Unit 3 Intermolecular Forces and Properties

Unit 3Intermolecular Forces and Properties Topic 3.3 Solids, Liquids, and Gases Text Reference Ch 11, section 11.1

Understanding: Matter exists in three states: solids, liquid, and gas, and their differences are influenced by variances in spacing and motion of the molecules. Learning Objective: Represent the differences between a solid, liquid, and gas phases using a particulate-level model. Essential Knowledge: Solids can be crystalline, where the particles are arranged in a regular three-dimensional structure, or they can be amorphous, where the particles do not have a regular, orderly arrangement. In both cases, the motion of the individual particles is limited, and the particles do not undergo overall translation with respect to each other. The structure of the solid is influenced by interparticle interactions and the ability of the particles to pack together. The constituent particles in liquids are in close contact with each other, and they are continually moving and colliding. The arrangement and movement of particles are influenced by the nature and strength of the forces (e.g., polarity, hydrogen bonding, and temperature) between the particles.

11.1 A Molecular Comparison of Gases, Liquids, and Solids • The fundamental difference between states of matter is the strength of the intermolecular (between molecules) forces of attraction. • Stronger forces bring molecules closer together. • Kinetic energy keeps them apart and moving. • Solids and liquids are referred to as condensed phases because their particles are fairly close together.

Table 11.1 Characteristic Properties of the States of Matter Properties of States of Matter *The atoms in a solid are able to vibrate in place. As the temperature of the solid increases, the vibrational motion increases.

Topic 3.1 Intermolecular Forces Text Reference Ch 11, sections 11.2, 11.3

Understanding: Intermolecular forces can explain the physical properties of a material Learning Objective: Explain the relationship between the chemical structures of molecules and the relative strength of their intermolecular forces when: a. The molecules are of the same chemical species.\Represent the differences between a solid, liquid, and gas phases using a particulate-level model. b. The molecules ar eof two different chemical species. Essential Knowledge: London dispersion forces are a result of the Coulombic interactions between temporary fluctuating dipoles. London dispersion forces are often the strongest net intermolecular forces between large molecules. a. Dispersion forces increase with increasing contact area between molecules and with increasing polarizability of the molecules. b. The polarizability of a molecule increases with an increasing number of electrons in the molecule and the size of the electron cloud. It is enhanced by the presence of pi bonding. c. The term “London dispersion forces” should not be used synonymously with the term “van der Waals forces.”

The dipole moment of a polar molecule leads to additional interactions with other chemical species. a. Dipole-induced dipole interactions are present between a polar and nonpolar molecules. These forces are always attractive. The strength of these forces increases with the magnitude of the dipole of the polar molecule and with the polarizability of the nonpolar molecule. b. Dipole-dipole interactions are present between polar molecules. The interaction strength depends on the magnitudes of the dipoles and their relative orientation. Interactions between polar molecules are typically greater than those between nonpolar molecules of comparable size because these interactions act in addition to London dispersion forces. c. Ion-dipole forces of attraction are present between ions and polar molecules. These tend to be stronger than dipole-dipole forces. The relative strength and orientation dependence of dipole-dipole and ion-dipole forces can be understood qualitatively by considering the sign of the partial charges responsible for the molecular dipole moment, and how these partial charges interact with an ion or with an adjacent dipole.

Hydrogen bonding is a strong type of intermolecular interaction that exists when hydrogen atoms covalently bonded to the highly electronegative atoms (N, O, and F) are attracted to the negative end of a dipole formed by the electronegative atom (N, O, and F) in a different molecule, or a different part of the same molecule. In large biomolecules, noncovalent interactions may occur between different molecules or between different regions of the same large biomolecule.

11.2 Intermolecular Forces • They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities. • The intermolecular attractions between molecules are not nearly as strong as the intramolecular attractions (bonds) that hold compounds together.

There are three types on intermolecular forces that exist between electrically neutral molecules: • Dispersion Forces • Dipole-Dipole Interactions • Hydrogen Bonding • The first two forces are collectively called van der Waals forces. • Another kind of attractive force, the ion-dipole force, is important in solutions. • All intermolecular interactions are electrostatic, involving attractions between positive and negative species, much like ionic bonds but much weaker

Table 11.2 Melting and Boiling Points of Representative Substances Relative Strengths of Attractions Intermolecular attractions are weaker than bonds. Hydrogen bonds are NOT chemical bonds.

1. Dispersion Forces • There are electrostatic interactions between electrically neutral atoms and/or molecules (such as liquified nonpolar gases). • The average distribution of electrons around a nucleus is symmetrical. However, the random motion of electrons in an atom or molecule can create an instantaneous, or momentary, dipole moment. • In He atoms, for example, if both electrons happen to be on one side of the atom, the atom has an instantaneous dipole. • The instantaneous dipole on one atom can induce an instantaneous dipole in an adjacent atom, causing the atoms to be attracted to each other.

