Types of Chemical Bonds

1. Types of Chemical Bonds

2. Bond Energy • Bond Energy – the energy required to break a bond • Atoms will bond in order to achieve the lowest energy configuration

3. Two Types of Bonds • #1: Ionic Bonds – When an atom with a high electron affinity reacts with an atom that loses an electron easily (metal with nonmetal) • #2: Covalent Bonds – When atoms share electrons (nonmetal with nonmetal)

4. Ionic Bond creates Ions • The energy between a pair of ions is calculated using Coulombs Law • Where • r is the distance between the ion centers • Q1 and Q2 are the ion charges

5. Size of the Ions • Remember Isoelectronic Ions? • Example 8.3: Arrange ions Se2-, Br-, Sr2+, Rb+ in order of increasing size

6. Bond Length • A molecule will position itself so the attractive forces are maximized and repulsive forces are minimized (energy is minimized) – this distance is called the bond length (from center of 2 atom to center of the other)

7. Covalent Bonding • Polar Covalent Bonds • Electrons are not shared equally due to electronegativity • Dipole moment is represented by the arrow pointing toward the negative side • Types of bonds with no dipole moments • Linear molecules with two identical bonds • Planar molecules with three identical bonds • Tetrahedral molecules with 4 identical bonds

8. Order the following bonds in terms of bond polarity • H-H, O-H, Cl-H, S-H, F-H

9. Percent Ionic Character • How can we tell the difference between a polar covalent bond and an ionic bond? • Percent ionic character of a bond = (measured dipole moment of X-Y) x 100%(calculated dipole moment of X+Y-)

10. Percent Ionic Character • Figure 8.13 – compounds with ionic character greater than 50% are normally considered to be ionic OR any compound that conducts an electric current when melted

11. % Ionic Character

12. Lattice Energy • Lattice Energy – the change in energy that takes place when gas ions are packed together to form a solid (energy released when an ionic solid forms) • Lattice Energy = k(Q1Q2/r) • Where: • k is a proportionality constant

13. Lattice Energy

14. Bond Energies • Scientists can calculate how much energy is required to break down a molecule • Depending on the bonds of the molecule, they modeled that each bond has a specific amount of energy – Bond Energies (this is purely a scientific invention)

15. Type of Bonds • Single Bond – sharing 1 pair of electrons (2 electrons) • Double Bond – sharing 2 pairs of electrons (4 electrons) • Triple Bond – sharing 3 pairs of electrons (6 electrons)

16. Bond Energy: table 8.4 pg 351 • Bond energy values can be used to calculate approximate energies for reactions • When bonds are broken – energy is added (endothermic) • When bonds are formed – energy is released (exothermic) • ΔH = ΣD(bonds broken) – ΣD(bonds formed) • Where: • Σ is the sum of terms • D is the bond energy per mole of bonds

17. Bond Energies

18. Use bond energies to determine the following: • Example: H2 + F2 2 HF

19. Use bond energies to determine the following: • Example: Calculate ΔH of methane with chlorine and fluorine to give Freon-12 (CF2Cl2), hydrofluoric acid, and hydrochloric acid.

20. Localized Electron Bonding Model – a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. • Lone pairs – localized on an atom • Bonding pairs – found between atoms

21. Application Rules • Description of the valence electron arrangement in the molecule using Lewis Structures • Prediction of the geometry of the molecule using the valence shell electron pair repulsion (VSEPR) model • Description of the type of atomic orbitals used by the atoms to share electrons or hold lone pair (Ch. 9)

22. Lewis Structure Steps for writing Lewis structures: • Sum the valence electrons from all atoms – this is the TOTAL number of electrons present

23. Lewis Structure Steps for writing Lewis structures: • Sum the valence electrons from all atoms – this is the TOTAL number of electrons present • Use a pair of electrons to form a bond between each pair of bound atoms • Arrange the remaining electrons to satisfy the octet rule for all elements • Compare the TOTAL number electrons to the number you drew – they must match – if not, add double bonds!

24. CH4 • CF4 • NH3 • BH3(watch out! This one has an exception!)

26. Exceptions to the Octet rule • C, N, O, F – always obey the octet rule! • B and Be often have fewer than 8 ve – they are very reactive • 2nd row elements cannot exceed the octet rule because their orbitals don’t allow it. • 3rd row elements can exceed the octet rule by using their empty valence d orbitals • Satisfy 2nd row elements first – then any left over electrons should be added to 3rd row elements that have an available d-orbital.

27. Exceptions to the Octet rule • C, N, O, F – always obey the octet rule! • B and Be often have fewer than 8 ve – they are very reactive • 2nd row elements cannot exceed the octet rule because their orbitals don’t allow it. • 3rd row elements can exceed the octet rule by using their empty valence d orbitals • Satisfy 2nd row elements first – then any left over electrons should be added to 3rd row elements that have an available d-orbital.

28. Exceptions to the Octet rule • C, N, O, F – always obey the octet rule! • B and Be often have fewer than 8 ve – they are very reactive • 2nd row elements cannot exceed the octet rule because their orbitals don’t allow it. • 3rd row elements can exceed the octet rule by using their empty valence d orbitals • Satisfy 2nd row elements first – then any left over electrons should be added to 3rd row elements that have an available d-orbital.

29. Exceptions to the Octet rule • C, N, O, F – always obey the octet rule! • B and Be often have fewer than 8 ve – they are very reactive • 2nd row elements cannot exceed the octet rule because their orbitals don’t allow it. • 3rd row elements can exceed the octet rule by using their empty valence d orbitals • Satisfy 2nd row elements first – then any left over electrons should be added to 3rd row elements that have an available d-orbital.

30. Exceptions to the Octet rule • C, N, O, F – always obey the octet rule! • B and Be often have fewer than 8 ve – they are very reactive • 2nd row elements cannot exceed the octet rule because their orbitals don’t allow it. • 3rd row elements can exceed the octet rule by using their empty valence d orbitals • Satisfy 2nd row elements first – then any left over electrons should be added to 3rd row elements that have an available d-orbital.

31. Exceptions to the Octet rule • C, N, O, F – always obey the octet rule! • B and Be often have fewer than 8 ve – they are very reactive • 2nd row elements cannot exceed the octet rule because their orbitals don’t allow it. • 3rd row elements can exceed the octet rule by using their empty valence d orbitals • Satisfy 2nd row elements first – then any left over electrons should be added to 3rd row elements that have an available d-orbital.

32. Draw Lewis Dot Structure for: • SF6 • ClF3 • XeO3 • RnCl2 • BeCl2 • ICl4-

33. What do I do with a charge? • First of all…what does the charge tell us? • So I just add or subtract from the total number of electrons! • Example: ICl4-

34. Formal Charge con’t • Atoms with a formal charge will.. • Try to achieve a charge close to zero • Any formal charges are expectred to reside on the most electronegative atoms

35. Resonance • Is invoked when more than one valid Lewis structure can be written for a particular molecule. • The resulting structure is an average of these resonance structures. • Ex. Nitrite Ion, Sulfate Ion

36. Example 8.10: • Give Possible Lewis Structures for XeO3, an explosive compound of xenon. Which Lewis structure or structures are most appropriate according to the formal charges?