CHM 101 – Chapter Eight

1. CHM 101 – Chapter Eight Chemical Bonds, Lewis Structures & the Octet Rule Ionic Bonding Covalent Bonding Bond Polarity & Electronegativity Drawing Lewis Structures Resonance Structures Exceptions to the Octet Rule Strengths of Covalent Bonds CHM 101 - Reeves

2. Lewis Symbols Lewis symbols display atoms using their symbol surrounded by its valence electrons depicted as dots. Potassium Bromine Carbon CHM 101 - Reeves

3. Ionic Compounds • When metals react with nonmetals, electrons are transferred from the metal to the nonmetal, forming a cation and an anion. • Ionic compounds such as sodium chloride {NaCl(s)} are large arrays of cations and anions arranged so ions of opposite charge are as close as possible. CHM 101 - Reeves

4. Ionic Compounds The lattice energy is the energy required to break a crystal apart into the independent ions. • The potential energy of attraction between the ions depends directly on the charge of the ions, and inversely on the distance between ions. CHM 101 - Reeves

5. Covalent Compounds • In covalent compounds, atoms share electrons to achieve a stable electronic configuration. • While ionic compounds are typically brittle, crystalline and possess high melting points, covalent compounds tend to be gases, liquids or solids with low melting points. • Bonds involve the interaction of charged species, with like charges repelling and opposite charges attracting. CHM 101 - Reeves

6. Bond Polarity & Electronegativity • In homonuclear diatomic molecules such as fluorine {F2(g)}, the bonding electrons are shared equally by both nuclei. • When two different atoms are bonded together, the electrons are often unequally shared. This is true for hydrogen fluoride {HF(g)}. • An unequal electron distribution results in a separation of positive and negative charge; the bond is said to be polar. CHM 101 - Reeves

7. Bond Polarity & Electronegativity • To quantify the extent of bond polarity, Pauling assigned each atom an electronegativity: The attraction for the shared electrons the atom displays when involved in a bond. CHM 101 - Reeves

8. Bond Polarity & Electronegativity • Electronegativity generally increases from left to right and from bottom to top of the periodic Table. Fluorine, the most electronegative element, is assigned a value of 4.0 • The larger the electronegativity difference between the atoms involved in the bond, the more polar the bond. Ionic bonds represent the extreme case CHM 101 - Reeves

9. Drawing Lewis Structures • Lewis structures depict molecules as collections of atoms bonded together by shared electron pairs (depicted as lines). • In most structures, the combination of shared and lone pairs of electrons provides each atom (except hydrogen) with an octet of valence electrons. These structures obey the "octet rule" CHM 101 - Reeves

10. Lewis Structure of CCl4 • Sum the valence electrons from all atoms. Add electrons to account for negative charge, subtract to account for postitive charge. • Choose a central atom and arrange the other atoms around it. • Use the valence electrons to complete the octets of each of the surrounding atoms except hydrogen, which only requires two electrons. • All of the valence electrons have been used, and each atom has an octet. Complete the structure by replacing the bonding electrons with lines. CHM 101 - Reeves

11. Lewis Structure of SO3 • Sum the valence electrons from all atoms. Add electrons to account for negative charge, subtract to account for postitive charge. • Choose a central atom and arrange the other atoms around it. • Use the valence electrons to complete the octets of each of the surrounding atoms except hydrogen, which only requires two electrons. • The central atom is left with six electrons. To complete its octet, share a lone pair from one of the surrounding atoms • All of the valence electrons have been used, and each atom has an octet. Complete the structure by replacing the bonding electrons with lines. CHM 101 - Reeves

12. Lewis Structure of SO32- • Sum the valence electrons from all atoms. Add electrons to account for negative charge, subtract to account for postitive charge. • Choose a central atom and arrange the other atoms around it. • Use the valence electrons to complete the octets of each of the surrounding atoms except hydrogen, which only requires two electrons. • The central atom is left with six electrons. To complete its octet, add the last lone pair to the central atom • All of the valence electrons have been used, and each atom has an octet. Complete the structure by replacing the bonding electrons with lines. CHM 101 - Reeves

13. Lewis Structure of PH3 • Sum the valence electrons from all atoms. Add electrons to account for negative charge, subtract to account for postitive charge. • Choose a central atom and arrange the other atoms around it. • Use the valence electrons to complete the octets of each of the surrounding atoms except hydrogen, which only requires two electrons. • The central atom is left with six electrons. To complete its octet, add the last lone pair to the central atom. • All of the valence electrons have been used, and Phosphorus has an octet and athe hydrogen atoms have two each. Complete the structure by replacing the bonding electrons with lines. CHM 101 - Reeves

14. Lewis Structure of CO2 • Sum the valence electrons from all atoms. Add electrons to account for negative charge, subtract to account for postitive charge. • Choose a central atom and arrange the other atoms around it. • Use the valence electrons to complete the octets of each of the surrounding atoms except hydrogen, which only requires two electrons. • The central atom is left with four electrons. To complete its octet, share lone pairs from two of the surrounding atoms • All of the valence electrons have been used, and each atom has an octet. Complete the structure by replacing the bonding electrons with lines. CHM 101 - Reeves

15. Resonance When multiple bonds are present, the choice of the surrounding atom that receives the extra bond can be ambiguous. Consider SO3: In sulfur trioxide, all bonds are equivalent in length and strength, suggesting that the double bond is shared among all three oxygens. The three representations, which differ only by the placement of the electrons, are called resonance structures, as indicated by the double arrows. CHM 101 - Reeves

16. Resonance Which of the following exhibit resonance? • Resonance requires: • At least one double bond • At least two surrounding atoms that can accommodate a double bond. CHM 101 - Reeves

17. Exceptions to the Octet Rule There are three cases that produce exceptions to the octet (8) rule: 1) The central atom has less than 8 electrons. Occurs with Be (4), B (6) and Al (6). 2) The compound has an odd number of valence electrons 3) The central atom has more than 8 electrons. Can only occur with row three (n = 3) and higher elements. CHM 101 - Reeves

18. Bond Energies Bond energy is the minimum energy required to break a bond. Bond Energy for N2 In general, the shorter the bond, the stronger the bond. CHM 101 - Reeves

19. Bond Energies Bond breaking is endothermic; bond making is exothermic. By combining the energy absorbed by breaking reactant bonds with the energy released by forming product bonds, the enthalpies of gas phase reactions can be estimated. Estimate the enthalpy change for the reaction of hydrogen gas and bromine gas to form gaseous HBr Break reactant bonds: Make product bonds: CHM 101 - Reeves

20. CH3CHCHCH2CH3 + HBr CH3CH2CHBrCH2CH3 Bond Energies Estimate the enthalpy change for the following reaction. Break reactant bonds: Make product bonds: CHM 101 - Reeves