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Functional Groups, Orbitals, and Geometry. Resonance Structures.

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Presentation on theme: "Functional Groups, Orbitals, and Geometry. Resonance Structures."— Presentation transcript:

1 Functional Groups, Orbitals, and Geometry

2 Resonance Structures

3 Bond Polarity - Part I  A bond is polar when the charge is not equally shared between the two atoms.  The more electronegative atom will have a partial negative charge ( δ - ). The arrow shows the dipole moment. Here we show partial charges.

4 Acids and Bases-Definitions  Arrhenius acid: A substance which dissolves in water to produce H +.  Brønsted-Lowry acid: a proton donor H + is a proton.  Lewis acid: an electron pair acceptor

5 Arrhenius Acids and Bases  Arrhenius acid: A substance which dissolves in water to produce H +.  Arrhenius base: A substance which dissolves in water to produce OH -. Limited to aqueous solutions. Does not explain a reaction such as NH 3 (g) + HCl(g)  NH 4 Cl(s)

6 Brønsted-Lowry Acids and Bases  B-L acids are proton donors. B-L bases are proton acceptors.  The emphasis is on the transfer of the H +. This links acids and bases. A B-L acid HB has a conjugate base: HB  H + + B: - This is the equation for HB acting as an acid.

7 Brønsted-Lowry Acids and Bases  HB  H + + B: - This is the equation for HB acting as an acid. B: - is the conjugate base.  B: - +H 2 O  HB + OH - This is the equation for B - acting as a base in water.  B: - + HA  HB + A - This is the equation for B- acting as a base with an acid other than water.  Be able to write these types of equations for any B-L acid or base.

8 Brønsted-Lowry Acids and Bases  Ammonia acting as an acid: NH 3  NH H +  Ammonia acting as a base: NH 3 (aq) + H 2 O  NH 4 + (aq) + OH - (aq)  What is the conjugate acid and what is the conjugate base of ammonia?  Is ammonia a conjugate acid or base?

9 Acid Strength and pK a HB H + + B: - K a = acid dissociation constant K a = [H + ][B - ] [HB] pK a = -log K a  The more completely an acid dissociates in water, the stronger it is. The stronger the acid, the larger its K a and the smaller its pK a.

10 Comparing Acid Strengths  Which is the stronger acid, ammonia or water?  There are two ways to find an answer: The quantitative way: compare pK a values. The qualitative way: compare the stabilities of the conjugate bases.

11 Comparing Acid Strengths  The quantitative way: compare pK a values. NH 3  NH H + pK a = 36 H 2 O(l)  H + (aq) + OH - (aq) pK a = 15.7  Water is the stronger acid.

12 Comparing Acid Strengths  The qualitative way: compare stabilities of the conjugate bases. NH 3  NH H + H 2 O(l)  H + (aq) + OH - (aq)  The more stable the conjugate base is in water, the stronger the acid. The amide ion is such a strong base it cannot exist in water, therefore ammonia is the weaker acid.

13 Comparing Acid Strengths  You will find it very helpful in studying organic chemistry to have a good idea of the relative strengths of some of the more common compounds acting as acids.  Please become VERY familiar with Table 1-5.

14 Comparing Acid Strengths by Comparing Structures  How does the structure of a compound affect its acid/base properties?  Look at the stability of the conjugate base. The more stable the conjugate base, the stronger its acid. Electronegativity Size/polarizability Resonance Stabilization Induction Hybrid orbital containing electrons

15 Comparing Acid Strengths by Comparing the Stabilities of the Conjugate Bases  Electronegativity (e.n.)  A more electronegative atom holds negative charge more easily. Many bases are anions. The more stable the anion, the weaker the base: e.n.(C) < e.n.(N)NH 2 - >OH - >F - Acid strength: CH 4

16 Comparing Acid Strengths by Comparing the Stabilities of the Conjugate Bases  Size  A larger anion is more stable: Size/stability: F - < Cl - < Br - < I - Acid strength: HF < HCl < HBr < HI Base strength: F - > Cl - > Br - > I -

17 Comparing Acid Strengths by Comparing the Stabilities of the Conjugate Bases  Resonance Stabilization  An anion stabilized by resonance has a stronger conjugate acid.

18 Comparing Acid Strengths by Comparing Structures  Induction  Look at nearby atoms. Electronegative atoms “pull” electron density away (induction). This can stabilize a negative charge. (Note: they must be very close to the negative charge to be effective.) Trichloroacetic acid is stronger than acetic acid. more stable

19 Comparing Acid Strengths by Comparing Structures  Hybrid orbital containing electrons  Acetylene (H-C ≡ C-H), believe it or not, can act as an acid with certain really strong bases. H-C ≡ C-H + B: -  H-C ≡ C: - + HB  The sp orbital is short (50% s character) and stabilizes the anion by holding the electrons closer to the nucleus.

20 Lewis Bases and Acids  Lewis looked at acid/base behavior from the viewpoint of the bonds that are formed instead of the transfer of a proton.

21 Lewis Bases and Acids  Lewis bases have nonbonding electrons that can be donated to form new bonds. Lewis bases are nucleophiles (lovers of nuclei +++).  Lewis acids accept these electrons. Lewis acids are electrophiles (lovers of electrons ---).

22 Two Bases Worth Knowing  NaH and NaNH 2 sodium hydride sodium ethoxide sodium amide sodium methoxide Given the reactants, be able to write the products of any acid/base reaction!

23 Identifying Bases  NaH and NaNH 2  Amines  Hydroxide ion, OH -  Alkoxide ions, e.g. CH 3 O -  Alcohols  Water

24 Identifying Acids  Inorganic (the seven strong acids)  Carboxylic acids  Phenols  Alcohols  Water  These are pretty much in order from strongest to weakest.


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