10.1 Intermolecular Forces | High School Chemistry

This paragraph introduces the concept of intermolecular forces, explaining that they are the forces between separate molecules. It contrasts these with the covalent bonds that hold molecules together, emphasizing that intermolecular forces are much weaker. The discussion also revisits the topic of ideal gases, clarifying that while they were previously described as having no attractive forces, there are indeed weak forces between molecules. The paragraph outlines three major categories of intermolecular forces: London dispersion forces (also known as van der Waals forces), dipole-dipole forces, and hydrogen bonding, with the latter being the strongest. The importance of opposite charges in chemical properties is highlighted, as is the attraction between positive and negative charges.

This paragraph delves deeper into dipole-dipole forces, explaining that these occur in polar molecules where there is a separation of charge, leading to partial positive and negative charges. The example of HCl is used to illustrate how the more electronegative chlorine atom becomes partially negatively charged, and the hydrogen partially positively charged. Hydrogen bonding is then introduced as a special case of dipole-dipole forces, occurring only with hydrogen atoms bonded to highly electronegative atoms like fluorine, oxygen, or nitrogen. The paragraph explains that hydrogen bonding is significantly stronger than other dipole-dipole forces and is crucial for understanding the unique properties of water, such as its high boiling point and expansion upon freezing.

This paragraph focuses on the importance of hydrogen bonding in water, detailing how it affects water's properties such as boiling point and its expansion when frozen. The explanation includes how hydrogen bonds are mediated through the lone pair of electrons on the partially negative atom, and how water molecules can act as both hydrogen bond donors and acceptors. The discussion also touches on the crystalline structure of ice, where each water molecule is hydrogen-bonded to four others, leading to the expansion of water when it freezes.

This paragraph explores how intermolecular forces influence the physical properties of liquids, such as boiling point, enthalpy of vaporization, viscosity, and surface tension. It explains that stronger intermolecular forces lead to higher boiling points and enthalpies of vaporization, as more energy is required to overcome these forces. Viscosity is also linked to intermolecular forces, with higher forces resulting in a thicker fluid. Surface tension is discussed in the context of water bugs being able to walk on water due to the strong hydrogen bonding at the liquid's surface. The paragraph concludes by contrasting these properties with vapor pressure, which is negatively correlated with intermolecular forces.

This paragraph discusses vapor pressure, explaining it as the measure of how many molecules have entered the gas phase above a liquid. It describes the equilibrium between the gas and liquid phases and how this is temperature-dependent. The paragraph clarifies that higher intermolecular forces lead to a lower vapor pressure, as fewer molecules can escape the liquid phase. The relationship between boiling point and vapor pressure is also explored, with the boiling point defined as the temperature at which a liquid's vapor pressure equals the external pressure.

This paragraph provides an in-depth comparison of intermolecular forces in different molecules, focusing on hydrogen bonding, dipole-dipole forces, and London dispersion forces. It explains that molecules capable of hydrogen bonding will have the strongest intermolecular forces, followed by those with dipole-dipole forces, and then those with only London dispersion forces. The discussion includes examples of how these forces affect the boiling points and vapor pressures of molecules like CH3OH, CH3F, and CH4. The paragraph also introduces ion-dipole forces, which occur in mixtures when an ionic compound is dissolved in a polar solvent like water.

This paragraph concludes the lesson by introducing the concepts of adhesion, cohesion, and capillary action. It explains that cohesion refers to the attractive forces between molecules of the same substance, while adhesion is the attraction between molecules of different substances. Capillary action is described as the ability of a liquid to rise in narrow tubes due to the adhesive forces between the liquid and the tube material. Examples of capillary action in nature, such as water traveling up tree roots, are provided, highlighting the biological relevance of these concepts.