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Functional Groups, Orbitals, and Geometry

Functional Groups, Orbitals, and Geometry. Resonance Structures. 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.

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Functional Groups, Orbitals, and Geometry

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  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 NH3(g) + HCl(g)  NH4Cl(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:- +H2O  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: • NH3 NH2-+ H+ • Ammonia acting as a base: • NH3(aq) + H2O  NH4+(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 pKa HB H+ + B:- Ka = acid dissociation constant Ka = [H+][B-] [HB] pKa = -log Ka • The more completely an acid dissociates in water, the stronger it is. • The stronger the acid, the larger its Ka and the smaller its pKa.

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

  11. Comparing Acid Strengths • The quantitative way: compare pKa values. • NH3 NH2-+ H+pKa = 36 • H2O(l)  H+(aq) + OH-(aq) pKa = 15.7 • Water is the stronger acid.

  12. Comparing Acid Strengths • The qualitative way: compare stabilities of the conjugate bases. • NH3 NH2-+ H+ • H2O(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)<e.n.(O)<e.n.(F) • Base strength: CH3->NH2->OH->F- • Acid strength: CH4<NH3<H2O<HF

  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 NaNH2 sodium hydride sodium methoxide sodium amide sodium ethoxide Given the reactants, be able to write the products of any acid/base reaction!

  23. Identifying Bases • NaH and NaNH2 • Amines • Hydroxide ion, OH- • Alkoxide ions, e.g. CH3O- • 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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