Chapter 6 CHEMICAL BONDING

WHAT IS ELECTRONEGATIVITY? WHY DOES IT MATTER?

Section 5.1 INTRODUCTION TO CHEMICAL BONDING

 Electronegativity- The ability for an atom to attract an electron to itself  There are 3 types of bonds  Ionic Bond  Covalent Bond  Polar Covalent Bond SECTION 6.1

 Ionic bond- Electrons are transferred forming ions. These ions are attracted to each other. The bond is a strong force between the two atoms.  Covalent bond- two or more atoms sharing pairs of electrons TYPES OF BONDS

 Polar Covalent- two or more atoms sharing a pair of electrons unevenly TYPES OF BONDS (CONT)

 The type of bond between two or more atoms depends on the difference in electronegativity of the atoms. ELECTRONEGATIVITY AND BONDING Type of BondDifference in Electronegativity Ionic Polar- Covalent Covalent

 What is the main distinction between ionic and covalent bonding?  How is electronegativity used in determining the ionic or covalent character of the bonding between two elements?  What type of bond would be expected between the following atoms?  Li and F  Cu and S  I and Bf SECTION 6.1: LEARNING CHECK

Section 6.2 COVALENT BONDING AND MOLECULAR COMPOUNDS

 Molecule- a neutral group of atoms that are held together by covalent bonds  Chemical formula- the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts  Octet- 8 electrons  Single bond- 1 electron pair shared  Double bond- TERMS TO KNOW

EXAMINE THESE COVALENT MOLECULES. WHAT DO YOU NOTICE?

 Overall goal:  Obtain an octet  Create a stable atom  Lowest amount to energy FORMATION OF A COVALENT BOND

 Bond length: distance between two bonded atoms  Bond energy: the energy required to break a chemical bond and form neutral isolated atoms. CHARACTERISTICS OF THE COVALENT BOND

 Atoms will gain, lose, or share electrons to get 8 electrons in their highest occupied energy level  Exceptions:  Helium and Hydrogen- only need 2 valence electrons  Boron- has 3 electrons and is stable with 6 electrons  Example: BF 3  Expanded Octet- Some atoms can hold more than 8 when they are bonded to extremely electronegative atoms.  Example: PF 5 and SF 6 THE OCTET RULE

 An electron configuration notation in which only the valence electrons of an atom of a particular element are shown. ELECTRON-DOT NOTATION

 Visual representation of molecules  Element symbol- nuclei and inner-shell electrons  Dot-pairs- non-bonding valence electrons  Dashes- bond between two elements LEWIS STRUCTURES

1.Determine the type and number of atoms in the molecule 2.Write the electron-dot notation for each type of atom in the molecule 3.Determine the total number of valence electrons available to be combined 4.Arrange the atoms 1.If carbon is present it is in the center 2.If not, the least electronegative atom is in the center 5.Add unshared pairs of electrons to each nonmetal atom (except hydrogen) Each atom needs to be surrounded by 8 electrons. 6.Count the electrons in the structure to be sure that the number of valence electrons used equals the number available. 7.Place non-bonding valence electrons in pairs around atoms without an octet HOW TO DRAW LEWIS STRUCTURES?

 NH 3  H 2 S  SiH 4  PF 3 EXAMPLES:

 Indicates the kind, number, arrangement, and bonds but not the unshared electrons. STRUCTURAL FORMULA H-Cl

 Double or triple bonds  Triple bonds are the shortest and strongest covalent bonds  The need for a multiple bond becomes obvious if there are not enough valence electrons to complete octets by adding unshared pairs. MULTIPLE COVALENT BONDS

 Compare the molecules H 2 NNH 2 and HNNH REVIEW

 Bonding in molecules or ions that cannot be correctly represented by a single Lewis structure. RESONANCE STRUCTURES