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# Comparing Acid Strengths by Comparing Structures - PowerPoint PPT Presentation

Comparing Acid Strengths by Comparing Structures. 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.

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Presentation Transcript

• 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

(CH3)3N + NH2-

• 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.

• A polar bond has a dipole moment μ:

μ(in debyes) = 4.8 δd

• δ is the charge at either end of the dipole

• d is the bond length in angstroms (charge separation) (1Å=10-10m)

dipole moment, μ

bond length, d

μ(in D) = 4.8 δd

• The dipole moment μ gives a quantitative measure of the polarity of a bond.

• C=O (2.4D) is more polar than C-O (0.86 D)

μ(in D) = 4.8 δd

• Knowing μ and d allows the charge separation δ to be calculated.

• C=O has a dipole moment of 2.4D and a bond length of 1.21Å.

δ= 2.4/(4.8x1.21)= 0.41

• C-O has a dipole moment of 0.86D and a bond length of 1.43Å.

δ= 0.86/(4.8x1.43)= 0.13

• The polarity (or dipole moment) of a molecule is the vector sum of the dipole moment for each bond in the molecule.

• A molecule with a significant dipole moment is polar.

• A molecule with little or no dipole moment is considered nonpolar.

• The dipole moment of a molecule can be measured.

• The dipole moments of the individual bonds can then be estimated.

• Lone pairs contribute to the dipole moments.

• arise from the charged nature of the subatomic particles (electrons and protons).

• are responsible for the cohesiveness of materials.

• are what determine physical properties of pure substances such asmelting point, boiling point, vapor pressure, and solubility.

• Substances that are gases at room temperature have weak intermolecular forces.

• Substances that are condensed (liquids or solids) at room temperature have much stronger intermolecular forces.

• If intermolecular forces did not exist, all substances would be gases, even at extremely low temperatures.

• Dipole-dipole

• generally attractive

• Hydrogen bonding

• a special category of very strong dipole-dipole force that involves the attraction between an electropositive H atom and nonbonding electrons on an electronegative atom (usually N, O, F, or Cl)

• London dispersion force

• instantaneous dipole-induced dipole

• increases with increasing surface area of the molecule

• present in all molecules

• Which will have the higher boiling point?

or

• Why does CCl4 have the higher boiling point?

chloroform, CHCl3 (μ = 1.0D)

or

carbon tetrachloride, CCl4 (μ = 0)

bp CHCl3 = 62°C

bp CCl4 = 77°C

• “Like dissolves like.”

• Polar substances dissolve in polar solvents.

• Nonpolar substances dissolve in nonpolar solvents.

• The other pairings (polar substance/nonpolar solvent and nonpolar substance / polar solvent) will not dissolve.

• For one substance to dissolve in another, there must be an attraction similar in magnitude to the forces holding the solvent together.

• In water, H bonding holds the molecules of water together pretty tightly.

• For a substance to dissolve in water, there must be an attraction between the substance and water that is close in magnitude to those H bonds.

• Ions, alcohols, and ethers all dissolve in water…can you show why?

• Carbon tetrachloride does NOT dissolve in water.

• Water is held together by H bonds, a strong intermolecular interaction.

• Carbon tetrachloride is nonpolar.

• The only force of attraction between CCl4 and H2O is dispersion, and that is not strong enough to push apart the H-bonded water molecules.

• Which are soluble in water and why?

http://www.agen.ufl.edu/~chyn/age2062/lect/lect_06/4_18.GIF