Chapter 8: Covalent Bonding

The goal of all Chemical Bonds Elements bond to have a full outer shell like noble gases. Ionic bonds: lose and gain electrons Covalent bonds: share electrons

Covalent Bonds Occur between two nonmetals Molecular compounds have covalent bonds. A bond in which two atoms share ONE pair of electrons between them is a SINGLE COVALENT BOND

Hydrogen gas (H2) Hydrogen gas wants to look like Helium. To do this it needs to get one more electron.

Diatomic Molecules Atoms that exist as a pair in nature There are 7 diatomic elements MEMORIZE!

Hydrogen = H2 Nitrogen = N2 Oxygen = O2 Fluorine = F2 Chlorine = Cl2 Bromine = Br2 Iodine = I2 Hydrogen gas Nitrogen gas Oxygen gas Fluorine gas Chlorine gas Bromine Iodine

Naming Covalent Use prefixes to distinguish the number of each atom present in the compound Example: C2H4 = dicarbon tetra hydride You still add the ending –ide to the last element

Prefixes 1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca

Examples

Common Names for some… Water = H2O Ammonia = NH3 Methane = CH4 Ozone = O3

Goal: to understand the shape of compounds Lesson Objective: To be able to draw Lewis Dot Structures Goal: to understand the shape of compounds

How to draw Lewis Dot Structures Determine the number of valence electrons for each atom by using the periodic table. Calculate the total number of electrons Be sure to multiply the number of valence electrons if there is more than one atom of the element in a compound Write the symbols for each atom and the correct amount for each element

What order do I write the symbols in for a molecule? Generally the first element in the formula goes in the center. Good rule of thumb – the least electronegative element goes in the center. Some elements NEVER go in the center (like: H, halogens)

Remember… Each element only needs 8 electrons 1 dot = 1 electron Only two dots per side * Remember, H and He are exceptions, they can have just 2 dots

Once you have the symbols written, draw a solid line between each element. This represents the SINGLE COVALENT BOND. For each line, subtract 2 from your valence electron total. Each line = 2 electrons

Distribute remaining dots remembering to satisfy the octet rule (8 electrons per symbol) If you run out of dots, you might need to make a double or triple bond. Double bonds: atoms share TWO PAIRS of electrons. (4 electrons total) Triple Bonds: atoms share THREE pairs of electrons. (6 electrons total)

Double and Triple Covalent Bonds Double bonds: atoms share TWO PAIRS of electrons. (4 electrons total) Triple Bonds: atoms share THREE pairs of electrons. (6 electrons total)

Example: HBr # Valence electrons = Write the symbols:

Examples: H2O NF3

Examples O2 CO2 CO N2

Resonance Sometimes there is more than one correct way to write a Lewis dot structure. Example: O3

Exceptions to the octet rule… P and S can hold more than 8 electrons in some molecules Ex) PBr5 Boron only requires 6 electrons to be stable Ex) BH3

Lewis Dot Structures for Ions Works the same way, BUT For each + charge subtract an electron. For each – charge add an electron Ex) Cl-  8 valence electrons

16.3 – Bond Polarity Atoms participating in covalent bonds share electrons, but they do not always share equally. Recall: Electronegativity: the tendency for an atom to attract electrons to itself when chemically combined with another element.

Nonpolar covalent bonds Electrons are shared equally  Examples: H2, O2, N2 How do we know? Think about tug of war…

Tug of War Both Hs have an electronegativity of 2.1 (table 14.2, p 405). Equal electronegativy = share electrons equally = nonpolar bond

Polar Bonds Don’t share electrons equally. Have a difference in electronegativity of .4 – 2 Ex) HCl H: 2.1 Cl: 3.0 Difference of .9 so polar bond.

Polar bonds have partial charges on atoms Polar Bonds - Notation More electronegative element d means partial O -- H H -- O Polar bonds have partial charges on atoms

Chemical Bonding Ionic Bonds Covalent Bonds >2 Non polar Polar 0-0.4 0.4-2

Non Polar Molecules Symmetric shapes with similar atoms. Ex) tetrahedral, linear, trigonal planar… To get this, most times you must draw lewis dot structures.

Polar Molecules Molecules that have partially positive end and one partially negative end. Dipole: molecule with two charged regions (or “poles”)

Remember…Just because it has polar bonds doesn’t make it polar… Polar bonds can “cancel” each other out and result in a non polar molecule.

Intermolecular Forces (IMF) Attractions BETWEEN MOLECULES van der Waals Forces Dispersion Forces Dipole Interactions Weakest type of intermolecular interactions

Dispersion Forces Weakest of all IMFs Caused by motion of electrons Strength of Dispersion forces increases with increasing size (molecular weight).

Dipole Interactions Occurs when polar molecules are attracted to each other

A hydrogen bond is 5% the strength of the average covalent bond. Hydrogen Bonding Very strong dipole interaction between H and F, N, or O. A hydrogen bond is 5% the strength of the average covalent bond.

Ionic vs. Covalent ** See table 16.5 ** Melting Point: ionic > covalent Electrical Conductivity Ionic > covalent Makeup of elements: Solubility: Covalent depends on make up of molecule