2. Bond Energy Bond Energy (BE): The energy required to break a bond. Bond Energy is important, bonds will be created if it allows the system (two or more atoms) to achieve the lowest possible energy state. Ionic substances are formed when an atom that loses electrons relatively easily reacts with an atom that has a high affinity for electrons (metal-nonmetal). Coulumb’s law can be used to determine the interaction between two ions: E = (2.31 x 10 -19 j-nm)(Q 1 Q 2 ) r

3. There is a limit to how close the atoms can get. The most stable distance is affected by electron repulsion and attraction between the two nuclei In Covalent bonds the most stable distance is the bond length. Bond Energy

4. Bond Energy (Covalent)

5. Electronegativity Electronegativity: The ability of an atom in a molecule to attract shared electrons to itself

6. Most bonds are a blend of ionic and covalent characteristics. Dependent on the differences in electronegativity. Bigger Difference  More Ionic Smaller Difference  More Covalent There are two types of covalent bonds: polar covalent and nonpolar covalent Nonpolar covalent bonds involve an even sharing of e -. Polar covalent bonds involve an unequal sharing of e -. Electronegativity

7. The range of electronegativity values is 4.0 (Fluorine) to 0.7 (Cesium) When the difference is negligible the electrons are shared equally creating a bond with covalent properties Between the extremes are cases when the electrons are shared unequally leading to a polar covalent bond with (fractional charges)

8. Electronegativity We can order bonds according to polarity: Order the following according to polarity: H-H, O-H, Cl-H, S-H, F-H H-H < S-H < Cl-H < O-H < F-H (2.1)(2.1) (2.5)(2.1) (3.0)(2.1) (3.5)(2.1) (4.0)(2.1) Electronegativity difference: 0 0.4 0.9 1.4 1.9

9. Electronegativity A: Covalent B: Polar Covalent (dipole) C: Separate Atoms

10. Electronegativity Differences Ionic bonds have an EN difference of approximately >1.7 Polar covalent bonds have an EN difference of approximately 0.3 to 1.7 Nonpolar covalent bonds have an EN difference of approximately <0.5

11. Polarity Dipole moment: When a molecule has a center of positive charge and a center of negative charge δ+δ+ δ-δ- H Cl

12. Polarity

14. Localized Electron Bonding Model LE model has three parts which we will apply: 1.Description of the valence electron arrangement in the model using Lewis structures 2.Prediction of the geometry of the molecule using the Valence Shell Electron Pair Repulsion (VSEPR) model 3.Description of the atomic orbitals used by the atoms to share the electrons or hold lone pairs

15. Lewis Structures Lewis structures of a molecule shows how the valence electrons are arranged among the atoms in a molecule. Example: For Ions only the valence electrons are shown K Br Notice no electrons are shown for K, it has no valence electrons left

16. Lewis Structures In covalent bonds the principle of achieving a noble gas configuration applies to the elements involved. Duet Rule: The formation of a bond to fill the 1s shell (applies to Hydrogen) Octet Rule: The second row non metals Carbon through Fluorine form molecules when their 2s and 2p (valence shells are filled.

17. Lewis Structures Steps for writing Lewis structures: 1.Sum the valence electrons from all the atoms. Do not worry about keeping track of which electrons come from which atoms. The TOTAL number is what is important. 2.Use a pair of electrons to form a bond between each pair of bound atoms. 3.Arrange the remaining electrons to satisfy the Duet rule for hydrogen and the octet rule for the second row elements.

18. Lewis Structures The second row elements C,N,O and F should always obey the octet rule The second row elements B and Be often have fewer than eight electrons around them. These electron deficient compounds are highly reactive The second row elements never exceed the octet rule (2s and 2p) can not accept more than 8 electrons 3 rd row and heavier elements often satisfy the octet rule but can exceed it by using the d orbitals When writing Lewis structures satisfy the octet rule for the atoms first then place the remaining electrons on elements with available d orbitals (central atom first)

19. Lewis Structures When two or more equivalent Lewis structures can be drawn then this is a Resonance situation. The actual structure is an average of the three resonance structures.

20. Lewis Structures Formal Charge: the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule. To determine formal charge we need to know two things: 1. The number of valence electrons on the free neutral atom (which has zero net charge because the number of protons) 2.The number of valence electrons “belonging” to the atom in a molecule

21. Lewis Structures Formal Charge: Formal Charge = (number of valence electrons on free atom) – (number of valence electrons assigned to the atom in the molecule) Assumptions: 1. Lone pair electrons belong entirely to the atom in question 2.Shared electrons are divided equally between the two sharing atoms (Valence Electrons) assigned = (# Lone pairs) + ½(number of shared electrons )

22. Lewis Structures Formal Charge