Unit 6 Introduction to Chemistry

1. Unit 6 Introduction to Chemistry Chapter 18 Bonding and Compounds

2. 18A – Principles of Bonding • Objectives: • Explain why atoms bond together • Differentiate between electron affinity and electronegativity and describe their general trends on the periodic table • Briefly describe covalent, ionic, and metallic bonds • Explain how relative electron affinities are important to the kind of bond formed • Assignments: Outline and Section Review page 444

3. Introduction to Chemistry • If it’s matter, it’s chemistry • Let’s Read… • Page 437

4. Bonding: Holding Atoms Together • Let’s Read… • Page 440

5. Octet Rule • The tendency for atoms to bond is governed by the 2nd law of thermodynamics • All natural processes move toward a state of minimum energy • Atoms are more unstable when they are not bonding to another atom • Atoms generally are most stable when they have a full eight electrons in their valence energy level • This principle is called the octet rule • Exceptions are hydrogen and helium • Two ways to bond • Sharing electrons – Covalent Bonding • Transferring electrons – Ionic Bonding

6. Electron Affinity & Electronegativity • Electron affinity • A measure of how well the atom attracts and holds electrons • Halogens and most nonmetals have the highest electron affinities • Most of the metals have low affinities • Noble gases have very low affinities • Linus Carl Pauling • Devised a scale to compare an atom’s ability to attract and hold electrons when bonded to other atoms • This is known as an element’s electronegativity • Electron affinity applies to unbonded atoms • Electronegativity applies to bonded atoms • Go together, large affinity means large electronegativity • Periodic Trend • Decreases going down a family • Increases across a period

7. Types of Bonds • Three Types • Covalent Bonds • Between two nonmetals • High electronegativity, so they SHARE electrons • Ionic Bonds • Between a nonmetal and a metal • One high electronegativity, one low • There is a TRANSFER of electrons • Metallic Bonds • Two metals • SEA of electrons

8. 18B – Covalent Bonds • Objectives: • Describe a covalent bond, including the kind of element that usually forms covalent bonds • Illustrate the pairing of electrons to form covalent bonds using electron dot notation • Draw Lewis structures of simple molecules • Compare and contrast single, double, and triple covalent bonds • Explain what make a bond polar and how it affects the properties of covalent compounds • Explain how bonding and nonbonding electrons affect a molecule’s shape • Assignments: Outline and Section Review page 451

9. Illustrating Covalent Bonding

10. Illustrating Covalent Bonding • All atoms want eight in the outer level • Share to make eight • The electrons that are shared are called the bonding pair • The ones that do not share are called the nonbonding pairs • Drawn in two dimensions through a Lewis Structure • Bonds can be represented by dots, but more often a single dash between the bonded atoms • When dashes are used, the nonbonding electrons are usually omitted • Let’s do chlorine…

11. Bonds in Diatomic Molecules • What are the diatomic molecules? • Seven of them • Oxygen, nitrogen, hydrogen, flourine, chlorine, bromine, iodine • Double bonds • Stronger than single bonds • Shorter that single bonds • Represented by two dashes • Share two pairs of electrons • Triple bonds • Stronger that double bonds • Shorter than double bonds • Represented by three dashes • Share three pairs of electrons • Let’s draw H2, O2, and N2

12. Bonds in Polyatomic Molecules • Polyatomic molecules • Molecules with three of more atoms • Electron dot notation for an element helps to determine how many covalent bonds can form • In general, for nonmetal elements, the number of unpaired dots is the number of covalent bonds the atom can form with the other atoms • These unpaired electrons are called bonding sites • How many does O, N, C, H have?

13. Covalent Bond Polarity • Covalent bonds share electrons, but do they share them equally? • It’s sometimes a tug of war! • Creates a polar bond • Uneven electron sharing between atoms with different electronegativity values • Diatomic elements form nonpolar bonds • No polarity, even sharing • Atoms of a polar bond acquire a partial electrical charge • Polar molecule • Intermolecular forces • Facet, page 448

14. Covalent Bonding & Molecular Structure • Let’s Read… • Page 448

15. Properties of Covalent Compounds • Covalent bonds make covalent compounds • Characteristics • Not as strong as ionic bonds • Relatively low melting and boiling points • Mostly gases and liquids at room temperatures • Some have great elasticity • Not many conduct electricity or heat which makes them excellent insulators • Some store tremendous energy, such as explosives • Our body is composed of covalent compounds • Why?

16. 18C – Ionic Bonds • Objectives: • Explain how an ionic bond forms • Tell which kinds of elements generally form ionic compounds • Use electron dot symbols to describe the net electron transfers in ionic bonds • Describe how ions assemble in crystal lattices • Identify the structural subdivisions that make up a crystal lattice • Describe the properties of ionic compounds • Assignments: Outline and Section Review page 456

17. Forming Ionic Bonds • Ionic bonds • A transfer of electrons • Opposite charges hold the atoms together • Ionic bonds makes an ionic compound • Usually between a nonmetal and a metal

18. Illustrating Ionic Bonds

19. Illustrating Ionic Bonds • Electron Dot Diagrams, not Lewis Structures • Na and F • Mg and S • Mg and F • Na and S

20. Structure of Ionic Compounds • Ionic compounds consist of innumerable ions • An orderly pattern appears as positive and negative ions line up • They eventually form an immense three dimensional pattern called a crystal lattice • Does not form molecules! • Instead makes a formula unit • The smallest ratio of elements that describes its chemical composition • Ionic crystal lattices are composed of repeating structural units called unit cells • Seven different shapes of unit cells, one is a cube

21. Properties of Ionic Compounds • At room temperature most are solid, brittle, crystalline substances • Beautiful crystals • High melting and boiling points due to the strength of the bonds • Many dissolve in water • Excellent conductors of electricity when dissolved or melted • Electrolytes • A substance that produces ions when it dissolves in a liquid

22. 18D – Metallic Bonds • Objectives: • Describe the simplest model of metallic bonding • Explain how metal atoms are arranged in elements and in compounds (alloys) • Explain how metallic bonds account for the properties of metals • Assignments: Outline and Section Review page 459

23. Metal Bonding Theory • Electron Sea Theory • Also called the free electron theory • Metals hold on to their valence electrons loosely which makes them able to be “shared” amongst a large number • Metallic bonds • Positive metal ions arranged in a crystal lattice immersed in a “sea” of negative electrons • These valance electrons are free to roam the entire lattice structure • All the nuclei share the former valence electrons • No electron dot notation can be used to represent metallic bonds • Compounds of different metals are called alloys

24. Metallic Lattice Structure • Similar to ionic structure • Most efficient layout for metallic structures is hexagonal close packed • Other types of structures • Face centered cubic • Body centered cubic

25. Properties of Metals • Most are solids at room temperatures • Malleable • Ability to be hammered or rolled into thin sheets • Ductile • The ability to be stretched without breaking or drawn into a wire • Shiny luster

26. Summary of Chemical Bonds