AP Chemistry 2014-2015

CH 10 INTERMOLECULAR

FORCES AND PHASES OF MATTER

 We’ve already looked at intramolecular forces (forces that hold together the atoms making up a molecule or compound). Intramolecular forces include ionic bonds, covalent bonds, and metallic bonds.

INTERMOLECULAR FORCES

 Intermolecular forces occur between molecules, rather than within them. The intermolecular forces present in a substance influence its characteristics. The table below summarizes the different types of intermolecular forces, from strongest to weakest.

Ion: a charged particle; can consist of one atom (monatomic ion) or several (polyatomic ion); can be positive (cation) or negative (anion).

Dipole: a separation of positive in negative charges; in this case, a dipole is found in polar molecules.

REVIEW TERMS

Induced dipole (new term): a temporary dipole that occurs in nonpolar molecules and atoms of noble gases due to an “unequal” distribution of electric charge. We usually assume that the electrons of an atom are uniformly distributed around the nucleus, but this is not always true. As electrons move around the nucleus, a momentary nonsymmetrical electron distribution can develop that produces a temporary (instantaneous) dipole. The attraction between induced dipoles and other induced

dipoles (or induced dipoles and ions or permanent dipoles) is weak and short-lived.

The tendency of an electron cloud to become distorted to give a dipolar charge distribution is called polarizability. Polarizability increases as the size of the atom increases, since larger atoms have more electrons. Think of (Captain) Jack Sparrow in Pirates of the Caribbean 2—when he discovers that “up is down”, and that he needs to turn the ship over, he can’t rock it very much by himself. When the crew catches on and starts running from side to side with him, the “polarizability” of the ship increases. Crew = electrons, ship = atom, tipping over = induced dipole. If you need to explain this on a test, make sure to mention the electrons!

Dipoles can also be induced in atoms, or nonpolar molecules, by the presence of ions or polar molecules. See the example on the right—iodine chloride is polar, so its permanent dipole is capable of inducing a temporary dipole in an atom of xenon.

Speaking of xenon…we’ve already seen that atoms of noble gases can have induced dipoles. Larger atoms are more polarizable , so the larger, heavier noble gases are more polarizable (and have stronger induced dipoles) than the lighter noble gases. You can see this trend in the table below. Helium is the lightest noble gas. It has the smallest number of

electrons, so it is the least polarizable, and its induced dipoles are relatively weak.

Xenon is a much heavier noble gas. It has a much larger number of electrons than helium, so it is much more polarizable, and its induced dipoles are much stronger. This is reflected in its higher freezing point

For London dispersion interactions to become strong enough to produce a solid, the motions of the atoms must be greatly slowed down. The noble gases generally have very weak London dispersion forces, which is why they have such low freezing points.

Element Freezing Point (°C)

Helium -269.7Neon -248.6Argon -189.4

Krypton -157.3Xenon -111.9

Hydrogen bonds: unusually strong interactions between two dipoles. Hydrogen bonds are found between partially positive hydrogen atoms and lone pairs on strongly electronegative atoms. They can occur between these atoms on two different molecules, or within the same molecule. The most familiar example of

hydrogen bonding occurs in water, where the partially positive hydrogen atom on one water molecule is attracted to the partially negative oxygen atom (and its lone pairs) on another water molecule.

Other elements that can interact with hydrogen atoms this way are nitrogen and fluorine. The hydrogen must also be covalently bonded to an electronegative atom (ex. NOT CARBON) in order for it to be able to form a hydrogen bond to an electronegative atom.

Hydrogen bonds are found in DNA. They link complementary base pairs together to make up the “rungs” of the DNA “ladder”.

Note that the “energy” listed in kJ/mol for each intramolecular/ intermolecular force (in the table on the first page) is the energy required to break or disrupt the corresponding attractive force. It’s generally higher for intramolecular forces than intermolecular forces. Here is one example—it takes 40.7 kJ to vaporize 1 mole of liquid water (disruption of intermolecular force, hydrogen bonds) white it takes 934 kJ to break the O-H bonds in 1 mole of water molecules (disruption of intramolecular force, covalent bonds). In other words, intramolecular forces (bonds) are stronger than intermolecular forces (interactions between molecules). Please note that hydrogen bonds are not true bonds;

they are intermolecular, not intramolecular. Hydrogen bonds are really just unusually strong dipole-dipole attractions.

High melting/boiling points of ionic compounds

Low melting/boiling points of molecular covalent compounds

High boiling point of water due to hydrogen bonding

An explanation for the gaseous state of the group 8A elements and BrINClHOF

SOME CONSEQUENCES OF IMFS

One important example of an ion-dipole interaction is the solubility of ionic compounds (and polar covalent compounds) in water, which is a polar liquid.

This helps to explain why polar or ionic compounds are soluble in water but nonpolar are not—the IMFs between ions and dipiole, as well as dipole and other dipoles, are stronger than the IMFs between dipoles (polar) and induced dipoles (nonpolar).

Identify the most important types of interparticle forces present in the solids of each of the following substances.a) HCl b) CaCl2

c) COd) BaSO4

e) Xef) C2H6

g) CsIh) P4

i) NH3

PRACTICE PROBLEM 1

Predict which substance in each of the following pairs would have the greater intermolecular forces.

SeO2 or SO2 CH3CH3 or H2CO

PRACTICE PROBLEM 2

The boiling point of HF is 20°C, while the boiling point of HCl is -85°C. Rationalize the difference between these two boiling points.

PRACTICE PROBLEM 3

In each of the following groups of substances, pick the one that has the given property. Justify your answer.

a) Highest freezing point: H2O, NaCl, or HF

b) Lowest freezing point: N2, CO, or CO2

c) Highest boiling point: HF, HCl, or HBr

PRACTICE PROBLEM 4

Structural model for liquids Gases have significant molecular motion but weak

intermolecular forces, while solids have little or no molecular motion but strong intermolecular forces. This makes gases and solids much easier to model than liquids, which have significant molecular motion AND strong intermolecular forces. Therefore, we won’t study the “structure” of liquids in depth.

THE LIQUID STATE

Surface tension is the resistance of a liquid to an increase in its surface area. Liquids with relatively large IMFs (such as those with polar molecules) tend to have relatively high surface tensions.

Capillary action is the spontaneous rising of a liquid in a narrow tube. Two different types of forces are responsible for capillary action: cohesive forces (cohesion) and adhesive forces (adhesion). Cohesive forces are the IMFs among

the molecules of a liquid Adhesive forces are the forces

between liquid molecules and their container; these occur when a container is made of a substance that has polar bonds (ex. glass). Adhesion allows a liquid to “pull itself” up a glass capillary tube to a height where the weight of the water just balances the liquid’s tendency to be attracted to the glass surface.

Polar liquids typically exhibit capillary action.

A concave meniscus shows that the liquid’s adhesive forces towards its container are stronger than its cohesive forces (ex. water).

A convex meniscus shows that the liquid’s cohesive forces are stronger than its adhesive forces towards its container (ex. nonpolar substances like mercury).

Viscosity is a measure of a liquid’s resistance to flow. Liquids with large IMFs tend to be highly viscous. Molecular complexity also leads to higher viscosity because very large molecules can become entangled with each other. Shorter hydrocarbon chains result in less viscous liquids than long hydrocarbon chains.

Viscosity is also affected by temperature—as you increase the temperature of a liquid, its viscosity decreases.

Hydrogen peroxide (H2O2) is a syrupy liquid with a relatively low vapor pressure and a normal boiling point of 152.2°C. Rationalize the diff erences of these physical properties from those of water.

PRACTICE PROBLEM 5

Crystalline solids are solids whose structure is very orderly/regular. The positions of the components in a crystalline structure are usually represented by a lattice, a three-dimensional system of points designating the positions of the components (atoms, ions, or molecules) that make up the substance. The smallest repeating unit of the lattice is called the unit cell.

SOLIDS

Ionic solids are solids composed of ions—in other words, they have ions at the points in their lattice structure. (Remember—ionic compounds contain metal cations and nonmetal anions). Ex. NaCl Ionic solids are stable, high-melting substances held

together by the strong electrostatic forces that exist between oppositely charged ions.

Molecular solids are solids composed of molecules—in other words, they have molecules at the points in their lattice structure. Ex. ice Molecular solids are characterized

by strong covalent bonding within the molecules but relatively weak forces between the molecules.

