Intermolecular Forces PDF

INTERMOLECULAR FORCES Background Information for the Learners (BIL) In the preceding lesson, we have noted the differences in properties of matter in gas phase from those in the liquid and solid phases can be attributed to the attractive forces in...

INTERMOLECULAR FORCES Background Information for the Learners (BIL) In the preceding lesson, we have noted the differences in properties of matter in gas phase from those in the liquid and solid phases can be attributed to the attractive forces in solid and liquid molecules while gas molecules have negligible or no attractions at all. The condensation of gaseous substance to form liquids which in turn form solids could be explained by the attractive forces called intermolecular forces. Intermolecular forces vs. Intramolecular forces It is important to note the difference between intermolecular forces and intramolecular forces. As discussed in General Chemistry 1, atoms can form stable units called molecules by sharing electrons. This is called the intramolecular bonding. Intramolecular force in water molecule which hold hydrogen and oxygen atom. Intramolecular (within molecules) forces holds atoms together in a molecule. Intramolecular forces stabilize individual molecules. Generally, these forces are simply the chemical bonds such as ionic and covalent bonding. On the other hand, Intermolecular forces are attractive forces between molecules. Intermolecular forces are responsible for the non-ideal Intermolecular force in water molecule which hold two molecules together. behavior of gases, but they exert more influence in the condensed phases of matter which are liquids and solids. Intermolecular forces are collectively known as van der Waals forces named after Dutch chemist, Johannes van der Waal. Van der Waals forces are electrical in nature; that is, they result in the attraction between centers of opposite charge in two molecules close to each other. It is important to recognize that when a substance such as water changes from solid to liquid to gas, the molecules remain intact. The changes in states are due to changes in the forces among the molecules rather than in those within the molecules. In ice, the molecules are virtually locked in place, although they can vibrate about their positions. If energy is added, the motions of the molecules increase, and they eventually achieve the greater movement and disorder characteristic of liquid water. The ice has melted. As more energy is added, the gaseous state is eventually reached, with the individual molecules far apart and interacting relatively little. However, the gas still consists of water molecules. It would take much energy to overcome the covalent bonds and decompose the water molecules into their component atoms. This can be seen by comparing the energy needed to vaporize 1 mole of liquid water (40.7 kJ) with that needed to break the -OOH bonds in 1 mole of water molecules (934 kJ). Types of Intermolecular Forces The intermolecular forces of attraction in substances includes Dipole-dipole, London dispersion forces, hydrogen bonding and ion-dipole forces. London Dispersion forces London dispersion forces, or simply dispersion forces, are intermolecular forces of attraction that exist between all atoms and molecules. In addition, dispersion forces are the only kind of intermolecular forces present among symmetrical nonpolar substances such as O2 and CO2 and monoatomic species such as noble gases. Without dispersion forces, such substances could not condense to form liquids or solidify to form solids.  Dispersion forces are weak attractive forces that results from the continuous movement of electrons in particles. Nonpolar molecules have zero dipole moment because their electron density is uniform and symmetrical. Nevertheless, the electrons have some freedom to move around the molecule. This induces temporary dipoles (instantaneous dipoles) in neighboring atoms or molecules. As electron clouds become larger and more diffuse, they are attracted less strongly by their own positive nuclei. Thus, they are more easily distorted, or polarized by the adjacent/nearby nuclei. Polarization increases with increasing numbers of electrons and therefore with increasing size of molecules. Therefore, dispersion forces are generally stronger for molecules that are larger or have more electrons. As an example, between helium and argon, two argon atoms will have greater dispersion force because they are bigger than helium atoms. Dipole-dipole Forces Dipole-dipole forces are attractive forces between polar molecules, that is, between molecules that possess dipole moments. Their origin is electrostatic, and they can be understood in terms of Coulomb’s law. The larger the dipole moment, the greater the force. Dipole-dipole forces are the attraction between the positive end of one molecule and the negative end of another. Two molecules of HCl interacts. HCl has both positive and Dipoles form when there is a large negative end so it is a polar molecule and exhibits a dipole moment. The solid lines represent intramolecular force while difference in electronegativity between the broken lines represent intermolecular force. two atoms joined by a covalent bond. Hydrogen bonding  Hydrogen bond is a special case of very strong dipole-dipole interaction. They are not chemical bonds in formal sense. Strong hydrogen bonding occurs among polar covalent molecules containing H and one of the three small, highly electronegative elements – F, O, or N. Like ordinary dipole-dipole interactions, hydrogen bonds result from the attractions between + (partial positive) atoms of one molecule, in this case H atoms and the − (partial negative) atoms of another molecule. The + H is attracted to a lone pair of electrons on an F, O, or N atom. Typically, a hydrogen bond is about five to ten times stronger than other dipole-dipole interactions. The image shows the unusual high boiling points of NH3, H2O and HF compared with those other hydrides of the same group because of hydrogen bonding Ion-Dipole Forces Ion-dipole force acts between an ion (either cation or anion) and a polar molecule. When an ionic compound is placed in a solution of water, the positive end of the ionic compound becomes surrounded with the partial negative end of the ionic compound, in turn, becomes surrounded by the partial positive hydrogen ion in water. In short, the positive pole is attracted to the negative ion (anion), while the negative pole is attracted to a positive ion (cation).  Ion- dipole interactions are involved in solution process like in the case of sodium chloride (table salt) dissolving in water. In this case, Na+ and Cl- ions are dispersed amidst water molecules. The Na+ will be surround by the partial negative oxygen of water molecule while the Cl- will be surround by the partial positive H of water molecule. The strength of this interaction depends on the charge and size of the ion and on the magnitude of the dipole moment and size of the molecule. The charges on cations are generally more concentrated because cations are usually smaller than anions. Therefore, a cation interacts more strongly with dipoles than does an anion having a charge of the same magnitude. These four intermolecular forces vary in strength. Ion dipole forces is the strongest of the four, followed by hydrogen bonding being a special type of dipole- dipole. Dipole-dipole is weaker than the ion-dipole and hydrogen bonding while London dispersion forces is the weakest. Learning Competency: Describe and differentiate the types of intermolecular forces (STEM_GC11IMFIIIa-c- 100)