Covalent and Ionic Bonding in Chemical Compounds

1. Covalent BondingChapter 6-1 (Old Text)Chapter 9-1 and 9-5 (New Text)

2. Na1+ Cl1- Ionic Bond Ideally, the electron clouds of each ion do not overlap. Electron transfer occurs from metal to nonmetal and the resulting ions attract only by electrostatic attraction.

3. Ionic Compounds have a rigid and strong crystal lattice structure in which the ions are arranged in an orderly, repeating pattern. Cl - ion Na+ ion

4. Covalent Bond Electrons are shared between two or more nonmetal atoms in a molecule. There are no ions! Electron clouds of adjacent atoms overlap and merge to form new bonding orbitals (a process called hybridization).

5. Most Covalent Substances Have a Molecular Form There are covalent bondsinside molecules but not between them! The attractions between molecules are called intermolecular forces (as opposed to intramolecular forces which are true covalent bonds). The strengths of the intermolecular forces determine the melting and boiling points of molecular substances. You don’t break any chemical bonds to melt molecular substances!

6. Electronegativity – A measure of the ability of an atom to attract shared electrons in a chemical bond , on a relative scale of 0 to 4. (American Chemist Linus Pauling developed this concept.) The noble gases are not ranked because they don’t form bonds. Notice that F is the best at attracting electrons in a bond.Finding the Difference in Electronegativity (ΔEN) between two bonded atoms can be used to predict whether a specific bond is more ionic or more covalent. In reality, most bonds are a little of both. ΔEN is calculated for each bond in a molecule. Just subtract the electronegativity values of the two bonded atoms.

7. EN > 2.1Ionic Bond – Electron Transfer Between 2 AtomsEN = 0.4 to 2.1Polar Covalent Bond – Unequal Sharing of the Shared Electrons Between 2 AtomsEN < 0.4Pure Covalent Bond (Nonpolar Covalent) – Essentially equal Sharing of the Shared Electrons Between 2 Atoms

8. ExamplesPredict the Bond Type based on EN:(EN values are always positive)The first one is done for you… • Na and Cl EN = 0.9 – 3.0 = 2.1, polar covalent bond • Ca and O • C and H • Al and Cl • Mg and H • S and O

9. H H O Br I Polar Covalent vs. Pure Covalent Bonds 2.1 2.1 The O-H bond (EN = 1.4) in water is a polar covalent bond. The shared electrons spend more time with oxygen than with hydrogen, thus the electron sharing is uneven. 3.5 2.8 2.5 The Br-I bond (EN = 0.3) in bromine iodide is a pure covalent (nonpolar) bond. The shared electrons essentially spend equal time with bromine and iodine.

10. Covalent Bonds are Like Springs! The atoms can vibrate back and forth to some extent without the bond breaking. The strength of a bond is related to its length, which is in turn related to the radii of the bonding atoms.

11. Bond Length • Bond length is the average distance between the nuclei of two bonded atoms. It is the position of highest stability that balances the forces of attraction and repulsion between two bonded nuclei. Energy is released when bonds form from atoms. The chemical potential energy of the chemical system is minimized (lowest) at the ideal bond length distance. Chemial Potential

12. Comparing Bond Length vs. Bond Energy Bond (Dissociation)Energy is the amount of energy needed to break a bond. It is the same amount of energy as was released when the bond formed. As bond length increases, bond energy decreases. In other words, longer bonds are weaker and easier to break than shorter bonds! (note: pm = picometer or 10-12 m)

13. Unstable Molecule – NI3 N I I I 2NI3(s)  ——>  N2(g)  +  3I2(g) Iodine is a much bigger atom than Nitrogen! The N-I bond is long and weak! Tickling from a feather is enough to break the bond, producing diatomic nitrogen gas and solid diatomic iodine (the purple smoke)! So, NI3 is very unstable, but N2 is very stable (triple bond)!

14. Table of Average Bond Dissociation Energies in kJ/mol(Bond Enthalpies) • These are average values because: • Atoms in a bond continually vibrate (move), and • The same type of bond may be measured inside different molecules and an average strength calculated

15. To calculate the energy needed to break all the bonds in a molecule…. • 1st - Draw the Lewis Structure of the molecule in order to see what specific bonds are in the molecule. • 2nd - Use a table of average bond energies to look up each type of bond. Pay specific attention to whether the bonds are single, double, or triple bonds. • 3rd - Add together the bond energies for each bond in the molecule. • Ex: Calculate the energy needed to break the bonds in Hydrocyannic Acid (HCN): Lewis Structure H – C  N : Total energy to break bonds in HCN is: 1 H-C bond (413 kJ/mol) + 1 CN bond ( 891 kJ/mol ) = 1304 kJ/mol (see the bond energies from the previous slide) Make sure you account for each and every bond in the molecule!

16. Single vs. Double vs. Triple BondsBond Length vs. Bond Energy When you compare bonds made between the same two atoms, triple bonds are generally stronger and shorter than double bonds which are generally stronger and shorter than single bonds!

17. You are done! • Now complete the study sheet using your freshly completed notes and this Power Point Presentation (to obtain electronegativity values and bond energies)