Instantaneous and induced dipoles in He atoms • The tendency of an electron cloud to distort is called its polarizability.

Factors That Affect Amount of Dispersion Force in a Molecule • Number of electrons: the more electrons in an atom or molecule, the greater the dispersion force. • Increasing molecular weight: dispersion forces increase as molecular weights increase. • Figure to right shows the differences in boiling points between nonpolar halogen molecules and noble gases in the same periods.

Molecular shape: the greater the intermolecular contact due to shape, the greater the dispersion forces. • For example, neopentane and n-pentane have same molecular formula but differ in boiling points due to their shapes (linear vs spherical).

2. Dipole–Dipole Interactions • The oppositely charged ends attract each other.

For molecules of approximately equal mass and size, the strength of the intermolecular attraction increases with increasing polarity. • The figure below shows increases in boiling points as dipoles increase.

Which Have a Greater Effect: Dipole–Dipole Interactions or Dispersion Forces? • If two molecules are of comparable size and shape, dipole–dipole interactions will likely be the dominating force. • If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

3. Hydrogen Bonding • The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong. • A hydrogen bond is an attraction between a hydrogen atom attached to a highly electronegative atom and a nearby small electronegative atom in another molecule or chemical group.

Boiling point as a function of molecular weight • The nonpolar series (CH4 to SnH4) shows an increase in boiling points as the molecular weights increase. • Heavier members of groups 5A, 6A, and 7A follow the same trend. • However, H2O, HF, and NH3 have higher boiling points than expected due to hydrogen bonding between molecules. • Water has about 0.4 times the molecular mass of CH4 (18 g/mol vs 44 g/mol), yet it has an anomalously higher boiling point due to its strong hydrogen bonding attractions between water molecules.

Hydrogen bonds are a special kind of dipole-dipole attraction • Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. • These atoms interact with a nearly bare hydrogen nucleus (which contains one proton but NO inner electrons).

Ice Compared to Liquid Water • Generally, the solid state of a substance is more dense than the liquid state. • However, in water hydrogen bonding keeps the molecules further apart in the solid state than in the liquid state. • In the liquid state water molecules have random motion as hydrogen bonds are made and broken between molecules. • In ice the water molecules have far less motion and become locked into orderly arrangements by hydrogen bond attraction. • The lower density of water ice is understood in terms of hydrogen bonding. • As temperatures increase, the motion of water molecules increases and the rigid structure collapses.

Ion–Dipole Interactions • Ion–dipole interactions are a type of intermolecular attraction between an ion and a polar molecule found in solutions of ions. • The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents, like water. • The magnitude of attraction increases as either the ionic charge or the magnitude of the of the dipole moment increases.

Comparing Intermolecular Forces • Effects of forces are additive. • Multiple attractive interactions are possible between molecules.

Acetic acid and 1-propanol, for example, have the same molecular weight, but the attraction between acetic acid molecules is greater because adjacent molecules can form two hydrogen bonds versus the single hydrogen bond between 1-propanol molecules. • Hence, the boiling point of acetic acid is greater. • Because these molecules are polar, dispersion forces, though present, have far less effect on the molecular attraction.

Generalizations about Relative Strengths of Intermolecular Forces • When two molecules have comparable molar masses and shapes, dispersion forces are roughly equal. • Differences in magnitudes of intermolecular forces are due to differences in strengths of dipole-dipole interactions (if present). • Molecules capable of H-bonding have the strongest attractions. • When two molecules have very different molar masses and there is no H-bonding, dispersion force determines the substance with stronger attractions.

11.3 Select Properties of Liquids • Liquid properties are affected by the strength of intermolecular forces: • Boiling point/melting point (previously discussed) • Viscosity • Surface tension • Capillary action

Viscosity • Resistance of a liquid to flow is called viscosity. • It is related to the ease with which molecules can move past each other. • Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

Table 11.3 Viscosities of a Series of Hydrocarbons at 20 degree Celsius

Surface Tension • Water acts as if it has a “skin” due to an imbalance of intermolecular forces at the surface of the liquid. • Surface molecules experience a net downward force, called surface tension, since they are not attracted equally in all directions as are interior molecules. • Molecules at the surface are slightly closer together than deeper down in the liquid. • It causes water to “bead up” when in contact with nonpolar surfaces (such as glass or a waxy leaf).

Capillary Action • Intermolecular forces that bind similar molecules to one another, such as hydrogen bonding in water, are called cohesive forces. • Intermolecular forces that bind a substance to a surface are called adhesive forces. • These forces are important in capillary action which is the rise of liquids up very narrow tubes.

Water has stronger adhesive forces with the glass of a capillary tube than it has between its own molecules. Consequently, it forms a concave surface called a meniscus and rises up in a narrow tube. • Mercury has stronger cohesive forces between its atoms than with the glass. Consequently, it forms a convex surface and does not rise up in a narrow tube.

Unit 3 Intermolecular Forces and Properties