Chapter 6 Chemical Bonding Table of Contents Section

Chapter 6 Chemical Bonding Table of Contents Section 1 Introduction to Chemical Bonding Section 2 Covalent Bonding and Molecular Compounds Section 3 Ionic Bonding and Ionic Compounds Section 4 Metallic Bonding Section 5 Molecular Geometry Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chemical Bond An attractive force that holds atoms or ions together Chapter menu Resources

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 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

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) Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

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 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 1 Introduction to 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? Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

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 menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds • A molecule is a neutral group of atoms that are held together by covalent bonds. • A chemical compound whose simplest units are molecules is called a molecular compound. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Visual Concepts Molecule Chapter menu Resources Copyright © by Holt, Rinehart and

Molecular Compounds One type of molecular compound is called a Diatomic molecule examples- O 2, this is composed of two atoms of oxygen others: Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds Bond Energies & Bond Lengths for Single Bonds- Inverse relationship Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Characteristics of the Covalent Bond When two atoms form a covalent bond, their outer orbitals overlap. This achieves a noble-gas configuration. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds The Octet Rule • Bond formation follows the octet 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds 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 with expanded valence, involving more than eight electrons. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds Electron-Dot Notation, continued Sample Problem B a. Write the electron-dot notation for hydrogen. b. Write the electron-dot notation for nitrogen. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds 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 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds Lewis Structures • The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash. • example: Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Rules for Drawing Lewis Structures 1. Determine the total valence electrons for all atoms. 2. Draw a skeleton. Central atom is C or least electronegative. Never H. 3. Give each atom an octet. (some exceptions) 4. Double or triple bond if you don’t have enough electrons. (COPSN) 5. Extra electrons? Put on the central atom. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds Lewis Structures, continued Examples: Draw the Lewis structure of a. iodomethane, CH 3 I. b. Ammonia, NH 3 c. Silane, Si. H 4 d. Phosphorous trifluoride, PF 3 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 2 Covalent Bonding and Molecular Compounds 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Lewis Structures with Multiple Bonds Examples: a. b. c. d. Formaldehyde, CH 2 O Carbon dioxide, CO 2 Hydrogen cyanide, HCN Ethyne, C 2 H 2 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 3 Ionic Bonding and Ionic Compounds 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 3 Ionic Bonding and Ionic Compounds Polyatomic Ions, continued • Some examples of Lewis structures of polyatomic ions • nitrate NO 3 -1 • ammonium, NH 4 +1 • sulfate, SO 4 -2 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 3 Ionic Bonding and 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, Na. Cl, • An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 3 Ionic Bonding and 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 3 Ionic Bonding and Ionic Compounds, continued • The chemical formula of an ionic compound represents the simplest ratio of the compound’s ions. • A formula unit is the simplest collection of atoms from which an ionic compound’s formula can be established. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 3 Ionic Bonding and Ionic Compounds Ionic Vs. Covalent Bonding Chapter

Chapter 6 Section 3 Ionic Bonding and Ionic Compounds Formation of Ionic Compounds Chapter

Chapter 6 Section 3 Ionic Bonding and Ionic Compounds 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 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 3 Ionic Bonding and Ionic Compounds Na. Cl and Cs. Cl

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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 4 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 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 4 Metallic Bonding Properties of Substances with Metallic, Ionic, and Covalent

Chapter 6 Section 4 Metallic Bonding 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 electrons around the metal atoms, which are packed together in a crystal lattice. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 4 Metallic Bonding, continued • Mobile electrons allow charge/heat to pass and allow the metal to be hammered because there is not a rigid structure. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry VSEPR Theory • As shown at right, diatomic molecules, like those of (a) hydrogen, H 2, 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry VSEPR Theory • VSEPR stands for • “valence-shell electron-pair repulsion. ” • VSEPR theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry 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 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry VSEPR Theory, continued Sample Problems Use VSEPR theory to predict the molecular geometry of a. BCl 3 b. NH 3 c. CH 4 d. NO 3 -1 e. CO 2 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry 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. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry VSEPR and Molecular Geometry Chapter menu Resources Copyright

Chapter 6 Section 5 Molecular Geometry Intermolecular Forces a. k. a. van der Waals Forces • The forces of attraction between molecules are known as intermolecular forces. • The boiling point is a good measure of the intermolecular forces between its molecules: • the higher the boiling point, the stronger the forces between the molecules. Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry Intermolecular Forces, continued • The strongest intermolecular forces exist between polar molecules. • Identical atoms attached = nonpolar • Different atoms attached = polar • Polar bonds are the result of an asymmetrical distribution of electrons- causes a dipole • Nonpolar bonds- equal distribution of electrons Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry Intermolecular Forces, continued • A dipole is represented by an arrow with its head pointing toward the negative pole and a crossed tail at the positive pole. The dipole created by a hydrogen chloride molecule is indicated as follows: Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Chapter 6 Section 5 Molecular Geometry Intermolecular Forces, continued • Determine if each of the molecules is polar or nonpolar: • CH 4 • Cl 2 • CCl 3 H Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Intermolecular Forces a. k. a-van der Waals forces Three types 1. Hydrogen Bonding 2. Dipole-Dipole 3. London Dispersion Forces Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Hydrogen Bonding -Strongest -occurs between molecules that contain O-H, N-H or H-F atoms bonding. -example: H 2 O vs. H 2 S Boiling point 100 C -61 C Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Dipole-Dipole -2 nd in strength - occurs in polar molecules - will see an increase in boiling point as the mass increases Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

London Dispersion Forces -3 rd in strength -created from “momentary” dipoles that occur between two nonpolar molecules -also-increase in strength with increase in mass Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.

Intermolecular Forces Let’s try some: 1. Draw Lewis Structure 2. Determine if the molecule is polar or nonpolar 3. Pick based on polarity or if you see an HF, NH or OH you will know. 1. HCl 2. NH 3 3. CH 3 Cl 4. Si. F 4 Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.