Section1 Introduction to Chemical Bonding Chapter 6 Objectives • Definechemical 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 An attractive force that holds atoms or ions together

Types of Chemical Bonds Ionic -electrical attraction between positive and negative ions -transfer of electrons -usually a metal and a nonmetal Cation- positive ion Anion- negative ion

Types of Chemical Bonds, cont’d Covalent -sharing of electrons between two atoms to make a bond -can be: nonpolar covalent (equal sharing) polar covalent (unequal sharing)

Types of Chemical Bonds, cont’d Determining the Type of Bond -bonds are not 100% ionic or covalent -use the table of electronegativities (pg.159) 3.3 – 1.8 ionic bonds 1.7- 0.3 polar covalent 0.2- 0 nonpolar covalent

Section1 Introduction to Chemical Bonding Chapter 6 Chemical Bonding, continued Sample Problem A Use electronegativity values listed on page 159, and Figure 2 in your book, on page 172, to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative?

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Objectives • Definemolecule and molecular formula. • Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy. • State the octet rule.

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Molecular Compounds • Amolecule is a neutral group of atoms that are held together by covalent bonds. • A chemical compound whose simplest units are molecules is called amolecular compound.

Visual Concepts Chapter 6 Molecule

Molecular Compounds One type of molecular compound is called a Diatomic molecule examples- O2, this is composed of two atoms of oxygen others:

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Formation of a Covalent Bond • Atoms have lower potential energy when they are bonded to other atoms • The figure below shows potential energy changes during the formation of a hydrogen-hydrogen bond.

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Characteristics of the Covalent Bond • The distance between two bonded atoms at their most stable energy state is the bond length. • In forming bond, atoms release energy. The same amount of energy must be added to separate the bonded atoms. This is called Bond energy .

Bond Energies & Bond Lengths for Single Bonds- Inverse relationship Section2 Covalent Bonding and Molecular Compounds Chapter 6

Characteristics of the Covalent Bond When two atoms form a covalent bond, their outer orbitals overlap. This achieves a noble-gas configuration.

Section2 Covalent Bonding and Molecular Compounds Chapter 6 The Octet Rule • Bond formation follows theoctet rule: • Compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest energy level. • Become like noble gases.

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Exceptions to the Octet Rule • Atoms have more or less than 8 electron in outer energy level • Hydrogen- only two electrons. • Boron– only six electrons. • Main-group elements in Periods 3-7 can form bonds withexpanded valence,involving more than eight electrons.

Electron-Dot Notation Bonding involves the outer s & p electrons (valence electrons). Using a dot notation shows the outer electrons each atom has for bonding.

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Electron-Dot Notation, continued Sample Problem B • Write the electron-dot notation for hydrogen. b. Write the electron-dot notation for nitrogen.

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Objectives, continued • Use the basic steps to write Lewis structures. • Determine Lewis structures for molecules containing single bonds, multiple bonds, or both. • Explain why scientists use resonance structures to represent some molecules.

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Lewis Structures • Electron-dot notation can also be used to represent molecules. • The pair of dots between the two symbols represents the shared electron pair

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Lewis Structures • The pair of dots between the two symbols represents the shared pair of a covalent bond. • In addition, each atom is surrounded by three pairs of electrons that are not shared in bonds. • An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Lewis Structures • The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash. • example:

Rules for Drawing Lewis Structures • Determine the total valence electrons for all atoms. • Draw a skeleton. Central atom is C or least electronegative. Never H. • Give each atom an octet. (some exceptions) • Double or triple bond if you don’t have enough electrons. (COPSN) • Extra electrons? Put on the central atom.

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Lewis Structures, continued Examples: Draw the Lewis structure of a. iodomethane, CH3I. b. Ammonia, NH3 c. Silane, SiH4 d. Phosphorous trifluoride, PF3

Section2 Covalent Bonding and Molecular Compounds Chapter 6 Multiple Covalent Bonds • A double bond- two pairs of electrons are shared between two atoms. • A triple bond-three pairs of electrons are shared between two atoms. • In general, double bonds have greater bond energies and are shorter than single bonds. • Triple bonds are even stronger and shorter than double bonds.

Lewis Structures with Multiple Bonds Examples: • Formaldehyde, CH2O • Carbon dioxide, CO2 • Hydrogen cyanide, HCN • Ethyne, C2H2

Section3 Ionic Bonding and Ionic Compounds Chapter 6 Polyatomic Ions • A charged group of covalently bonded atoms is known as a polyatomic ion. • Like other ions, polyatomic ions have a charge that results from either a shortage or excess of electrons.

