1. Chapter 6 – Modern Chemistry
Chemical Bonding
Chapter 6 – Modern Chemistry

2. 6.1 Types of Bonds
A chemical bond is a mutual electrical attraction between the nuclei of one atom and the valence electrons of another atom
Occurs in several ways
Gaining and losing electrons
Ionic Bonds
Sharing electrons
Covalent Bonds
Overlapping electron clouds
Metallic Bonds

3. Determining which is which?
Covalent or Ionic?
Electronegativity differences between atoms will determine whether it is ionic or covalent
Covalent bonds share electrons and will have an electronegativity difference between
Ionic bonds form from ions that have gained or lost electrons and their difference is between 1.7 – 3.3

4. Polarity
Covalent bonds can also show partial charges which is known as polarity
Covalent compounds can be classified as polar or nonpolar depending on their shape and electronegativity differences
Differences between 0 – 0.3 will be nonpolar (shared equally)
Differences between 0.3 – 1.7 will be polar
Example would be water in which we have two O-H single bonds
Electronegativity of O is 3.5 and H is 2.1
Difference of 1.4 which makes the bond polar
Values are found on page 161

6. 6.2 Covalent Compounds
Atoms will form covalent compounds when the valence electrons of one atom become attracted to nucleus of another atom
This attraction will in return have to same affect on the second atoms valence and the first atoms nucleus

7. Formation of a Covalent Bond
This cross attraction that occurs in covalent compounds is formed by the need of all atoms to fulfill the octet rule
This rule states that all atoms are trying to fill their outer energy shell with 8 electrons
The need to do this is a result of achieving their lowest potential energy state
This amount of energy that is given off is the same about of energy that is required to break the bond
This is known as bond energy

8. Electron Dot Diagrams
Electron Dot Diagrams are a method of illustrating valence electrons
They are created by taking the symbol of the element and treating it as if it is inside an imaginary box creating four sides
Place the valence electrons around the symbol one on each side of the box before going back and pairing them up
This will show how many unpaired electrons are available for bonding
Covalent bonds will typically form at the unpaired electrons

9. Lewis Structures
Lewis structures are the illustrations of covalent compounds using the dot diagrams
When creating a Lewis structure, it is important that every unpaired electron is paired up in some fashion
Atoms can create several types of covalent bonds in order to pair these up: single, double, triple, and coordinate covalent bonds
Single – each atom shares one electron
Double – each atom shares two electrons
Triple – each atom shares three electrons
Coordinate – one atom shares two while the other contributes none

10. Examples – Electron Dot Diagrams

11. Examples – Lewis Structures

12. 6.3 Ionic Bonding Formation of Ionic Compounds
Electron Dot Diagrams (not Lewis Structures) can be used to illustrate the formation of ionic compounds
Illustrates the losing of one electron from sodium to form a positive ion and the gaining of one electron by chlorine to form a negative ion

13. Characteristics of Ionic Bonding
Ionic compounds form a repeating pattern of negative and positive ions that will all interact with each other through the entire substance
This is different than covalent compounds that share their electrons to form a single molecule
This interaction forms a crystalline structure known as a crystal lattice
With such a high level of interaction between ions, there is a lot of potential energy stored in the lattice structure
This is known as the lattice energy
It explains the high melting and boiling points of ionic compounds as well as why many of them are solids at room temperature

14. 6.4 Metallic Bonding
In metals, the vacant p orbitals of the their outer energy levels will overlap and allow the valence electrons to move freely throughout the entire substance
Creates a sea of electrons
Because of this overlap, metals have their unique properties
Ductility
Malleability
Good conductors of the heat and electricity

15. 6.5 Molecular Geometry VSEPR Theory
The arrangement of shared pairs of electrons surrounding atoms must be oriented as far apart as possible
If an atom has two shared pairs of electrons, then the arrangement must be 180º separation
If an atom has three shared pairs of electrons, then the arrangement must be 120º separation
If an atom has four shared pairs of electrons, then the arrangement must be 109.5º separation
These occur without any unshared pair of electrons
If an atom has unshared pairs of electrons, it will cause the angles to lessen just slightly due to the unshared pair having a greater force of repulsion than the shared pair
For instance, in water, there are two shared and two unshared but instead of having angles at º, the angle is 105º

16. Molecular Geometry Hybridization
Occurs when two or more atomic orbitals of similar energies are mixed together to produce a hybrid orbital
Carbon is a good example of this process
Outer configuration is 2s2 2p2 but instead of only having two unpaired electrons these two orbitals blend to form an sp3 orbital at the second energy level
This will also help explain why certain atoms can break the octet rule and why shape sometimes will affect polarity

17. Intermolecular attractions
Molecules, which have atoms covalently bonded to each other, are also attracted to other molecules by various forces
These forces are dipole-dipole interaction, hydrogen bonding, and London dispersion forces
Dipole-dipole interaction is the attraction between the partial positive and negative ends of a polar molecule
Hydrogen bonding is the attraction between an unshared pair of electrons and the hydrogen atom (when the hydrogen is connected to a highly electronegative atom)
London dispersion forces is the weakest form of attraction that occurs when instantaneous dipoles occur due to the movement of the electrons