1. Review of Bonding and Bond Types

2. Valence Electrons
The electrons in the highest occupied level of an element’s atoms.
The number of valence electrons largely determines the chemical properties of an element.
Can be determined by looking at the electron configuration or dot diagram for an element.
Alkali metals have 1, Alkaline Earth metals have 2, Transition metals have 2 or 3, Boron group has 3, Carbon group has 4, Nitrogen group has 5, Chalcogens have 6, Halogens have 7, and the Noble Gases have 8 (an octet).
Chemist Gilbert Lewis explained that atoms form certain kinds of bonds due to atoms tendency to achieve the electron configuration of a noble gas (octet rule).

3. Metallic Bonding
Metals are made up of closely packed cations rather than neutral atoms. The cations are surrounded by mobile valence electrons (delocalized electrons), which drift freely from one part of the metal to another.
The attractions of the free-floating valence electrons for the positively charged metal cations creates the bond that holds the metal together.

4. Metallic Bond Characteristics
Good conductors of electrical current because electrons can flow freely in them. As electrons enter one end of a bar of metal, an equal number leave the other end.
Ductile and malleable due to the metals ability to be forced through a narrow opening to produce a wire or shaped with a hammer without shattering. A sea of drifting valence electrons insulates the metal cations from one another. When a metal is subjected to pressure, the metal cations easily slide past one another like ball bearing immersed in oil.
Also crystalline solids at room temperature (except for mercury).

5. Electronegativity
The tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of other elements.

6. Ionic Bonding
Two elements with very different electronegativities (a difference greater than 1.7).
So… if you have Lithium and Chlorine, what’s the difference? Is it greater than 1.7? If so, then it is an ionic bond.
Elements with high electronegativity pull electron(s) more strongly and therefore steal the electron(s) from the element with a low electronegativity.
The stronger elements therefore becomes negatively charged (anion) since it has gained electrons and the weaker elements becomes positively charged (cation) since it has lost electrons.
Since they have opposite charges, they attract one another by electrostatic forces which creates the bond.
Usually between one non-metal and one metal.

7. Drawing Lewis Structures for Ionic Bonds
The overall charge on the compound must equal zero, that is, the number of electrons lost by one atom must equal the number of electrons gained by the other atom.
The Lewis Structure (electron dot diagram) of each ion is used to construct the Lewis Structure (electron dot diagram) for the ionic compound.
Examples:
Lithium fluoride, LiF
Lithium atom loses one electron to form the cation Li+
Fluorine atom gains one electron to form the anion F-
Lithium fluoride compound can be represented as
Li+
Lithium oxide, Li2O
Each lithium atom loses one electron to form 2 cations Li+ (2 electrons in total are lost)
Oxygen atom gains two electrons to form the anion O2-
Lithium oxide compound can be represented as
Li Li+
Practice…

8. Different Day

9. Covalent Bonding
Two elements with similar electronegativities (difference is less than 1.7).
Since they have similar pull on the electron(s) they end up both holding onto it and the electron(s) and sharing them, which creates the bond.
Usually both non-metals.

10. Structure of Covalent Bonding
Count the number of total valence electrons in the substance.
Write first element as central atom and put all other elements around (except if H is first, then put second element as central).
Draw a line to represent shared pair of electrons between central atom and other elements.
Fill in electrons around all elements until you have 8 (Except H only gets 2, Be only gets 4, B only gets 6)
Count the total number of electrons on your drawing
If equal to calculated electrons from step 1, you are done!
If more electrons on drawing then calculated, take away two electrons from central atom and one other element and make a double bond
If less electrons on drawing then calculated, add extra needed to central atom.
Practice problems: CF4, H2O, CO2

11. Polarity
Polar bonds are when one atoms has a slightly stronger pull on the electrons that are being shared (difference between 1.7 and 0.4). Therefore, one end of the bond is slightly more negative (the one that is slightly stronger) and one is slightly more positive (the one that is slightly weaker).
All the diatomic elements (Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, Iodine and Hydrogen) have non-polar covalent bonds since they share the electrons equally.

12. Many days later…

13. Ionic Bond Characteristics
Compounds are called formula units
Most are crystalline solids at room temperature (pg 422 has some great pictures)
In solids NaCl, each sodium cation is surrounded by six chloride anions and each chlorine anion is surrounded by six sodium cations. Giving them both a coordination number (the number of ions of opposite charge that surrounds the ion in a crystal) of 6.
In this arrangement, each ion is attracted strongly to each of its neighbors and repulsions are minimized, which results in a very stable structure.
High melting point (usually above 300 °C)
High solubility (dissolving ability) in water since water is polar.
Good conductor of electricity either when melted (the orderly crystal structure breaks down allowing cations and anions to move freely through the liquid) or when mixed in water (ions also become free to move around allowing electrons to flow).
Brittle, not very malleable because when struck with a hammer, the blow tends to push ions of like charge into contact and they repel, shattering the crystal.

14. Covalent Bonding Characteristics
Compounds are called molecules.
Can be solid, liquid or gas at room temperature.
Low melting point (usually below )
High to low solubility in water (depending on the molecules’ polarity)
Poor to nonconducting when dissolved in water (aqueous).
A greater range of physical properties is found due to the widely varying intermolecular attractions (polar and nonpolar bonding)
Can have single, double or triple covalent bonds between two atoms.

15. Polar Bonds vs. Polar Molecules
The presence of a polar bond in a molecule often makes the entire molecule polar, where one end of the molecule has a slightly more negative charge and the other end is slightly more positive (called a dipole).
(Ex: H20 and NH3)
However, if the polar bonds are arranged symmetrically around the molecule the entire molecule is actually non-polar.
(Ex: CO2 and CH4)

16. Other Intermolecular Forces
Molecules are often attracted to each other by a variety of other forces, although weaker than ionic or covalent bonds.
Van der Waals forces:
London Dispersion forces-the weakest of all molecular interactions caused by the motion of electrons.
Dipole-induced dipole interactions-occurs when polar molecules cause electrons in another molecule to move to one side of the molecule as they are attracted to the dipole of the other molecule.
Dipole-dipole interactions- occur when polar molecules are attracted to eachother since they have different charges on each side.
Hydrogen bonds:
Attractive forces in which a hydrogen covalently bonds to a very electronegative atom is also weakly bonded to an unshared pair of electrons of another electronegative atom. Since hydrogen only has one electron when it bonds with a highly electronegative atom, the nucleus of the hydrogen atom becomes very electron deficient. Therefore, the sharing of a nonbonding electron pair on a nearby electronegative atom compensates for the deficiency.