Chapter 9- Covalent Bonds Agenda- Lab - Review - Quiz – Review –Chapter 8 / 9 Test – Chapter 8/9

Section 1  Why do atoms bond?  To become noble or stable  To achieve an octet (are exceptions) Covalent Bonds

What is a covalent Bond?  Elements share electrons  Majority form between nonmetallic elements Result?  A Molecule is formed Covalent Bonds

Lewis Dot Review: In your notes draw the following dot structures  H  N  O  C  S  Cl  Ar

Groups and Bonds  Group 15 = 3 Bonds PH 3  Group 16 = 2 Bonds H 2 S  Group 14 = 4 Bonds  CCl 4  Group 17 = 1 Bond HCl Lewis structures

Sigma Bond  The single covalent bond is calls the… “Sigma Bond”  Shared electrons between two atoms

Multiple Covalent Bonds Why Multiple Bonds? Hint: Think Noble.  To achieve an Octet! Example: C 2 H 4  Draw the central atoms  Attach the surrounding atoms  Make sure each atom has an octet

Sigma and pi Bonds Sigma Bonds  Two atoms share electrons pi Bonds  Parallel orbitals over lap  Forms double bonds Example: C 2 H 4

Let’s take a closer look

Lets take a closer look

Strength and Energy Bond Strength  The shorter the bond length, the stronger the bond, the greater the bond-dissociation energy Bond Energy  Endothermic – more energy is needed to break the bond than is released  Exothermic – more energy is released during bond formation than is required to break it.

Section 2: Naming Covalent Molecules  Different than Ionic 1. First element = entire name 2. Second element = root + ide 3. Prefixes used to indicate the # of each type present in compound Naming Covalent

Prefixes for Covalent Molecules Prefixes

Example P205P205  Follow your rules! Phosphorusoxidedipent

Naming Acids 2 types of acids 1. Binary Acids  HCl, H 2 S, HBr, HCN 2. Oxyacids  H 2 SO 4, HClO 3, HClO 2

Binary Acids Example: HCl 1. Hydro + root of second element  Hydrochlor… 2. Add –ic then acid  Hydrochloric acid Name the following: HBr, HI, HF, HCN

Oxyacids Example: H 2 SO 4 and H 2 SO 3 1. Root of oxyanion present  Sulfur… 2. If oxyanion ends in …ate add -ic to the end  H 2 SO 4 = sulfuric acid 3. If oxyanion ends in …ite add –ous to the end  H 2 SO 3 = sulfurous acid

Section 3  Molecular Structures

Molecular Structures  Structural Formula  Uses letter symbols and bonds to show relative positions of atoms. Section 3:

Lewis Structures  Determining Lewis Structures Step 1…  Predict the location of atoms  H is always terminal (end)  Central atom has the least attraction for shared electrons (closer to the left of the periodic table)

Lewis Structures Step 2…  Find total # of valence electrons Step 3…  Determine # of bonding pairs. Divide # of valence electrons by 2 Step 4…  Place 1 pair (single bond) between the central atom and terminal atoms

Lewis Structures Step 5…  Subtract pairs used from total possible pairs (step 3)  Place remaining pairs around terminal and central atom (octet) Step 6…  If central atom does not have octet, use lone pairs as double bonds.

Example: Carbon dioxide  Step 1- predict location  Step 2 – Total Valence Electrons =  Step 3 – Divide by 2 = pairs  Step 4 – Central Atom bonds  Step 5 – Place remaining pairs  Step 6 – Check Octet rule Lewis Structures COO : : : : : : : : :: : : : : : : 16 8

OOC :: : : : : Lewis Structures WHAT’S WRONG WITH THIS PICTURE? Carbon ~ octet? Move electron pairs on each O to achieve octet around C

Charge Molecules Positive Charges…  You must remove electrons from the total electrons available for bonding according to the charge.  Example NH 4 + Total Valence Electrons = Subtract the charge Divide by 2 (8/2 = 4) ~ bonding pairs N H H H H + (9-1 = 8) 9

Charged Molecules Negative Charges…  You must add electrons to the total electrons available for bonding according to the charge.  Example PO 4 3- = Total Valence Electrons = ADD the charge Divide by 2 (32/2 = 16) ~ bonding pairs P O OO O : : :: :: : : : : : : 3- ( = 32) 29

VSEPR Model V alence S hell E lectron P air R epulsion  Electrons are located as far apart as they can be  Shared electron pairs repel one another  Lone pairs also repel (even more) Hybrid Orbitals  S and p orbitals change to form new equal orbits  Each bond between atoms represents an s, p or d orbit

Visualizing the Models Example #1: BeCl 2  1 st Draw the Lewis dot.  Determine the # of shared pairs and lone pairs around the central atom.  Shared pairs = 2  Lone pairs = 0  2 Total hybrid bonds  S and p (sp)

Example #1: AlCl 3 (Exception to the Octet Rule)  1 st Draw the Lewis dot. Visualizing the Models  Determine the # of shared pairs and lone pairs  Shared pairs = 3  Lone pairs = 0  3 Total hybrid bonds  s, p and another p (sp 2 )

Example #1: CH 4  1 st Draw the Lewis dot. Visualizing the Models Determine the # of shared pairs and lone pairs around the central atom! Shared pairs = 4 Lone pairs = 0 4 Total hybrid bonds s, p, p and another p (sp 3 )

Example #1: PH 3  1 st Draw the Lewis dot. Visualizing the Models  Determine the # of shared pairs and lone pairs  Shared pairs = 2  Lone pairs = 1  2 Total hybrid bonds  S and p (sp 3 )  Shape  Trigonal Pyramidal PH H H :

Electronegativity and Polarity Even or uneven sharing of electrons. Determined by the electronegativity Identical atoms share evenly Bonds between to different atoms.  one atom pulls the electrons closer  creates a relative negative and positive side of the molecule RULES: difference between electronegative #s = Nonpolar covalent = polar covalent > 1.7 = ionic bond

Properties of Covalent Compounds Solubility  Polar Molecules soluble in polar substances  Non-polar in non-polar Intermolecular force between molecules is called the “van der Waals” force 3 types of intermolecular foces 1. Nonpolar (weak) = dispersion forces 2. Polar (weak)= dipole-dipole force 3. Hydrogen Bonds  Very Strong  between H and (F, N, O)

Molecular Shapes sp sp 2 sp o 120 o o o o

Molecular Shapes sp 3 d sp 3 d 2 90 o 90 o / 120 o

Ionic or Covalent Does the compound contain a metal? Is the metal a transition metal? Use I, II, III, IV, V – to indicate the charge of the metal Don’t use roman numerals: Don’t use prefixes The compound is covalent; use prefixes Example: N 2 O – dinitrogen monoxide P 2 O 5 – diphosphorus pentoxide Example: FeO – Iron (II) oxide Cu 2 S – Copper (I) Sulfide Example: NaCl – sodium chloride CaCl 2 – calcium chloride YES NO YESNO