Chemical Bonding

1. Chemical Bonding Chapter 8 AP Chemistry

2. Types of Chemical Bonds • Ionic – electrons are transferred from a metal to a nonmetal • Covalent – electrons are shared between 2 nonmetals • Metallic – between metal atoms

3. Lewis Symbols • Since bonding involves the element’s valence electrons, Lewis symbols are used. • Lewis symbols consist of the element symbol plus a dot for each valence electron. • Example: [Ne]3s23p4 (sulfur)

4. Which is the correct Lewis Structure for Chlorine? Either of these two!

6. Octet Rule • Atoms tend to gain, lose or share electrons until they are surrounded by 8 valence electrons. • In doing so, they achieve the configuration of a noble gas and achieve stability.

7. Ionic Bonding • Metals will lose electrons to achieve a stable outer valence level, and nonmetals will take those electrons to also achieve a stable outer valence. • NaCl is composed of Na+ and Cl- ions, arranged in a 3-dimensional crystal.

9. Formation of NaCl • There is an electron transfer between Na and Cl. • The metal, Na, has a low ioniziation energy and the nonmetal, Cl has a high electron affinity.

12. Energetics of Ionic Bonds • The formation of NaCl is extremely exothermic (ΔH°f = -410.9kJ). Why? • Loss of e- is always endothermic – takes energy to remove an electron from an orbital. When an e- is gained, the process is exothermic.

13. Lattice Energy • Lattice energy is a measure of just how much stabilization results from arranging oppositely charged ions in an ionic compound. Lattice energy is the energy required to completely separate a mole of solid ionic compound into its gaseous ions.

15. Lattice Energy • The magnitude of lattice energy depends on the charges of the ions, their sizes, and their arrangement in the solid. • For a given arrangement of ions, the lattice energy increases as the charges on the ions increases and as their radii decrease. • In forming ions, transition metals lose the valence shell e- first, then as many d e- as needed to reach the charge of the ion.

16. Covalent Bonding • A chemical bond formed by sharing a pair of electrons is called a covalent bond. • The attractive forces of the nuclei and electrons must overcome the repulsion between electrons and nuclei.

18. Lewis Structures for Covalent Bonds

23. Multiple Bonds

24. Multiple Bonds

25. Bond Polarity and Electronegativity • Polar Covalent Bond – one atom pulls harder for the shared electrons than the other atom, forming a dipole. • The difference in electronegativity between the atoms must be less than 1.7.

26. Bond Polarity • Non-polar Covalent Bonds – the atoms pull equally on the electrons. • The atoms must have the same electronegativity • The elements that have non-polar covalent bonds are the ones which are diatomic.

28. Dipoles F has a higher electronegativity, so it pulls harder on the (-) electrons creating an overall negative pole. C has a lower electronegativity, so the electrons are farther from it, and it has a partial positive charge.

29. Dipole Moment • The quantitative measure of the magnitude of a dipole is called the dipole moment, μ. • If a distance, r, separates two equal and opposite charges, Q+ and Q-, the magnitude is the product of Q and r. • μ = Qr

32. Dipole Moments • Reported in debyes (D), which equals 3.34 x 10-30 coulomb-meters (C-m) • Charge is measured in e (electronic charge), 1.60 x 10-19 C, and distance is in Å. • Suppose two charges 1+ and 1- are separated by 1.00Å. • μ = Qr = (1.60 x 10-19C)(1.00Å)(10-10 m/1Å) (1D/3.34 x 10-30) = 4.79 D

33. Drawing Lewis Structures • Sum the valence electrons from all atoms. • Write the symbols for the atoms to show which atoms are attached to which and connect them with a single bond. • Complete the octets around all of the atoms bonded to the central atom. • Place any leftover electrons on the central atom. • If there are not enough electrons to go around, try multiple bonds.

34. Resonance Structures • Multiple, but equally good representations for individual Lewis structures can be drawn for a molecule. • The resonance structures are “averaged”to give a more accurate description of the molecule.

35. Resonance Structures

38. Resonance in Benzene

39. Exceptions to the Octet Rule • The octet rule has limits when dealing with some of the transition metals and some covalent compounds. • Molecules and polyatomic ions with an odd number of electrons • Molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons • Molecules and polyatomic ions in which an atom has more than an octet of valence electrons

40. Strengths of Covalent Bonds • The stability of a molecule is related to the strengths of the bonds it contains. • Bond enthalpy is the enthalpy change, ΔH, for breaking a particular bond in one mole of a gaseous substance. • ΔHrxn = Σ(bond enthalpies of bonds broken) – Σ(bond enthalpies of bonds formed)

41. Bond Enthalpies, continued • H—CH3 + Cl—Cl → Cl—CH3 + H—Cl • ΔHrxn = [D(C—H) + D(Cl—Cl)] – [D(C—Cl) + D(H—Cl)] • = (413 kJ + 242 kJ) – (328 kJ + 431 kJ) = -104 kJ • This reaction is exothermic because the bonds in the products are stronger than than the bonds in the reactants, and the ΔHrxn value is negative. • These are often averaged values and provide an estimate.