Molecular Polarity & Intermolecular Forces

How to Predict Whether a Molecule is Polar or Nonpolar

You already learned how to determine whether a BOND is polar or nonpolar using a Table of Electronegativity. You learned to determine and draw the correct molecular geometry of a molecule. You can put these two pieces together and determine whether the overall molecule is polar.

Although you might think that a molecule which has at least one polar bond would automatically be a polar molecule, this would be an incorrect assumption. There are lots of nonpolar molecules which contain polar bonds. However, it is true that a molecule with no polar bonds and no lone pairs must be nonpolar.

How is this possible? Because of the overall 3-D shape! If the molecule and the polar bonds are symmetrically arranged, they may cancel out. Thus, the molecule would be nonpolar with polar bonds. Let’s take a look: CCl4 CH3Cl NH3 PH3 H2S SiH4

So not only are the individual bond dipoles important, but lone pairs of electrons affect the molecular polarity as well. Lone pairs of electrons pull electron density away from the central atom, so they have a dipole as well. We can draw an overall dipole for the entire molecule, called the dipole moment. This dipole moment is the vector addition of the individual bond dipoles as well as the pull of lone pairs on electron density. If you look at a table of molecular geometries, there are some generalizations about molecular polarity which can be made:

Your Turn: Determine whether CO2, CF4 , CH2Cl2, and H2O are polar or nonpolar

Partial Ionic Character We often say that a molecule like HF has “partial” ionic character. Or we can say that an “ionic” compound like AlN has “partial” covalent character. This can actually be calculated.

The dipole moment is defined as: μ= Qr If a molecule is polar, it has a nonzero dipole moment. If a molecule is nonpolar, it has a dipole moment of zero. The dipole moment is defined as: μ= Qr where μ is the Dipole moment in debye units (D), Q is the charge in coulombs (C), and r is the distance between the charges. The higher the dipole moment, the more polar the molecule is or you could say that the bonds have more ionic character.

We can use this equation to calculate a dipole moment, but as it is actually easy to experimentally measure a molecule’s dipole moment, we more often use the dipole moment to calculate the partial charge or %-ionic character of a molecule. HBr has a bond length of 141 pm and a dipole moment of 0.79 D. Given that the full charge of an electron is 1.60x10-19C and 1 D = 3.336x10-30Cm, what is the %-ionic character of the H-Br bond?

Intermolecular Forces

Intermolecular Forces You already know that CO2 is a gas at room temperature, while water is a liquid, and sucrose is a solid. Why? They are all molecular species. What holds water molecules together in the liquid phase at room temperature? Or what determines what phase or state a compound will exist in at room temperature? Let’s review the 3 states of matter:

Gas Liquid Solid no fixed volume fixed volume no fixed shape fixed shape, so rigid very low densities high densities density varies with T and P not compressible rapid, random motion fluid, motion little motion high kinetic motion some kinetic energy little kinetic energy

Intermolecular Forces So what determines the state of a compound at room temp? The strength of the attractions between separate molecules: the stronger the attraction, the more likel the compound will be a solid, the weaker the attractions, the more likely it will be a gas.

Intermolecular Forces These attractive forces “glue” solids or liquids together. What are these forces?

Intermolecular Forces The forces that hold individual molecules together in the solid or liquid phases are called intermolecular forces. They are also called van der Waal forces (although there is also a repulsive van der Waal force) They are responsible for many physical properties, including MPt and BPt

Intermolecular Forces There are several types DEPENDING on the MOLECULAR POLARITY! Here are the 3 Main Types Ion-Dipole (Ch 12) Dipole-Dipole (and H-Bonding) London Dispersion

Dipole Dipole Forces Attractive force between POLAR molecules (have a dipole moment) Electrostatic attraction of partial positive end of molecule to partial negative end of another molecule In liquid or solid, molecules align themselves so are attracted to several other molecules These dipole-dipole forces are much weaker than a real covalent bond, about 3-4kJ/mol. So they may be broken with a low amount of energy, so the solid melts, and the liquid evaporates!

431 kJ/mol 16 kJ/mol

Strength of Dipole-Dipole Forces There are 2 main factors in the strength of dipole-dipole forces: Distance between molecules (the closer they are, the stronger the dipole-dipole forces) Molecular Polarity (the more polar, the stronger the dipole-dipole force)

Hydrogen Bonding: A special case of dipole-dipole forces Some molecules have such a strong dipole-dipole force, that this extra-strong dipole force was given its own name: Hydrogen Bonding It is NOT really a bond, it is just an extra strong dipole-dipole force. It occurs under certain circumstances.

H-Bonding Occurs when a molecule is small and very polar Occurs when have a N-H, O-H, or H-F bond The molecules have a very large dipole moment, and they can get very close to one another due to the small size of H, N, O, F H-Bonding is responsible for the very high melting and boiling point of water H-Bonding is responsible for shape of proteins and DNA

London Dispersion Forces But CO2 is a nonpolar molecule, so how can it ever be a solid? There is another force, called the London Dispersion Force. The London Force acts upon ALL molecules, polar or not!

Loondon Dispersion Forces But London Forces are the ONLY intermolecular force that operates in nonpolar molecules. So how does it work when there are no dipoles to create an attraction? Instantaneous, induced dipoles!

London Dispersion Forces London Forces seem weak, but they can be very important, and many nonpolar molecules are solids or liquids. Waxes, oil, gasoline are examples.

London Dispersion Forces There are 2 factors in the strength of London Force: Size and mass of molecule or atom Shape of molecule

Effect of Size/Mass on London Dispersion Forces The larger the size or mass of a molecule, the more electrons it has. These electrons are also generally further from the nucleus. So the electrons can be distorted or pushed and pulled from one end of the molecule more readily, creating a larger temporary dipole.

Effect of Shape on London Dispersion Forces The closer molecules can align, the stronger the London Force, and the higher the melting and boiling points. Bulky molecules with “branching” can’t get as close, and so have lower London Forces and will have lower MPt and BPt than an “unbranched” molecule of the same mass.

Intermolecular Force Summary London Forces act upon all molecules, but are most important for larger, heavier molecules Dipole-dipole forces are found in polar molecules and depend on the size and polarity of the molecules H-bonding occurs when very small and very electronegative N, O, or F atoms are present along with N-H, O-H, or H-F bonds

Intermolecular Force Summary Although they are weak compared to covalent bonds, these intermolecular forces greatly affect the boiling and melting points of a compound The stronger the forces, the higher the melting and boiling points The stronger the forces, the more likely a compound will be a solid or liquid They also mean more deviations of a gas from the Ideal Gas Law!