2. A. Ionic Bonding 1. attraction between large numbers of (+) ions and (-) ions 2. results when there is large electronegativity differences 3. generally involves a metal and a nonmetal

3. B. Covalent Bonding 1. sharing of electron pairs between two atoms 2. results when small electronegativity difference 3. generally involves two nonmetal atoms 4. nonpolar covalent bond a. Electrons are shared equally b. Same two nonmetals or very small electronegativity difference 5. polar covalent bond a. The more electronegative atom attracts the shared electrons more strongly b. Results in partial charges on the atoms

4. C. Greater electronegativity difference (see p.151) results in a more polar bond or greater ionic character

5. A. Molecule – a neutral group of nonmetal atoms joined together by covalent bonds B. Molecular Compound 1. generally contains only nonmetals 2. has covalent bonds 3. particles are called molecules 4. represented by a molecular formula which gives the actual number of atoms or each element

6. C. The formation of a covalent bond results from the attraction of the shared electrons by nuclei of both atoms D. Characteristics of a covalent bond 1. bond length – the distance between the nuclei of the two atoms 2. bond energy – the energy needed to break a bond

7. E. Octet Rule 1. atoms lose, gain, or share electrons to have 8 in the outer level; (except H=2) 2. most compounds follow the octet rule 3. exceptions to the octet rule a. Odd number of valence electrons for the compound b. BF 3 (also some other halides of B or Be) c. Halogen atoms attached to P or S; more than 8 electrons around the central atom

8. F. Electron Dot Symbol a. Use dots to represent the valence electrons b. KMgAlSi NSBrAr

9. G. Lewis Structure (electron dot structures) can be used to show how atoms share electrons in a compound H. Single, Double, and Triple Covalent Bonds a. single bond – sharing one pair of electrons b. double bond – sharing 2 pairs of electrons c. Triple bond – sharing 3 pairs of electrons

10. I. Rules for Lewis Structures 1. count the total number of valence electrons (write them down) 2. determine the central atom (usually the element closest to the center of periodic table) 3. draw skeleton structure using only single bonds 4. distribute the remaining electrons to have 8 around each atom (except H=2) 5. carbon always has 4 bonds (except when bonded to only 1 atom) 6. only one bond to F, Cl, Br, I

11. J. Draw Lewis Structure watermethane ammoniacarbon dioxide

12. K. Resonance Structures – when more than one Lewis structure can be written for the molecule 1. ozone (O 3 ) 2. sulfur trioxide (SO 3 )

13. A. Ionic Compounds a. Generally contain a metal and a nonmetal b. (+) and (-) ions combined so that charges are equal c. Most form crystalline solids d. Represented by a formula unit; simplest ratio of the ions

14. B. Electrons Dot Symbols and Ionic Bonding 1. sodium and fluorine 2. magnesium and chlorine

15. C. Ionic Compounds Metal and nonmetal Ionic bonds High melting points All are solids Hard, but brittle Nonconductors as solids Conduct when melted or dissolved Molecular Compounds Contain only nonmetals Covalent bonds Low melting and boiling points Solids, liquids, and gases Nonconductors

16. D. Polyatomic Ion – a group of covalently bonded atoms with a (+) or (-) charge Draw Lewis Structures ammonium ionsulfate ion

17. A. Outer electrons of a metal move freely from atom to atom B. Metallic bonding is the result of the attraction between metal atoms and the surrounding “sea of electrons” C. Metallic Properties 1. good conductors of electricity and heat 2. shiny appearance 3. malleable – can be hammered into different shapes 4. ductile – can be drawn into a wire D. Higher heat of vaporization results from stronger metallic bonding

18. A. VSEPR Theory (Valence Shell Electron Pair Repulsion) 1. electron pairs repel and are as far apart as possible 2. determines the shape of the molecule (molecular geometry)

19. B. Atoms Attached Unshared Pairs Bond Shape to Central Atom On Central AtomAngles 40109.5 tetrahedral 31109.5 pyramid 30120 triangle 22109.5 bent 21120 bent 20180 linear

20. C. Draw Lewis Structure and determine bond angle and shape methaneammonia sulfur trioxidewater sulfur dioxidecarbon dioxide

21. D. Hybridization 1. s and p atomic orbitals mix to form hybrid orbitals 2. Type of Hybrid OrbitalFound if atom has sp 3 only single bonds sp 2 one double bond spone triple bond (or two double)

22. D. Hybridization 3. Show type of hybridization for each carbon atom

23. E. Polar Molecule (dipole) 1. lines up a certain way in an electric field 2. molecule with one end slightly positive and the other end slightly negative 3. hydrochloric acidcarbon dioxide watercarbon tetrachloride chloroform

24. F. Intermolecular Forces (attraction between molecules) 1. Stronger attraction between molecules results in higher melting/boiling point; determines whether the substance is solid, liquid or gas 2. Types of Intermolecular Forces a. London Dispersion Forces (weakest) b. Dipole Forces c. Hydrogen Bonding (strongest)

25. G. London Dispersion Forces 1. Weakest intermolecular force 2. Caused by the motion of electrons; having a majority of electron on one side of molecule causes a temporary unbalanced charge 3. All molecules have dispersion forces, but they are important only if it does not have dipole forces or hydrogen bonding 4. Strength increases with greater molar mass

26. H. Dipole Forces 1. attraction between polar molecules 2. H – Cl --------H – Cl

27. I. Hydrogen Bonding 1. strongest intermolecular force 2. hydrogen attached to N, O, F is also attracted to an unshared pair on a nearby molecule 3. water has hydrogen bonding