Section3 Ionic Bonding and Ionic Compounds Chapter 6 Polyatomic Ions, continued • Some examples of Lewis structures of polyatomic ions • nitrate NO3-1 • ammonium, NH4+1 • sulfate, SO4-2

Section3 Ionic Bonding and Ionic Compounds Chapter 6 Objectives • Compare a chemical formula for a molecular compounds with one for an ionic compound. • Discuss the arrangements of ions in crystals. • Definelattice 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.

Section3 Ionic Bonding and Ionic Compounds Chapter 6 Ionic Compounds • Most of the rocks and minerals that make up Earth’s crust consist of positive and negative ions held together by ionic bonding. • example:table salt, NaCl, • Anionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal.

Section3 Ionic Bonding and Ionic Compounds Chapter 6 Ionic Compounds • Most ionic compounds exist as crystalline solids. • In contrast to a molecular compound, an ionic compound is not composed of independent, neutral units that can be isolated.

Section3 Ionic Bonding and Ionic Compounds Chapter 6 Ionic Compounds, continued • The chemical formula of an ionic compound represents the simplest ratio of the compound’s ions. • A formula unitis the simplest collection of atoms from which an ionic compound’s formula can be established.

Section3 Ionic Bonding and Ionic Compounds Chapter 6 Ionic Vs. Covalent Bonding

Section3 Ionic Bonding and Ionic Compounds Chapter 6 Formation of Ionic Compounds

Section3 Ionic Bonding and Ionic Compounds Chapter 6 Formation of Ionic Compounds, continued • In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice. • The combined attractive and repulsive forces within a crystal lattice determine: • the distances between ions. • the pattern of the ions’ arrangement in the crystal

Section3 Ionic Bonding and Ionic Compounds Chapter 6 NaCl and CsCl Crystal Lattices

Section4 Metallic Bonding Chapter 6 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.

Section4 Metallic Bonding Chapter 6 Metallic Bonding • Chemical bonding is different in metals than it is in ionic, molecular, or covalent-network compounds. • The unique characteristics of metallic bonding gives metals their characteristic properties, listed below. • electrical conductivity • thermal conductivity • malleability • ductility • shiny appearance

Section4 Metallic Bonding Chapter 6 Properties of Substances with Metallic, Ionic, and Covalent Bonds

Section4 Metallic Bonding Chapter 6 The Metallic-Bond Model • In a metal, the vacant orbitals in the atoms’ outer energy levels overlap. • This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal. • The electrons are delocalized,which means that they do not belong to any one atom but move freely about the metal’s network of empty atomic orbitals. • These mobile electrons form a sea of electronsaround the metal atoms, which are packed togetherin a crystal lattice.

Section4 Metallic Bonding Chapter 6 Metallic Bonding, continued • Mobile electrons allow charge/heat to pass and allow the metal to be hammered because there is not a rigid structure.

Section5 Molecular Geometry Chapter 6 Objectives • Explain VSEPR theory. • Predict the shapes of molecules or polyatomic ions using VSEPR theory. • Explain the shapes of molecules or polyatomic ions using VSEPR theory.

Section5 Molecular Geometry Chapter 6 Molecular Geometry • The properties of molecules depend not only on the bonding of atoms but also on molecular geometry:the three-dimensional arrangement of a molecule’s atoms.

Section5 Molecular Geometry Chapter 6 VSEPR Theory • As shown at right, diatomic molecules, like those of (a)hydrogen, H2, and (b)hydrogen chloride, HCl, can only be linear because they consist of only two atoms. • To predict the geometries of more-complicated molecules, one must consider the locations of all electron pairs surrounding the bonding atoms. This is the basis of VSEPR theory.

Section5 Molecular Geometry Chapter 6 VSEPR Theory • VSEPR stands for • “valence-shell electron-pair repulsion.” • VSEPR theorystates that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible.

Section5 Molecular Geometry Chapter 6 VSEPR Theory • Find the VSEPR formula • A=central atom • B=bonded atoms • E=lone pairs of e- on the central atom • Use the chart on pg.200 to determine the molecular shape

Section5 Molecular Geometry Chapter 6 VSEPR Theory, continued Sample Problems Use VSEPR theory to predict the molecular geometry of a. BCl3 b. NH3 c. CH4 d. NO3-1 e. CO2

Section5 Molecular Geometry Chapter 6 VSEPR Theory, continued • Unshared electron pairs repel other electron pairs more strongly than bonding pairs do. • This is why the bond angles in ammonia and water are somewhat less than the 109.5° bond angles of a perfectly tetrahedral molecule.

Section5 Molecular Geometry Chapter 6 VSEPR Theory, continued • Treat double and triple bonds the same way as single bonds. • Treat polyatomic ions similarly to molecules. • The next slide shows several more examples of molecular geometries determined by VSEPR theory.

Table of Contents