Guide- Ib Chem Organic Intro

Organic Chemistry: IB Chem III Introductory Test 26 Sept. 2008

The test will cover:

1. Electron configuration and orbitals 2. Atomic/Mass Numbers and Weighted averages 3. Electronegativity and covalent/ionic bond types 4. Hybridization 5. Lewis structures 6. Acids and Bases – the three definitions 7. Ka and pKa 8. Orbital diagrams

1. Electron Configuration and Orbital Notation Electron configuration is written based on energy levels and number of electrons present in each orbital. Example:

In the above diagram, Sulfur (S, element 16) is described using electron configuration. One way to check the correctness of this type of notation is to count the “exponents,” which indicate the number of electrons in an element, and make sure that the number matches the element’s atomic number. Here, Sulfur’s atomic number is 16, and the sum of the superscripts is 16. Correct! There are 3 rules to remember for orbital notation: 1. Aufbau principle: Start with lowest energy level (1s) and move up. 2. There are never more than 2 electrons in a single orbital. 3. Hund’s Rule: Degenerate levels (same energy, same shape) get half-filled evenly before any remaining electrons are added to double them up. This makes the atom more stable. Illustration:

Oxygen, #8. Electron configuration 1s22s22p4: The first three p orbitals on the same energy level are half-filled, then they start to be completely filled, beginning with the lowest.

2. Atomic mass numbers and Weighted Averages

Weighted average = [(Mass # of isotope1)(% abundance)] + [(Mass # of isotope2)(% abundance)] 3. Electronegativity Using the electronegativity table, subtract e- values of each bonding element. If the difference is 2 or greater, the bond is ionic. Above about .3, the bond is polar covalent. Below that, the bond is nonpolar covalent. In general on the periodic table, electronegativity decreases toward the left and down the list.

4. Hybridization Number of bonds or electron pairs around an atom. Double and triple bonds are still only counted once! Electron pairs count!

Example: The carbon has two double bonds, so it is sp. Each oxygen has a double bond and two electron pairs, a total of three hybridizations—so it is sp2.

sp: 180 degrees: Two bonds/pairs sp2: 120 degrees: Three bonds/pairs sp3: 109.5 degrees: Four bonds/pairs (usually the maximum)

5. Lewis structures Draw them using dots for electron bonds (you can also draw lines, like in the structural formula), and depict electron pairs with double dots around the atom.

Eg. H:C:::C:H is the same as H-C≡C-C

6. Acids and Bases Definitions Type Arrhenius Bronsted-Lowry Lewis

Def. of Acid Increases [H+] Donates proton Accepts electrons

Def. of Base Increases [OH-] Accepts proton Donates electrons

If something has an extra electron pair, it’s probably a Lewis base. If it has an extra positive charge, it’s probably a B-L or Lewis acid. If it has an extra H, it’s probably an Arrhenius acid. Strong acid  Weak Conjugate Base Weak acid  Strong Conjugate Base Strong base  Weak Conjugate Acid Weak base  Strong conjugate acid

7. Ka and pKa Just as pH = -log [H+],

pKa = -log Ka. Reactions tend to produce (favor) the weaker acid, whether conjugate or not.

Strong acid = High Ka = Low pKa Weak acid = Low Ka = High pKa That is, pKa is directly related to weakness. So if one acid has a pKa of 16, and the conjugate acid has a pKa of 25, the reaction will proceed toward the right (the conjugate acid, the product) because that acid is weaker.

8. Orbital Diagrams Example: C=C (ignoring other attached atoms/electrons)

Steps: