1. Bonding Chapter 8

2. Types of Chemical Bonds Ionic Bonds – metals/nonmetals o Electrons are transferred o Ions paired have lower energy (greater stability) than separated ions o Electrostatic forces Covalent Bonds – nonmetals o Electrons are shared by nuclei o Pure covalent – non-polar covalent Electrons are shared evenly (F-F) o Polar Covalent Electrons are shared unequally Atoms end up with fractional charges o δ+ or δ-

3. E = 2.31 x 10 -19 J·nm Q 1 Q 2 r o E = Energy (J) o Q 1 and Q 2 = numerical ion charges o r = distance between ion centers in nanometers Negative sign indicates an attractive force o Release of energy ( ) Coulomb’s Law

4. Covalent Bond Length Distance at which the system energy is at a minimum Forces at work o Attractive forces – protons and electrons o Repulsive forces – electron-electron and proton-proton Energy is given off when two atoms achieve greater stability together than apart o Bond energy

5. Electronegativity The ability of atom in a molecule to attract shared electrons to itself Trend – increases across and up

6. Electronegativity & Bonds Greater electronegativity difference between two elements means less covalent character and greater ionic character Any compound that conducts an electric current when melted is an ionic compound If the electronegativity difference < 1.67, then the atoms will share electrons.

7. Bond Polarity & Dipole Moments Dipolar Molecules o Molecules with a somewhat negative and a somewhat positive end Dipole moment o Molecules with preferential orientation in an electric field o Slight negative side will be attracted to positive o All diatomic molecules with a polar covalent bond are dipolar

8. Bond Polarity & Dipole Moments Some molecules have polar bonds, but no dipole moment o Linear, radial, or tetrahedral symmetry of charge distribution o Charge balances/evens out CO 2 CCl 4

9. Bonding & Noble Gas e - Configurations Ionic bonds – electrons are transferred until each species attains a noble gas configuration Covalent bonds – electrons are shared in order to complete the valence configurations of both atoms Predicting Formulas of Ionic Compounds o Based on placement in the periodic table o Na  Na + Sizes of ions o Cations are smaller than parent ion o Anions are larger o Isoelectronic ions – size decreases as the nuclear charge increases

10. Binary Compounds Lattice Energy

11. Binary Ionic Compounds Lattice energy – change in energy that takes place when separated gaseous ions are packed together to form an ionic solid M + (g) + X - (g)  MX (s)

12. Determining ΔH f ° Metal Nonmetal Step 1: Sublimation o Solid  Gas Step 2: Ionization Energy o Gas  Ion Step 3: Bond Energy o Eg: Diatomic  Single Step 4: Electron Affinity o X + e -  X - Step 5: Formation of solid compound (LE) Sum = ΔH f °

13. Example – formation of LiF

14. Binary Ionic Compounds The formation of ionic compounds is endothermic until the formation of the lattice The lattice formed by alkali metals and halogens (1:1 ratio) is cubic except for cesium salts

15. Lattice Energy Calculations Lattice Energy = k Q 1 Q 2 r o k = constant dependent on the solid structure and the electron configurations o Q 1 and Q 2 = numerical ion charges o r = shortest distance between centers of the cations and the anions Lattice Energy increases as the ionic charge increases and the distance between anions and cations decreases Charge has more impact than distance

16. Partial Ionic Character of Covalent Bonds

17. ( ) % Ionic Character = measured dipole moment of X-Y x 100% Calculated dipole moment of X + Y - Ionic compounds generally have > 50% ionic character % ionic character is difficult to calculate for compounds containing polyatomic ions Calculating Percent Ionic

18. Covalent Chemical Bond Strengths of the Bond Model o Associates the quantities of energy with the formation of bonds between elements o Allows the drawing of structures showing the spatial relationship between atoms in a molecule o Provides a visual tool to understanding chemical structure Weaknesses of the Bond Model o Bonds are not actual physical structures o Bonds can not adequately explain some phenomena Resonance structures

19. Multiple Bonds Single bonds – 1 pair of shared electrons Double bonds – 2 pairs of shared electrons Triple bonds – 3 pairs of shared electrons As the number of shared electrons increases, bond length shortens Multiple bonds typically have higher bond energy

20. Bond Energy & Enthalpy ΔH = sum of energies required to break old bonds (endothermic) - sum of the energies released in forming new bonds (exothermic) ΔH = ΣD(bonds broken) – ΣD(bonds formed) o D = bond energy per mole o D always has a positive sign

21. Localized Electron Bonding Model Lone electron pairs o Electrons localized on an atom Bonding electron pairs o Electrons found in the space between atoms o Shared pairs Localized Electron Model o “A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.” Derivations o Valence electron arrangement using Lewis structures o Prediction of molecular geometry using VSEPR o Description of the type of atomic orbitals used to share or hold lone pairs of electrons