Chapter 6 Chemical Bonding Chapter 6 Section 1

Chapter 6 Chemical Bonding

Chapter 6 Section 1 Introduction to Chemical Bonding Objectives • Define chemical bond. • Explain why most atoms form chemical bonds. • Describe ionic and covalent bonding. • Explain why most chemical bonding is neither purely ionic nor purely covalent. • Classify bonding type according to electronegativity differences.

Chemical Bond – atoms bound together by a mutual electrical attraction between nuclei and valence electrons • Ionic Bond – bond between anions (-) and cations (+) caused by the transfer of electrons between atoms – usually a metal and a nonmetal • Covalent Bond – bond between atoms that share valence electrons – usually 2 nonmetals • Metallic Bonding – attraction between metals with delocalized electrons – (not a real bond) • Hydrogen Bond – attraction that occurs between hydrogen atoms in a compound and strongly electronegative atoms with lone pairs of electrons (oxygen, nitrogen, etc. ) (not a real bond) Section 1 Introduction to Chemical Bonding

Predicting Bond Type by Electronegativity Difference

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds Objectives • Define molecule and molecular formula. • Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy. • State the octet rule.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds Objectives, continued • List the six basic steps used in writing Lewis structures. • Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both. • Explain why scientists use resonance structures to represent some molecules.

Polar Covalent Bonding results in an unequal sharing or uneven electron distribution Section 2 Covalent Bonding and Molecular Compounds

• Molecule – group of atoms that are covalently bonded • Compound – group of atoms that are ionicly bonded • Chemical Formula – tells the kind and number of atoms in a compound • Molecular Formula – tells the kind and. number of atoms in a molecule

Covalent Bond Characteristics • Bond Length – average distance between two bonded atoms at their minimum potential energy • Bond Angle – angle between three bonded atoms • Bond Energy – energy required to break a chemical bond and form neutral isolated atoms

Lewis Structures • Structural Formula – shows the kind, number and arrangement of atoms in a molecule (may or may not show lone pairs) • Octet Rule – compounds tend to form so that both atoms have eight electrons in their outer energy level – Exceptions – hydrogen & expanded octets • Electron Dot Notation – Lewis Dots – shows valence electrons

Rules for Drawing Lewis Structures 1. Determine the type and number of atoms 2. Determine the total number of valence electrons available to bond 3. Draw skeletal structure (least electronegative atom is usually central) 4. Draw bonds 5. Place extra electrons to fill octet 6. Check or recount electrons to make sure everything is “HAPPY”

Simple Structures

More Single Bonds

Multiple Covalent Bonds

Resonant Structures – two or more possible structures for the same molecule Ozone Nitrate Ion

Chapter 6 Section 3 Ionic Bonding and Ionic Compounds Objectives • Compare a chemical formula for a molecular compounds with one for an ionic compound. • Discuss the arrangements of ions in crystals. • Define lattice energy and explain its significance. • List and compare the distinctive properties of ionic and molecular compounds. • Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information. (we already worked on this one)

• Ionic Compounds – positive and negative ions combined in such a way that the net charge is 0 • Formula Unit – simplest collection (ratio) of atoms in an ionic compound Section 3 Ionic Bonding and Ionic Compounds

I o n i c B o n d s

Crystal Lattice – orderly arrangement of an ionic crystal

Basic Lattice Systems

Properties of Bond Types • Covalent Molecules – Low melting points – Low boiling points – insulators • Ionic Compounds – High melting points – High boiling points – Conductors when molten

Polyatomic Ions – a group of covalently bonded atoms that has a charge nitrate ion, NO 3 -

Chapter 6 Section 4 Metallic Bonding Objectives • Describe the electron-sea model of metallic bonding, and explain why metals are good electrical conductors. • Explain why metal surfaces are shiny. • Explain why metals are malleable and ductile but ionic-crystalline compound are not.

• Properties – Good conductors of heat and electricity – Good reflectors (shiny) – Malleable – Ductile – High tensile strength Section 4 Metallic Bonding – attraction between metals with delocalized electrons – (not a real bond)

Chapter 6 Section 5 Molecular Geometry Objectives • Explain VSEPR theory. • Predict the shapes of molecules or polyatomic ions using VSEPR theory. • Explain how the shapes of molecules are accounted for by hybridization theory.

Chapter 6 Section 5 Molecular Geometry Objectives, continued • Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces and their effects on properties such as boiling and melting points. • Explain the shapes of molecules or polyatomic ions using VSEPR theory.

Molecular Geometry • – Valence Shell Electron Pair Repulsion Theory – Predicts shapes of molecules based on the repulsion of unshared pairs of electrons (lone pairs) Least Repulsion Bonded Pair – Unbonded Pair Most Repulsion Unbonded Pair – Unbonded Pair Section 5 Molecular Geometry

Examples

Examples with Bond Angles

– mixing of two or more orbitals of similar energies to make orbitals with the same energy Hybridization of Carbon

– molecules that have a positive and negative ends due to the arrangement of bonds and lone pairs H F

force s of attraction between molecules that are weaker than bonds attraction between polar molecules attraction that occurs between hydrogen atoms in a compound and strongly electronegative atoms with lone pairs of electrons (oxygen, nitrogen, etc. ) very small, short lived attractions and repulsions caused by the motion of electrons

Dipole -Dipole Hydrogen Bonding London Dispersion Forces

Homework • Page 209 -211 • Numbers 3, 5, 6, 15, 16, 20, 21, 23, 24, 25, 28, 29, 31, 38, 46 48, 49