Intermolecular Forces Glenn V. Lo Department of Physical Sciences Nicholls State University

Explaining Properties Properties of pure substances depend on strength of attractive forces among particles (molecules or ions) Stronger attractive forces leads to Higher melting point higher boiling point Higher viscosity (resistance to flow) Higher surface tension (energy required to increase surface area)

Example: boiling point What does the fact the O 2 is a gas while H 2 O is a liquid at room temperature imply about the strength of intermolecular forces?

Intermolecular Forces Van der Waals forces Dipole-dipole interaction London dispersion forces (predominant) Dipole-induced dipole interaction Hydrogen bonding interaction Significantly stronger than van der Waals forces for small molecules

Van der Waals Forces Dipole-dipole interaction: + end of polar molecule tends to be attracted to – end of another polar molecule

Bond Polarity A polar bond is one where shared electrons spend more time near one of the two atoms. The shared electrons are said to be polarized towards the more electronegative atom. Bond dipole moment = a vector quantity; a measure of bond polarity:  = (  )(1.602x C)( r ),  = partial charge (or “ionic character”), r = bond length Unit for  ? 1 Debye = 3.336x C m Example: HF, experimentally found to have r = 92 pm,  =1.9 Debye Calculate . Ans. 0.43

Molecular Polarity Molecular dipole moment = vector sum of bond dipole moments; nonzero sum = “polar molecule” Polar Molecule Has at least one polar bond If >1 polar bonds: they must not be symmetrically oriented Example:  CO 2 is nonpolar  H 2 O is polar For which pair of molecules does dipole-dipole interaction exist: CO 2 and CO 2, H 2 O and H 2 O?

Measuring Molecular Polarity Put sample between electrically charged plates (capacitor). Polarity of sample affects the ability of the plates to store a charge. See PhET for animation of how polar molecule behaves between electrically charged plates.

Electronegativity scale Linus Pauling (1948): Electronegativity scale: 0.7 (Fr) to 4.0 (F), based on bond dissociation energies. Example: Bond dissociation energy of HF is larger than the average for H 2 and F 2. The difference is due to difference in electronegativity of H and F.  EN(H-F) = (eV) -1/2 [D H-F – (D H-H + D F-F )/2] 1/2 D H-F = eV, D H-H = eV, D F-F =2.75 eV EN of F was arbitrarily set to 4.0 Based on this, show that EN of H = 2.3 Based on various data, Pauling settled on a value of 2.1 for H. 1961, Allred and Pauling revised EN of H to 2.2 Other scales: Mulliken = EN = (IE + EA)/2, scaled to “match” Pauling’s values.

Predicting Bond Polarity Comparing with known  for HCl, HBr and HI, Pauling came up with:  = 1-exp(-  EN 2 /4), where  EN = difference in electronegativity Classifying bonds If  EN = 0,  = 0: “pure covalent” or “nonpolar” bond If  EN > 0, “polar covalent” bond, BUT typical rules of thumb are that:  If  EN > 2.0 (  > 63%): bond is essentially “ionic”  If  EN < 0.5 ((  < 6%), the bond is “essentially nonpolar”

Bond Polarity Classify the bond between C and H, Li and Br, Na and Cl, Na and Na.

Test Yourself Which pair of atoms would form a bond where the polarity is directed towards the first atom? A. C and C, B. H and F, C. O and C, D. I and Cl

Test Yourself Which pair of atoms would form a polar covalent but essentially nonpolar bond? A. N and N, B. H and C, C. Na and Cl, D. Be and I

Test Yourself Which of the following is/are true? 1) A nonpolar molecule can have polar bonds 2) A molecule with no polar bonds is nonpolar A. 1 only, B. 2 only, C. both, D. neither

Test Yourself Which of the following molecules is polar? A. H 2 O, B. BF 3, C. CH 4, D. PF 5

Test Yourself For which pair of molecules does dipole-dipole interaction occur? A. H 2 O and CO 2, B. H 2 O and HF, C. CO 2 and CO 2, D. CH 4 and CO 2

Van der Waals Forces London dispersion forces Motion of electrons around nuclei can temporarily lead to polarity. Attractions result from “fluctuating dipoles” Larger molecules are more polarizable Example: For which of the following pairs of molecules are attractions due London dispersion forces stronger: A. CH 4 and CH 4 B. C 2 H 6 and C 2 H 6 ?

