Basic Concepts of Chemical Bonding

1. Basic Concepts of Chemical Bonding Chapter 8 AP Chemistry

2. Bonding • Ionic – Electrostatic forces that exist between two ions of opposite charges transfer of electrons ( metal by non-metal) • Covalent – sharing of e’s between two atoms (two non-metals) • Metallic – found in metals where the atoms are bound to their neighbors but allow free movement of electrons (chapter 23)

3. G.N. Lewis Valence e’s – available to form bonds – reside in incomplete outer shell Lewis – developed system for noting val e’s – one dot for each val electron Place dots next to element following Hund’s rule. Lewis Symbols and the Octet Rule

4. Electrostatic attraction of Cations + by Anions – In order to maximize the attraction ionic solids exist in lattices Lattice energy – energy required to separate the crystalline solid into a gas (a mole of solid into gasious ions) Halite NaCl Rock Salt Ionic Bonding and Energy Relationships

5. Energy Relationships • removing an e from a metal is endothermic • when a non-metal gains an e it is exothermic • If e transfer was the only factor all ionic cmps would be endothermic • Because of the stability of the lattice structure (going to a lower energy state) releasing energy, all ionic cmps are exothermic.

7. Lattice Energy – tends to increase as the charge of the ions increases and the size of the ions decreases E = k Q1Q2 d Q1Q2 = charge of part d = distance between k = constant Bauxite Energy Relationships

8. Ionic Bond Formation • The representative elements form ions that have the noble gas configurations. Metals lose electrons and non-metals gain electrons. • The transitional metals do not always form electron configurations of the noble gases. The s sublevel electrons are the first to be transferred followed by the d sublevel electrons if necessary.

9. Ionic Bond Formation • Many stable ions are formed by emptying the s orbital or by leaving the d sub-shell full d10, half empty d5, or empty d0. • Polyatomic Ins – form when molecules gain or lose electrons. It is often unclear how (which atom gains or loses electrons) but the overall charge of the Ion is greater than the molecule. (has more of less electrons)

10. Atoms increase in size going from left to right in a period. Cations are smaller than their parent atoms. Due to the + charge and fewer screening electrons Anions are larger than their parent atoms . Size of Ions

11. The effective nuclear charge decreases because of the increase of screening electrons. Size of Ions

12. Is a series of ions or atoms that have the same electron configuration. Within a series the greater the atomic # (greater the number of protons- effective nuclear charge) the smaller the species. Isoelectric Series

13. Covalent Bonding • Electron pair sharing between two atoms – electrons attracted to both nuclei at the same time – covalent bonding results • Lewis described bonding patterns using electron dot symbols (Lewis structures) • Multiple bonds are formed when more than one pair of electrons is shared between atoms.

14. Single bonds – covalent bond one pair of shared electrons – ex: Cl-Cl structural, Cl2 molecular formula Double bonds – 2 pair of shared electrons, O=O structural, O2 molecular Triple bonds – 3 pair of shared electrons N N, Structural, N3 molecular Single, Double, and Triple Bonds

15. Bond Length and Strength • In general the distance between bonded atoms decreases as the number of shared electron pairs increases. • A triple bond is stronger than a single bond but not three times stronger. The second and third bonds are weaker than the first bond. The first bond in a triple bond is stronger than a “normal” single bond.

16. Strength of Covalent Bonds • Bond dissociation energy – Bond energy – enthalpy required to break a covalent bond is the average of the bond enthalpies of different molecules • Delta H (rxn) = Sum (bond energies of bonds broken) – Sum (bond energies of bonds formed) The aprox enthalpy of a reaction can be predicted by bond energies. This relates well to the heat of formations data that is more accurate.

17. Strength of Covalent Bonds • Between atoms of comparable size the greater the bond strength the shorter the bond length.

18. Covalent bonds are the result of electron sharing – types - Non-polar – when e’s are shared equally between nuclei – diatomic elements are an example Polar covalent – when e”s are not shared equally Bond Polarity and Electronegativity

19. If the electronegativities are equal (i.e. if the electronegativity difference is 0), the bond is non-polar covalent If the difference in electronegativities between the two atoms is greater than 0, but less than 2.0, the bond is polar covalent If the difference in electronegativities between the two atoms is 2.0, or greater, the bond is ionic Electronegativity – is a quantity that describes an elements ability to compete for electrons in a covalent bond. The greater the number the better it competes for e’s. Ways of noting partial +,- charges Electronegativity

20. Drawing Lewis Structures of Covalent molecules • Sum the valence e’s from all atoms in species • Write the atomic symbols showing how atoms are connected – drawl a single bond between atoms • Complete the octets of the atoms bonded to the central atom (peripheral atoms) • Place left over electrons on the central atom even if it results in the central atom having

21. Drawing Lewis Structures more than an octet. • If there are not enough e’s to give the central atom an octet, form multiple bonds by pulling terminal e’s from the peripheral atoms and placing them into the bond with the central atom. • Examples

22. Formal Charges • Formal charges are a way to assign all the valence elcectrons in a molecule to a “parent atom” • Rules 1. all bonding e’s are divided equally between atoms that form bonds. 2. all non-bonding e’s are assigned to the atom on which they reside.

23. Formal Charges • The formal charge is the number of valence e’s in the isolated atom (usually the group number in the periodic table) minus the number of electrons assigned by the rules • When several different Lewis structures are plausible, the one in which the formal charges are minimized is generally the preferred one.

24. Calculating Formal Charges

25. Examples

26. Practice Questions • Identify formal charges that are not zero

27. Answers

28. Exceptions to the Octet Rule • Most of the second period elements (C,H,O) are always observed with octets. Other elements do not easily achieve or ever achieve octets. • Molecules that contain odd numbers of e’s although they are uncommon and tend to be reactive are exceptions. • Light elements (H,Li,Be,B) tend to be surrounded by less than and octet of electrons.

29. Exceptions to Octet Rule • Third period elements and below in the periodic table are capable of expanding their octets because of the unfilled d orbitals thus having greater than eight electrons. • Examples on page 285-288