1. Introduction Presentation
Chapter 8 Bonding
Introduction Presentation

3. 8.1 Types of Bonding
Ionic Bonding – Type of bonding where electrons are transferred from a metal ion to a non-metal ion
Based on differences in electronegativity and ionization energy

6. Coulomb’s law
Used to calculate the amount of energy that is involved in the interaction between a pair of ions
Equation:

7. Coulomb’s Law
Negative sign indicates an attractive force between the ions
Tells us that the ion pair has lower energy when bonded than when apart as individuals
The closer the pair is when bonded, the more exothermic the process is
Can also be used to calculate the repulsive forces that come into play between like charged atoms
In the case of diatomic molecules, we can use coulomb’s law to find the exact spot where bond energy favors the formation of a chemical bond

8. Bond Length Distance at which the energy of a system is at a minimum
All systems will bond in a way that maximizes the attractive forces and minimizes the repulsive forces between atoms/ions
At short distances the repulsive forces rise because of the closeness of the particles
At large distances the attractive forces are weak and so atoms are unable to bond

9. Covalent BOnd
Formed when two atoms (non-metals) share 1 or more electrons
Electrons can be shared equally (non-polar) or unequally (polar)
All polar covalent molecules will have partial charges, and thus a positive and negative pole

14. 8.2 Electronegativity
Ability of an atom to attract electrons to itself, based on nuclear pull/strength
Metals have low electronegativity values, non-metals high
Trend occurs for the same reasons as all other trends we’ve observed

16. Characterizing Bonds
Greater electronegativity difference between bonding elements leads to a less covalent bond, and more ionic bond
We do not use the values from Regents chemistry
We use % ionic character and so if a compound is able to conduct an electric current when melted, its ionic
Generally your % ionic character is greater than 50%

17. 8.3 Bond Polarity and Dipole Moments
Dipolar Molecules -- Molecules with a slight negative end, and a slight positive end due to a difference in electronegativity values between atoms
In an electric field these molecules will orient themselves

18. Molecules with Polar Bonds but NO dipole
These molecules are oriented in such a way that the charge differences are equally distributed over the entire molecule
In essence, the charge differences are “cancelled” out
No one end of the molecule is greater in charge than the other
Linear, tetrahedral, and planar molecules all have the ability to follow this rule

20. 8.4 Electron configurations and sizes
Placement of the elements on the periodic table suggests how many electrons are lost or gained to achieve a noble gas configuration
Group 1 always lose 1
Group 2 always lose 2
Group 16 gains 2
Group 17 gains 1

22. Ionic Radii
Anions– gain electrons, have negative charges, and are larger in radius than the atom they started out as
Cations – lose electrons, have positive charges, and are smaller than the atom they started out as.
Cations are significantly smaller because they lose an entire energy level when bonding

23. Isoelectronic Series
Ions that have the same number of electrons when done with a reaction
All ions in this series share the same noble gas configuration when bonded to another atom
Size decreases as the nuclear charge increases across any isoelectronic series