2.3 Macromolecules Study Guide

2.3 Macromolecules (Chem.2.3)

Explore this Phenomenon

Water forms clumps or beads on the surface of a car, but soaks into the towel.

1. Why does the water bead up on the car, but it soaks into the towel?
2. What difference in properties do you expect will cause this difference?

Standard Chem.2.3
Engage in argument supported by evidence that the functions of natural and designed macromolecules are related to their chemical structures . Emphasize the roles of attractive forces between and within molecules. Examples could include non-covalent interactions between base pairs in DNA allowing it to be unzipped for replication, the network of atoms in a diamond conferring hardness, or the nonpolar nature of polyester (PET) making it quick-drying. (PS1.A)

In this section, look for the ways in which the structure of atoms and molecules determine the properties and functions of those molecules. Pay attention to how the structure of molecules relates to the forces of attraction between those  molecules and other molecules.

Forces of Attraction
The first type of force between molecules we will consider are called van der Waals forces, after Dutch chemist Johannes van der Waals (1837-1923). Van der Waals forces are the weakest intermolecular force and consist of dipole-dipole forces and dispersion forces.

Dipole-Dipole Forces
Dipole-dipole forces are the attractive forces that occur between polar molecules . A molecule of hydrogen chloride has a partially positive hydrogen atom and a partially negative chlorine atom. In a collection of many hydrogen chloride molecules, they will align themselves so that the oppositely charged regions of neighboring molecules are near each other.

Dipole-dipole forces are a result of the attraction of the positive end of one dipole to the negative end of a neighboring dipole.

Dipole-dipole forces are similar in nature, but much weaker than ionic bonds.

London Dispersion Forces
Dispersion forces are also considered a type of van der Waals force and are the weakest of all intermolecular forces. They are often called London forces after Fritz London (1900-1954), who first proposed their existence in 1930. London dispersion forces are the intermolecular forces that occur between atoms and between nonpolar molecules as a result of the motion of electrons.

The electron cloud of a helium atom contains two electrons, which can normally be expected to be equally distributed spatially around the nucleus. However, at any given moment the electron distribution may be uneven, resulting in an instantaneous dipole. This weak and temporary dipole subsequently influences neighboring helium atoms through electrostatic attraction and repulsion. It induces a dipole on nearby helium atoms (see Figure below ).

A short-lived or instantaneous dipole in a helium atom.

The instantaneous and induced dipoles are weakly attracted to one another. The strength of dispersion forces increases as the number of electrons in the atoms or nonpolar molecules increases.

The halogen group consists of four elements that all take the form of nonpolar diatomic molecules. Table below shows a comparison of the melting and boiling points for each.

The dispersion forces are strongest for iodine molecules because they have the greatest number of electrons. The relatively stronger forces result in melting and boiling points which are the highest of the halogen group. These forces are strong enough to hold iodine molecules close together in the solid state at room temperature . The dispersion forces are progressively weaker for bromine, chlorine, and fluorine and this is illustrated in their steadily lower melting and boiling points. Bromine is a liquid at room temperature, while chlorine and fluorine are gases, whose molecules are much further apart from one another. Intermolecular forces are nearly nonexistent in the gas state, and so the dispersion forces in chlorine and fluorine only become measurable as the temperature decreases and they condense into the liquid state.

Hydrogen Bonding
The attractive force between water molecules is a dipole interaction. The hydrogen atoms are bound to the highly electronegative oxygen atom (which also possesses two lone pair sets of electrons, making for a very polar bond. The partially positive hydrogen atom of one molecule is attracted to the oxygen atom of a nearby water molecule (see Figure below).

A hydrogen bond in water occurs between the hydrogen atom of one water molecule and the lone pair of electrons on an oxygen atom of a neighboring water molecule.

A hydrogen bond is an intermolecular attractive force in which a hydrogen atom that is covalently bonded to a small, highly electronegative atom is attracted to a lone pair of electrons on an atom in a neighboring molecule. Hydrogen bonds are very strong compared to other dipole interactions. The strength of a typical hydrogen bond is about 5% of that of a covalent bond .

Hydrogen bonding occurs only in molecules where hydrogen is covalently bonded to one of three elements: fluorine, oxygen, or nitrogen. These three elements are so electronegative that they withdraw the majority of the electron density in the covalent bond with hydrogen, leaving the H atom very electron-deficient. The H atom nearly acts as a bare proton , leaving it very attracted to lone pair electrons on a nearby atom.

The hydrogen bonding that occurs in water leads to some unusual, but very important properties. Most molecular compounds that have a mass similar to water are gases at room temperature . Because of the strong hydrogen bonds, water molecules are able to stay condensed in the liquid state. Figure below shows how the bent shape and two hydrogen atoms per molecule allows each water molecule to be able to hydrogen bond
to two other molecules.

Multiple hydrogen bonds occur simultaneously in water because of its bent shape and the presence of two hydrogen atoms per molecule.

In the liquid state, the hydrogen bonds of water can break and reform as the molecules flow from one place to another. When water is cooled, the molecules begin to slow down. Eventually, when water is frozen to ice, the hydrogen bonds become permanent and form a very specific network (see Figure below ).

When water freezes to ice, the hydrogen bonding network becomes permanent. Each oxygen atom has an approximately tetrahedral geometry – two  covalent bonds and two hydrogen bonds.

The bent shape of the molecules leads to gaps in the hydrogen bonding network of ice. Ice has an unusual property that its solid state is less dense than its liquid state. Ice floats in liquid water. Virtually all other substances are denser in the solid state than in the liquid state.

Putting It Together

Let us revisit this phenomenon:
Water forms clumps or beads on the surface of a car, but soaks into the towel.

1. How do the intermolecular forces explain why the water soaks into the towel and beads on the car?
2. Would you expect nonpolar oil to have the same interaction between the car and the towel? Explain.