1. 11.2-11.3 Solids, Liquids and Gases, and Intermolecular Forces (IMFs)

2. Solids, Liquids and Gases On a molecular level, solids and liquids are more alike than gases. – That is, the molecules are touching in both the solid and liquid state, but are far from each other in the gaseous state. What is a major difference, on a molecular level, between a solid and liquid? The freedom of movement of the constituent molecules and atoms. Thermal energy partially overcomes the attraction in between the particles of a liquid. A quick review: Which states of matter have definite shape, definite volume and which are compressible? Solids have definite shape. Solids and liquids have definite volume (their molar volumes are very similar). Gases are compressible… liquids are slightly compressible, but since the particles are already touching, there is not much room for them to compress.

3. Solids Solids may be crystalline or amorphous. Crystalline solids, such as solid salts, have an atoms or molecules that are arranged orderly, while amorphous solids, such as rubber or glass, do not have atoms or molecules that are arranged orderly. In both cases, the motion of individual particles is limited.

4. Changes Between States What are two ways that we can change states of matter? Change temperature and/or pressure. Heating a solid past it’s melting point will transform the solid into its liquid state. A liquid can be transformed into a gas, both by increasing temperature and or reducing pressure. The exact opposite can be done to condense to the more dense state.

5. Intermolecular Forces (IMFs) The structure of a substance determines the IMFs that holds the substance together, which in turn determines state that the substance is in. In solids and liquids, IMF strength is moderate to strong, and for gases, IMFs are weak. IMFs form between charges, partial charges and temporary charges on molecules (or atoms and ions). Just as with ions, molecules attract to lower their potential energy, but IMF are much weaker than bonding forces. Again, here Coulombs law can be used. – Bonding forces are the attraction between large charges differences (attraction between electrons and protons) at short distances, while IMFs result from much smaller charges at further distances.

6. Types of Intermolecular Forces Dispersion forces Dispersion Forces (a.k.a. London Forces) the weakest IMF and is present between all molecules and atoms, but are the only IMF between nonpolar molecules.. – Result from fluctuations in electron distribution around a molecule or atom. For example, the reason that two molecules of chlorine, or two atoms of helium can attract to each other, is by the instantaneous (or temporary) dipole created by the mobility of electrons. The temporary dipole of one molecule induces a temporary dipole of a neighboring molecule. – The larger the electron cloud, the stronger the dispersion force. – The shape, not always the molar mass, of the molecule also determines the strength of the dispersion force. If two molecules have the same molar mass, but one is long and one is round and bulky, the longer molecule will have a higher boiling point because it has more areas of contact. – Molar masses can be used as a guide to determine magnitude of dispersion forces or boiling points within a group (or family). – The longer an alkane, the higher the boiling point.

7. Dipole-Dipole Forces These exist in all molecules that are polar. – Polar molecules have electron rich regions and electron deficient regions. Here, the negative end of the permanent dipole of one molecule, interacts with the positive end of the permanent dipole of another molecule. Remember, even polar molecules have dispersion forces. Here is an example of a dipole-dipole interaction between two molecules of formaldehyde (CH 2 O). Dipole-dipole interactions are stronger than dispersion forces so molecules with these interactions between them will have higher melting and boiling points. BP’s and MP’s increase with increasing dipole moments. Miscibility, the ability for one liquid to mix with another, is also based on polarity. Remember this saying, “likes dissolve likes”. Polar-polar mix, nonpolar-nonpolar mix, but polar-nonpolar do not.

8. Let’s Try a Practice Problem! Which molecules have dipole-dipole forces? a.) CI 4 b.) CH 3 Cl c.) HCl b.) Chloromethane and c.) hydrochloric acid both contain dipole-dipole interaction (and dispersion forces), but a.) carbon tetraiodide, only contains dispersion forces.

9. Hydrogen Bonding The interaction between the hydrogen of one highly polar molecule and the highly electronegative F, O, N of a neighboring molecule. This is the strongest intermolecular force, that can occur in pure substances. A sort of super dipole-dipole force. Still, is only 2-5% as strong as a covalent bond. Water molecules attracting to each other is an example of a hydrogen bond.

10. Let’s Try a Practice Problem Which has the higher boiling point, HF or HCl? Why? HF has the higher boiling point because the hydrogen bonds that hold hydrofluoric acid molecules together are stronger than the dipole- dipole interactions that hold the molecules of hydrochloric acid together. So it takes more energy to break the attraction between HF molecules that to break the attraction between molecules of HCl.

11. Ion-Dipole Interaction Occurs when an ionic compound is mixed with a polar molecule. – An example, table salt in water. Here, the positive sodium ions interact with the partially negative oxygen of a water molecule, and the negative chloride ions interact with the partial positive hydrogens of a water molecule, and the ions separate. This is the process of solvation (dissolving).

12. Let’s Try a Practice Problem! Which substance has the highest boiling point? a.) CH 3 OH b.) CO c.) N 2 a.) methanol (methyl alcohol), because hydrogen bonds (and dipole-dipole interactions and dispersion forces) exist between those molecules. Only dipole- dipole interaction and dispersion forces exist between molecules of carbon monoxide, and only dispersion forces exist between molecules of nitrogen gas.

13. 11.2-11.3 pgs.536-537 #’s 50, 54, 56, and 60 Read 11.4-11.5 pgs. 497-509