Chemistry 102 – Chapter 9 Summary and Review 1. Explain the correlation between valence electrons and group number. Answer: For the main group or representative elements, the group number is equal to the number of valence electrons. 2. What are the formal charges for the nitrate ion? How are these different from the oxidation numbers? Answer: Formal charges assume 100% sharing of electron density in a bond (100% covalent, or nonpolar covalent bond) while oxidation number is determined by assuming 100% ionic character of the bond (where the more electronegative atom “gets” the electrons). 3. Does breaking a chemical bond release energy? Answer: Breaking chemical bonds does not release energy but requires energy or an endothermic process. 4. The lattice energy for magnesium oxide is 3890 kJ∙mol–1. What is the chemical equation that describes this? Answer: Lattice energy is defined as the energy required to completely separate one mole of a solid ionic compound into its gaseous ions. Therefore, this will be a positive value (and endothermic process). For magnesium oxide, the equation is: MgO(s) → Mg2+(g) + O2–(g) 5. How many resonance structures are there for nitrous oxide (N2O)? Using formal charges, which structure is most reasonable? Answer: There are three resonance structures for nitrous oxide. The first structure is most valid since O is more electronegative than N. -1
O
:N
-2
:
N
+1
:
:
N
-1
:
:
O:
:
N
+1
:
:N
:
N
+1
O:
+1
6. What is the enthalpy of combustion of 1 mol of hydrogen gas (forming water vapor) using the bond enthalpies in the table? Bond Enthalpy, kJ∙mol–1 H–H 436.4 Answer: The equation is: 1H2 + ½O2 → H2O O=O 498.7 Or: H–H + ½O=O → H–O–H O–H 460 Breaking 1 mol (H–H) and ½ mol O=O and making 2 mol (O–H) ∆Horxn = 1 mol (436.4 kJ∙mol–1) + 0.5 mol (498.7 kJ∙mol–1) – 2 mol (460 kJ∙mol–1) = –234.3 kJ 7. With the Lewis dot structures of carbon dioxide and carbon monoxide, explain the oxidation numbers of carbon and oxygen using electronegativity. Answer: The Lewis dot structures for CO and CO2 are shown. Carbon is less electronegative than oxygen, so when it is assumed that it is behaving totally ionic (the premise for determining oxidation numbers), the more electronegative atom will get all of the electrons. CO2: Because the valence on oxygen is 6, with two extra electrons, the oxidation number is –2. Carbon has a valence of 4 but has “lost” all of these and now has an oxidation number of +4. CO: Again, oxygen has a valence of 6 and gains two electrons – giving it an oxidation number of –2. Carbon has a valence of 4 and loses two electrons, giving an oxidation number of +2. 8. What are the thermochemical equations for the bond enthalpies of nitrogen, oxygen and chlorine? What is the correlation between bond length, number of electrons in a bond, and bond strength? Answer: N ≡ N(g) → 2N(g) ∆H = 941.4 kJ∙mol–1 ∆H = 498.7 kJ∙mol–1 O = O(g) → 2O(g) Cl – Cl(g) → 2Cl(g) ∆H = 242.7 kJ∙mol–1 The more electrons in the bond (from single (2) to double (4) to triple (6)) the shorter the bond and the stronger the bond. So a triple bond is stronger and shorter than a double bond which is stronger and shorter than a single bond.