Covalent Bonding Day 1

Ionic Bonding vs Covalent Bonding While we have been writing Ionic bonds as single formulas (ex. NaCl), there is no such thing as a single molecule of NaCl Ionic bonds keep repeating to form a crystal, with the cation and anion repeating again and again. Ionic bonds keep repeating, forming a crystal, until the crystal breaks off. No such thing as 1 NaCl molecule. Cation and Anion size

Ionic Bonding vs Covalent Bonding The big difference between ionic and covalent is that covalent molecules are individual molecules Is such thing as individual molecules with covalent bonds. Individual molecules are not chemically bonded to each other. Separate entities.

Covalent Bonds Covalent bonds are formed when two or more atoms (usually non-metals) SHARE electrons Atoms share electrons instead of giving or taking because there is a small difference (or no difference) in electronegativity

Lewis Structures Shows the arrangement of valence electrons in a bond Electrons are not transferred using arrows, but drawn in-between two atoms to show that they are shared IONIC COVALENT Na* Cl*******  Na+ and Cl- which  NaCl whose electron configuration does Na+ have? Cl-? Ca and Cl  CaCl2

Steps for Drawing Lewis Structures Add up the total number of valence electrons from the formula. Arrange the atoms so there is a central atom with the rest surrounding it. Draw in single bonds between the atoms using 2 electrons per bond. Now figure out how many electrons are left over. Fill in leftover electrons around each atom to give them each 8 electrons. Having 8 Electrons is known as having a “full octet” Hydrogen can only have 2 electrons- is still full Pool together electrons, can think of it like a bank account

Example: CH4 Step 1: Add up the total number of valence electrons from the formula. CH4 = 1 C and 4 H’s C = 4 electrons H = 1 electron x 4 = 4 electrons Total = 8 Valence Electrons

Example: CH4 Step 2: Arrange the atoms so there is a central atom with the rest surrounding it. This is usually the first atom in the formula Hydrogen is NEVER a central atom H H C H

Example: CH4 Step 3: Draw in single bonds between the atoms using 2 electrons per bond.

Example: CH4 Step 4: Figure out how many electrons are left over. Fill in leftover electrons around each atom to give them a full octet. Started with 8 Electrons to work with Used 8 Electrons for initial bonds = 0 Electrons left Don’t need to fill in any leftover electrons Carbon has 8 electrons around it Hydrogen has 2 electrons around it

Lewis Structures If you have enough electrons to create all of the bonds and give every atom a full octet (2 for Hydrogen, 8 for the rest) with no electrons left over then you have the right structure. Too few electrons– you will need a double or triple bond Too many electrons– double check your math!

Practice CF4 NH3 H2O

CCl4 CH2F2 HF PBr3 C2H6 H2 SiCl4 H2S

Covalent Bonds Day 2

Write the Lewis Structure for SO2 What’s the problem?

What if we don’t have enough? If you do not have enough electrons, you need to add double bond(s) and/or triple bond(s) Don’t go too fast – try ONE double bond first, and if that doesn’t work try another double bond or try a triple bond instead. Let’s try SO2 again…

General Information Always try a formula with single bonds first and count how many electrons you used… If you have enough for all single bonds and dots you are fine If you used 2 more electrons than you had available you need a double bond If you used 4 more electrons then you had available you need 2 double bonds OR a triple bond Hydrogen can only have 1 single bond, it will NEVER have a double bond or triple bond

Few more practice problems GeO SiO2 C2H2Cl2

Polyatomic Ions Still draw lewis structures the same way for polyatomic ions The charge on the polyatomic ion gets added or subtracted to the total number of valence electrons you have to work with! Polyatomic Ions with a negative charge: Add that many electrons to the total at the beginning Add because negative charge means it has extra electrons Example: SO42- Total Valence Electrons = 30 Valence Electrons: (add charge) +2 S = 6 New Total = 32 O = 6 x 4 = 24

Polyatomic Ions Polyatomic Ions with a positive charge: Subtract electrons from the beginning total. Subtract because positive charge means it has lost electrons Example: NH4+1 Total Valence Electrons = 9 Valence Electrons: (subtract charge) -1 N = 5 New Total = 8 H = 1 x 4 = 4

Draw the lewis structure for… SO3 2- Valence Electrons: S = O =