Test Yourself For which pair of molecules is London Dispersion strongest? A. H 2 O and H 2 O, B. H 2 O and H 2 S, C. H 2 S and H 2 S

Test Yourself Which of the following has the lowest boiling point? A. O 2, B. F 2, C. Cl 2

Test Yourself Which of the following explains why the melting point of iodine is higher than that of iodine chloride? A. London dispersion forces are stronger among I 2 molecules, which are larger and more polarizable B. ICl molecules are capable of dipole-dipole interaction, whereas I 2 molecules are not C. Both, D. Neither

Van der Waals forces Dipole - induced-dipole interaction: between a polar molecule and a nonpolar molecule Polar molecule causes temporary polarization of nearby nonpolar molecule Example: For which of the following pairs do we expect dipole - induced dipole interaction? A. H 2 O and H 2 O B. CO 2 and CO 2 C. H 2 O and CO 2

Hydrogen bonding Exists between molecule with H bonded to F, O, or N, and another molecule that has a lone pair. Example:

Hydrogen bonding Extensive hydrogen bonding interaction explains why H 2 O has an anomalously high boiling point.

Test Yourself How many atoms in the acetic acid molecule (CH 3 COOH) are capable of hydrogen bonding interaction with water molecules? A. 1, B. 2, C. 3, D. 4

Test Yourself Explain why O 2 is a gas while H 2 O is a liquid at room temperature.

Test Yourself Watch the video at: Which liquid is more volatile: water or acetone? Why?

Explaining Solubility Interaction of two different substances depend on Strength of attraction among like molecules Strength of attraction among unlike molecules Rule of thumb for solubility: “Like dissolves like” Since H 2 O is polar, substances made of small polar molecules tend to be soluble in water than substances that are made up of nonpolar molecules. For nonpolar solvents, polar solutes tend to be less soluble than nonpolar solutes.

Example Explain why ethyl alcohol is soluble in water, but vegetable oil is not.

Example Consider the model shown below for a molecule of stearic acid, C 18 H 36 O 2. Based on this information, is stearic acid likely to be soluble or insoluble in water?

Test Yourself Which of the following is least soluble in water? A. NH 3, B. HF, C. CH 4

Test Yourself In which solvent is I 2 more soluble: H 2 O or CCl 4 ?

Large Molecules Typically: long chains of C atoms with hydrogens + some O, N, or other atoms O and N with lone pairs are “hydrophilic” --- attract water molecules Rest of the molecule = “hydrophobic” If molecule has no hydrophilic parts: insoluble in or immiscible with water Is the molecule shown on the right soluble or insoluble in water?

Solubility of Ionic Compounds in water What happens when an ionic compound dissolves in water? Ions have very strong attractions for water molecules. (“Ion-dipole” interaction) Why are some ionic compounds insoluble in water? Positive and negative ions also have strong attractions It takes a lot of energy to separate ions in a solid (“lattice energy”) Competition: ionic bonding in the lattice vs. ion-dipole interaction

Direct Observation of Molecular Polarity Questions to answer: What were the four liquids streams tested with the charged rod? Which two liquid streams were not attracted to the charged rod? Why? Which two liquid streams were attracted to the charged rod? Why? Why did it not matter whether the charged rod was positive or negative? Which of the two substances (cis-dichloroethylene and trans- dichloroethylene) caused an increase in the charging time of the electrically-charged plates? Why?

Test Yourself Watch the video at: Which liquid is polar?