Chapter 8
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Chapter 8. Covalent Bonding VSEPR Theory Molecular Shape Polar or NonPolar Properties of Molecular Substances. Why Do Atoms Bond?. The stability of an atom, ion or compound is related to its energy lower energy states are more stable.

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Chapter 8

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Chapter 8

Chapter 8

Covalent Bonding

VSEPR Theory

Molecular Shape

Polar or NonPolar

Properties of Molecular Substances

Why do atoms bond

Why Do Atoms Bond?

  • The stability of an atom, ion or compound is related to its energy

    • lower energy states are more stable.

  • Metals and nonmetals gain stability by transferring electrons (gaining or losing) to form ions that have stable noble-gas electron configurations.

    • Ionic Bonding

  • Another way atoms can gain stability is by sharing valence electrons with other atoms, which also results in noble-gas electron configurations.

    • Covalent Bonding

Chapter 8

  • 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

The Covalent Bond

  • Atoms will share electrons in order to form a stable octet.

  • Covalent bond : the chemical bond that results from the sharing of valence electrons

  • also called a molecular bond

The molecule

The Molecule

  • formed when two or more atoms bond covalently

  • The smallest piece in a covalent compound

  • Formed when the proton of one atom is attracted to the electron cloud of another atom.

Chapter 8



Single covalent bonds

Single Covalent Bonds

  • 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

Chapter 8

  • Bonding pair: a pair of electrons shared by two atoms

  • Lone pair: an unshared pair of electrons on an atom

Formation of water

Formation of Water

Group 17 elements will form one covalent bond

Group 17 elements will form one covalent bond.

Group 16 elements will form two covalent bonds

Group 16 elements will form two covalent bonds.

Group 15 elements will form three covalent bonds

Group 15 elements will form three covalent bonds.

Group 14 elements will form four covalent bonds

Group 14 elements will form four covalent bonds.

Sigma bonds

Sigma Bonds

  • Single covalent bonds are also called sigma bonds:

  • the electron pair is centered between two atoms.

Multiple covalent bonds

 Multiple Covalent Bonds    

When more than one pair of electrons is shared, a multiple covalent bond is formed

Multiple bonds are made up of sigma bonds and pi bonds: formed when parallel orbitals share electrons.

Double covalent bond

Double Covalent Bond

Two pairs of electrons are shared

Contains one sigma and one pi bond.

Triple covalent bond

Triple Covalent Bond

Three pairs of electrons are shared

Has one sigma and two pi bonds.

Strength of covalent bonds

Strength of Covalent Bonds

  • The strength of covalent bonds is determined by the bond length:

    • distance between the bond nuclei

  • 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

Bond Dissociation Energy

  • 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

Bond Energy and Bond Length

  • 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

Energy Changes

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.

Naming molecules

Naming Molecules

  • Molecular Formula

    • Shows what atoms and how many are in a molecule

    • Examples:

    • Nonmetal-Nonmetal Combinations

Naming binary compounds

Naming (Binary Compounds)

The first element is always named first using the entire element name

The second element is named using its root and adding the suffix -ide

Prefixes are used to indicate the number of atoms of each element that are present in the compound

Common names

Common Names

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

Binary acids

Binary Acids

An acid that contains hydrogen and one other element Ex. HCl

ion ends –ide.

Name the acid with hydro-root of the anion-ic

HCl (hydrogen and chloride ) becomes hydrochloric.

HCl in a water solution is called hydrochloric acid.

Naming acids

Naming Acids

  • Acids contain hydrogen as the first element

    • Binary Acids: H bonded to one other element

      • An ion that ends –ide

      • Name the acid with hydro-root-ic

      • Example: HCl

        • Hydrogen ion and chloride ion

        • Hydrochloric acid



  • An acid that contains both a hydrogen atom and an oxyanion.

  • Example: HNO3

  • Identify the oxyanion present.

  • name ends with the suffix –ate, replace it with the suffix –ic.

  • If the name ends with suffix –ite, replace it with suffix –ous,

  • NO3 is the nitrate ion so the acid is nitric acid.

Acid naming summary

Acid Naming Summary

Structural formulas

Structural Formulas

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

Lewis structures

Lewis Structures

Used to predict the structural formula

Show arrangement of the atoms and un-bonded electrons

Five steps to draw lewis structures

Five steps to draw Lewis structures:

  • Count the total number of valence electrons in all atoms involved.

  • Decide how the elements are arranged in the structure and draw it out.

    • 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.

Chapter 8

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 



Chapter 8

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.

  • If the polyatomic ion is negatively charged, ADD the charge to the number of valence electrons.

  • If the ion is positively charged, SUBRACT the charge from the number of valence electrons.

  • Follow the rest of the steps to drawing Lewis structures.



Resonance structures

Resonance Structures

When a molecule or polyatomic ion has both a double bond and a single bond, it is possible to have more than one correct Lewis structure:



a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion.

The structures are called resonance structures.

A molecule that undergoes resonance behaves as if it has only one structure.

Exceptions to the octet rule

Exceptions to the Octet Rule

Three Ways Molecules Might Violate the Octet Rule

Odd number of valence electrons

Odd Number of Valence Electrons

  • Some molecules have an odd number of valence electrons and cannot form an octet around each atom

  • Example: NO2

Sub octet

Sub Octet

  • Some compounds form with fewer than 8 electrons present around an atom.

  • Boron

  • BF3

Coordinate covalent bond

Coordinate covalent bond

  • when one atom donates an entire pair of electrons to be shared with atoms or ions that need two more electrons.

  • Boron compounds often do this

Expanded octet

Expanded Octet

  • Some elements can have more than eight electrons in their valence shell

  • Because of d-level electrons

  • PCl5

Chapter 8

  • How?

    • The d orbital starts to hold electrons.

    • This occurs in atoms in Period 3 or higher.

    • When you draw Lewis structures for these compounds, extra lone pairs are added to the central atom OR the central atom will form more than four bonds.

Things to remember

Things to Remember

  • Any exceptions to the Octet Rule are on the central atom

Molecular shape

Molecular Shape

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



This model is based on the fact that electrons pairs will stay as far away from each other as possible






How to apply vsepr

How to apply VSEPR

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

Chapter 8

  • Atoms will assume certain bond angles: the angle formed by any two terminal atoms and the central atom

  • Lone pairs take up more space than bonded pairs do.

Molecular polarity

Molecular Polarity

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 Nonpolar

Polar or nonpolar

Polar or NonPolar

  • 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

Solubility of Polar Molecules

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

Intermolecular Forces

  • Nonpolar Molecules: Van der Waals intermolecular Forces

    • Very Weak forces between molecules

  • Polar Molecules: have dipole-dipole intermolecular bonding.

    • Stronger intermolecular Forces

  • 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

    • Very strong intermolecular Forces

    • High boiling points, high melting points

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