1. Chapter 8: Basic Concepts of Chemical Bonding
Advanced Chemistry
Chapter 8: Basic Concepts of Chemical Bonding
Sections 8.1 – 8.5 Notes

2. What’s a Chemical Bond?
Whenever atoms or ions are strongly attached to one another, we say that there is a chemical bond
Three types of Chemical Bonds
Metallic
Ionic
Covalent

3. Bond Types – Brief Review
Ionic
Refers to the electrostatic forces that exist between ions of opposite charge
Covalent
Results from the sharing of electrons between nonmetals
Metallic
Are relatively freely moving electrons found between metals

4. Bond Polarity
Bond polarity is determine by differences in electronegativity
Non-Polar Covalent 0.0
Polar Covalent
Ionic Bond
>2.0

5. Dipole Moments
Covalent compounds with differences in electronegativity produce dipole moments, that’s why they are polar molecules!
Dipole moment increases with charge!
u = Qr

6. Calculating Dipole Moments
u = Qr
Where:
u = dipole moment measured in debyes (D)
Q = product of charges of atoms involved in bond
r = separation of charge in meters (m)
Convert using 1D = 3.34 x C·m
The distance between the centers (bond length) of H and Cl atoms
in the HCl molecule is 1.27 Angstrom.
A) Calculate the dipole moment using a +1 and -1 charge respectively

7. Ionic Bonding
Attraction between ions..generally metals and nonmetals..this you know..
But what happens to energy in ion formation?
Forming a cation (metals losing an electron) is an endothermic process..meaning energy is put in to remove an electron
Forming a anion (done by nonmetals) is an exothermic process
The difference between the processes is the overall energy change for one mole of reactant

8. Ionic Bonds and Lattice Formation
The ions are drawn together and a lattice of ionic structure is formed
The Lattice Energy is the measure of how much stabilization results from the arranging of oppositely charged ions in an ionic solid
It is, the energy required to completely separate a mole of a solid ionic compound into its gaseous ions

9. Lattice Energy Example
NaCl(s) Na+(g) + Cl- (g)
Hlattice = +788 kJ/mole
This means the forming of NaCl is highly
exothermic, H = -788 kJ/mole

10. Lattice Energies r2 Q1Q2 E = k Potential energy of two
The large positive endothermic lattice energies makes ionic bonds strong..
The strong attractions also make the compounds hard, brittle materials with high melting points
Potential energy of two
interacting charged particles
relates by this equation
Q1Q2
E = k r2

11. Practice Problem
Which substance would you expect to have the greatest lattice energy?
AgCl, CuO, or CrN
The greatest lattice energy results from the largest
product of the ionic charges…thus
CrN has (+3)(-3) = 9

12. Electron Configuration of Ions
Ions like loose or gain electrons to form noble-gas electron formations
This results in the most energy-favorable and stable formation
Even though an increase in ionic states would result in a higher lattice energy, it is not enough to remove an electron from a completed energy level or add to an unfavorable higher energy level

13. e- configuration of Transition Metals
Transition metals (d block) cannot reach the noble gas configuration due to their location on the table
So…they achieve stability by loosing electrons from the highest n shell..
So, they loose valence electrons first, then as many d electrons as are required to reach the charge on the ion

14. Example
Fe [Ar] 4s2 3d6
In forming the Fe3+ ion, 2e- are lost from the 4s subshell and 1 from 3d so…
Fe3+ [Ar] 3d5

15. Practice Problem
Write the electron configuration for Cr3+
[Ar] 3d3

16. Sizes of ions
Ionic size plays a crucial role in determining the structure and stability of ionic solids
It determines both the lattice energy of the solid and the way in which the ions pack in a solid
Ionic size also determines the properties of ions in solutions

17. Ion Size
Ion size depends on nuclear charge, the number of electrons it possesses, and the orbitals in which the outer electrons exist
Cations are smaller than their parent atoms
Anions are larger than their parent atoms
For Ions with the same charge, size increases as we go down a group in the PT

18. Ion Size and isoelectronic series
The term isoelectronic means that the ions possess the same number of electrons
Ex: O2-, F-, Na+, Mg2+ and Al3+
All have the configuration of Neon
The nuclear charge increases while # of e remain the same…so
Radius decreases due to larger attractive force between nucleus and electrons
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