Chapter 6 Review

1. Chapter 6 Review Test: Friday December 09, 2011

2. Be sure to know the following terms • Chemical bond • Nonpolar covalent bond • Polar covalent bond • Chemical formula • Unshared (lone) pair • Resonance • Polyatomic ions

3. What is the most electronegative element on the periodic table? • Fluorine (F) • What is the least electronegative element on the periodic table? • Cesium (Cs)

4. According to the rules of drawing Lewis structures, which atom is the central atom? • The least electronegative atom

5. Use electronegativity differences to determine whether each of the following bonds is nonpolar covalent, polar covalent, or ionic. Determine the direction of polarity. H—F Na—Cl 4.0 – 2.1 =1.9 3.0 – 0.9 = 2.1 H—O N—N 3.5 – 2.1 = 1.4 3.0 – 3.0 = 0 H—C H—N 2.5 – 2.1 = 0.4 3.0 – 2.1 = 0.9 Ba—O C—O 3.5 – 0.9 = 2.6 3.5 – 2.5 = 1.0

6. Define single bond, double bond, and triple bond. List them in order of increasing bond energy…Increasing bond length. • Single bond: bond in which one pair of electrons is shared (C--C) • Double bond: bond in which two pairs of electrons are shared (C=C) • Triple bond: bond in which three pairs of electrons are shared • BE single > double > triple • BL triple > double > single

7. Compare and contrast ionic and covalent compounds.

8. The manner in which ions lower their potential energy and stabilizing ionic compounds by combining in an orderly arrangement is called…. • Crystal lattice

9. What is the octet rule? • Outer energy level with 8 electrons • Name exceptions to the octet rule. • Hydrogen, boron • Oxygen, fluorine, chlorine • What is meant by expanded valence? Which elements can have an expanded valence? • An outer energy level with more than 8 electrons • The most electronegative elements Cl, F, O

10. What is the amount of energy required to separate bonded atoms into neutral atoms? • Bond energy • What is the amount of energy required to separate an ionic compound into its constituent ions? • Lattice energy

11. Metallic Bonding • The highest energy levels of most metal atoms are occupied by very few electrons. For example s-block metals contain only 1 or 2 electrons and all three p orbitals are empty. • Within 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 do not belong to any one atom but move freely about the metal’s network of empty atomic orbitals (sea of electrons)

12. Metallic Bonding (cont’d) * The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons The freedom of electrons to move in a network accounts for the properties of metals: electrical and thermal conductivity, absorption/reflection of light, luster, malleability, ductility

14. Enthalpy of vaporization • Enthalpy = heat • Just as lattice energy represents the strength of an ionic bond, enthalpy of vaporization represents the strength of a metallic bond *amount of energy (as heat) required to vaporize the metal

15. Intermolecular Forces(forces of attraction between molecules) -vary in strength, but generally weaker than covalent, metallic, and ionic bonds • dipole: created by equal but opposite charges separated by a short distance -occurs when a slightly negatively charged atom in a polar bond is attracted to a slightly negatively charged atom in a nearby molecule

16. Intermolecular Forces (cont’d) 2. Hydrogen bonding: intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule

17. Intermolecular Forces (cont’d) 3. London dispersion forces: intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles *act between ALL molecules and atoms *ONLY intermolecular forces acting among noble gases and nonpolar molecules * Strength of LDFs increasing with increasing atomic or molar mass (more specifically—increasing # of electrons)

18. Arrange the following attractions in order of increasing strength: covalent, ionic, metallic, dipole-dipole, hydrogen bonding, London dispersion • London dispersion • Dipole-dipole • Hydrogen bonding • Nonpolar covalent • Polar covalent • Ionic • Metallic

21. Which of the following pairs of elements do you expect to be most similar? Why? (a) Ti and Ga (b) N and O (c) Li and Na (d) Ar and Br (e) Ge and Ga

22. Predict the ion formed by each of the following. How many electrons will be gained OR lost? • Rb • K • Al • O • F • N • Mg • C • B

23. Classify the following elements as atomic or molecular. • Xenon • Iodine • Oxygen • Nickel • Classify each of the following compounds as ionic or molecular. • CS2 CuO • KI PCl3 • PtO2 CO • SO3 H2O

24. What are the basic units—atoms, molecules, or formula units—that compose each of the following substances? BaBr2 Rb2O NO Ne CF4 N2 I2 N2F4 O3 CO NaClHBr

25. Use electron-dot notation to illustrate the number of valence electrons present in one atom of each of the following elements. Rb C O P B Cl Li N Te ArCa As GaGe At

27. In writing Lewis structures, how is the need for multiple bonds generally determined? • The need for a multiple bond becomes obvious if there are not enough valence electrons to complete octets by adding unshared pairs. • Multiple bonds are possible in Lewis structures containing C, N, or O.

28. What is a formula unit? What are the components of one formula unit of CaF2? • Smallest whole unit of an ionic compound • Ca2+ and 2 F-

29. What type of energy best represents the strength of an ionic bond? What type of bonding holds a polyatomic ion together? How does bond length relate to bond energy? What is the relationship of heat of vaporization of a metal and the strength of the bonds that hold the metal together?

30. A chemical bond between atoms results from the attraction between the valence electrons of one atom and the ________ in the _________ of another atoms. • Protons; nucleus • Atoms with a strong attraction for electrons they share with another atom exhibit ________. • High electronegativity • Bonds that possess between 5% and 50% ionic character are considered to be _________. • Polar covalent • The greater the electronegativity difference between two atoms bonded together, the greater the bond’s percentage of _________. • Ionic character

31. The electrons involved in the formation of a chemical bond are called ____________. If a bond’s character is more than 50% ionic, then the bond is called a(n) ________________. How can electronegativity be used to distinguish between an ionic and covalent bond? In a crystal of an ionic compound, each cation is surrounded by a number of _________. Compared with the neutral atoms involved in the formation of an ionic compound, the crystal lattice that results is _________ in potential energy.

32. VSEPR valence shell electron-pair repulsion • Identify the major assumption of the VSEPR • theory that is used to predict the shape of • atoms. • Electrons are located as far apart as possible because of electron-electron repulsion • In the VSEPR theory, double and triple bonds are treated the same as single bonds.

33. BeF2 BF3 ONF CH4 NH3 H2O AB2 AB2E AB3 AB2E2 AB4 AB3E