Chapter 8 Covalent Bonding VSEPR Theory Molecular Shape

Chapter 8 Covalent Bonding VSEPR Theory Molecular Shape Polar or Non. Polar Properties of Molecular Substances

Why Do Atoms Bond? l The stability of an atom, ion or compound is related to its energy – lower energy states are more stable. l Metals and nonmetals gain stability by transferring electrons (gaining or losing) to form ions that have stable noble-gas electron configurations. – Ionic Bonding l Another way atoms can gain stability is by sharing valence electrons with other atoms, which also results in noble-gas electron configurations. – Covalent Bonding

v The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

The Covalent Bond v Atoms will share electrons in order to form a stable octet. v Covalent bond : the chemical bond that results from the sharing of valence electrons v also called a molecular bond

The Molecule vformed when two or more atoms bond covalently v. The smallest piece in a covalent compound v. Formed when the proton of one atom is attracted to the electron cloud of another atom.

M d o s l e M ole cu les

Single Covalent Bonds l l In a single covalent bond a single pair of electrons is shared This can be represented with a Lewis structure A single line represents a single covalent bond A single pair of electrons

• Bonding pair: a pair of electrons shared by two atoms • Lone pair: an unshared pair of electrons on an atom

Formation of Water

Group 17 elements will form one covalent bond.

Group 16 elements will form two covalent bonds.

Group 15 elements will form three covalent bonds.

Group 14 elements will form four covalent bonds.

Strength of Covalent Bonds l The strength of covalent bonds is determined by the bond length: – distance between the bond nuclei l Bond length is determined by: – The size of the atoms involved —larger atoms have longer bond lengths – How many pairs of electrons are shared—the more pairs of electrons shared, the shorter the bond length is.

Bond Dissociation Energy v v v the amount of energy required to break a bond Indicates the strength of a covalent bond When a bond forms, energy is released; When a bond breaks, energy must be added Each covalent bond has a specific value for its bond dissociation energy.

Bond Energy and Bond Length l A direct relationship exists between bond energy and bond length – – – Shorter Bond Stronger Bond Higher Bond Dissociation Energy Longer Bond Weaker Bond Lower Bond Dissociation Energy

Energy Changes l l An endothermic reaction is one where a greater amount of energy is required to break a bond in reactants than is released when the new bonds form in the products. An exothermic reaction is one where more energy is released than is required to break the bonds in the initial reactants.

Common Names l Many compounds were discovered and given common names long before the present naming system was developed (water, ammonia, hydrazine, nitric oxide).

Structural Formulas l A structural formula uses letter symbols and bonds to show relative positions of atoms.

Lewis Structures l l Used to predict the structural formula Show arrangement of the atoms and unbonded electrons

Five steps to draw Lewis structures: 1. 2. Count the total number of valence electrons in all atoms involved. Decide how the elements are arranged in the structure and draw it out. 1. 2. 3. 4. 3. Hydrogen is always an end atom. Central atom is usually written first in compound Central atom has least attraction for the electrons Usually closer to left on periodic table Subtract the # of electrons used in the bonds.

4. 5. 6. Satisfy the octets of the terminal atoms. Place any remaining electrons around the central atom to satisfy its octet. If the central atom cannot be satisfied, make a multiple bond using a lone pair from the terminal atoms. Check your work

Examples

Drawing Lewis structures for polyatomic ions is very similar to drawing Lewis structures for covalent compounds EXCEPT in finding the number of electrons available for bonding Count the total number of valence electrons in all atoms involved. v If the polyatomic ion is negatively charged, ADD the charge to the number of valence electrons. v If the ion is positively charged, SUBRACT the charge from the number of valence electrons. v Follow the rest of the steps to drawing Lewis structures. v

Examples

Sub Octet v Some compounds form with fewer than 8 electrons present around an atom. v Boron v BF 3

Molecular Shape l l How a molecule “looks” Determines properties The shape of a molecule determines whether or not two molecules can get close enough to react We describe shape using the VSEPR model

VSEPR This model is based on the fact that electrons pairs will stay as far away from each other as possible Valence Shell Electron Pair Repulsion

How to apply VSEPR 1. 2. 3. 4. Draw the Lewis Structure for a Molecule Count the pairs of bonded electrons Count the pairs of unbonded electrons Match the information with the VSEPR chart to classify the shape of the molecule

v Atoms will assume certain bond angles: the angle formed by any two terminal atoms and the central atom v Lone pairs take up more space than bonded pairs do.

http: //www. chem. purdue. edu/gchelp/vsepr/structur. html

Molecular Polarity l l Molecules are either polar or nonpolar depending on the bonds in the molecule. We must look at the shape (geometry) of a molecule to determine polarity. Symmetric molecules are nonpolar. Asymmetric Molecules are Polar

Polar or Non. Polar l Determine if a molecule is polar or nonpolar by – Looking at a model of the molecule – Looking at a Lewis Structure of the molecule

Solubility of Polar Molecules l l l Bond type and shape of the molecule determine solubility Polar substances and ionic substances will dissolve in polar solvents Nonpolar substances will only dissolve in nonpolar substances

Intermolecular Forces v Nonpolar Molecules: Van der Waals intermolecular Forces v Very Weak forces between molecules v Polar Molecules: have dipole-dipole intermolecular bonding. v Stronger intermolecular Forces v Polar Molecules with Hydrogen Bonding: hydrogen bonded to nitrogen, oxygen or fluorine, it will have hydrogen bonding between molecules. A very strong dipole-dipole interaction v Very strong intermolecular Forces v High boiling points, high melting points