CsF
(d) Na2SO4, CaSO4
(b) NaBr, NaI
(e) KF, K2O
(c) BaCl2, BaO
(f) Li2O, Na2S
HOMEWORK 10/21 : BOND ENERGY
Use bond energy values (Table 8.4) to estimate ∆H for each of the following reactions:
1. H-C ≡ N + 2 H2  CH3NH2
2. N2H4 + 2 F2  N≡ N + 4 HF
HOMEWORK 10/22 : LEWIS STRUCTURES
Draw Lewis Structures for the following molecules/ions:
1. nitrogen trifluoride
2. sulfur dioxide
3. chlorite ion
4. perchlorate ion
HOMEWORK 10/23 :LEWIS STRUCTURES
Draw Lewis Structures for the following molecules/ions:
1. BeF2
2. XeF4
3. SF4
4, Br3-
HOMEWORK 10/24 : LEWIS STRUCTURES, RESONANCE, AND FORMAL CHARGE
1. Which of the following have center atoms that obey the octet rule? Draw Lewis Structures
for each.
(a) AsF3
(b) ICl4(c) XeO4
2. (a) Draw 2 possible Lewis structures that obey the octet rule for nitrosyl chloride, NOCl.
(b) Using formal charges, explain which Lewis structure is more likely to be correct.
3. (a) Draw 2 resonance structures for sulfur dioxide.
(b) What is the bond order of the sulfur- oxygen bond?
HOMEWORK 10/25: MOLECULAR GEOMETRY
1) Draw the Lewis structure and predict the molecular structure (geometry) – including
bond angles for each of the following:
(a) SiF4
(b) ICl3
(c) PCl3
(d) PCl5
2) Draw the Lewis structure and predict the molecular structure (geometry) – including
bond angles for each of the following:
(a) ICl5
(b) XeCl4
(c) SeCl6
HOMEWORK 10/28: MOLECULAR POLARITY
Write Lewis Structures and predict both the molecular structure and polarity for the following:
(a) SO3
(b) NF3
(c) IF3
(d) COS
(d) CF2Cl2
(e) SeF6