The forces that exist among the molecules in a molecular solid depend on the nature of the molecules. Many molecules such as CO2, I2, P4, and S8 have no dipole moment, and the IMFs are London dispersion forces. When molecules do have dipole moments, their IMFs are significantly greater, especially when hydrogen bonding is possible.

Atomic solids have atoms at the lattice points that describe the structure of the solid. Ex. carbon (as graphite, diamond, and fullerenes), boron, silicon, and all metals. Metallic solids specifically contain

metal atoms that are joined through a special type of delocalized nondirectional covalent bonding.

Network solids specifically contain nonmetal atoms bonded to each other with strong directional bonds leading to giant molecules of atoms called networks. Covalent network solids are very hard and have high melting points.

Group 8A solids specifically (and obviously) contain atoms of the noble gases which are attracted to each other with London dispersion forces.

Amorphous solids are solids whose structure is disordered. One example of an amorphous solid is glass.

Classify each of the following substances according to the type of solid it forms.a) Gold b) Carbon dioxidec) Lithium fl uorided) Kryptone) Silicon dioxidef) Methane, CH4

g) Iodine, I2

h) Wateri) Uraniumj) Phosphine, PH3

k) Diamondl) Hydrogen, H2

m) Quartzn) Ammonium nitrateo) C6H12O6

PRACTICE PROBLEM 6

Metals are characterized by high thermal and electrical conductivity, malleability, and ductility. These properties can be traced to the nondirectional covalent bonding found in metallic crystals.

The bonding in metals is both strong and nondirectional. Although it is diffi cult to separate metal atoms, it is relatively easy to move them, provided the atoms stay in contact with each other.

METALS AND ALLOYS

The simplest picture that explains these observations is the electron sea model, which envisions a regular array of metal cations in a “sea” of valence electrons. The mobile electrons can conduct heat and electricity, and the metal ions can be easily moved around as the metal is hammered into a sheet or pulled into a wire.

The band model (molecular orbital model) gives a more detailed view of the electron energies and motions in a metal. The electrons are assumed to travel around the metal crystal in molecular orbitals formed from the valence atomic orbitals of the metal atoms. These molecular orbitals are so closely spaced that they are essentially a continuum (called a band).

An alloy is a substance that contains a mixture of elements and has metallic properties. Alloys typically retain a sea of mobile electrons, and so remain conducting.

In a substitutional alloy, some of the host metal atoms are replaced by other metal atoms of similar size. For example, in brass, approximately 1/3 of the atoms in the host copper metal have been replaced by zinc atoms. The density of a substitutional alloy lies between those of its component metals.

An interstitial alloy is formed when some of the interstices (holes) in the host metal are occupied by small atoms. For example, steel contains carbon atoms in the holes of an iron crystal. The presence of the interstitial atoms changes the properties of the host metal. This is because the small atoms (ex. carbon) form strong directional bonds (ex. carbon-iron), which make the resulting alloy harder, stronger, less malleable, and less ductile. The more carbon you add, the harder the steel becomes.

Label each of the following structures by the type of alloy it represents.

PRACTICE PROBLEM 7

CARBON AND SILICON: NETWORK

ATOMIC SOLIDS

The two most common forms of carbon, diamond and graphite, are typical network solids. In diamond, the hardest naturally-occurring substance, each carbon atom is surrounded by a tetrahedral arrangement of other carbon atoms to form a huge molecule. Each carbon is sp3 hybridized. On the other hand, each carbon atom in graphite is sp2 hybridized. Graphite is soft because it is composed of layers that slide past each other easily.

NETWORK SOLIDS

Glass is an amorphous solid that exhibits a good deal of disorder. Glass more closely resembles a very viscous solution than it does a crystalline solid. The properties of glass can be varied greatly by varying the additives. For example, the addition of B2O3 produces borosilicate glass (used for labware and cooking) that expands and contracts very little under large temperature changes.

A glass menorah…with a base…that looks like the molecular structure of glass.

Yay science!

A semiconductor is a substance conducts only a sl ight electric current at room temperature, but shows increased conductivity at higher temperatures. Silicon is one such substance; it forms a 3d network (similar to a diamond) and its conductivity increases as its temperature increases . (Most conductive substances become less conductive as they heat up.) Remember that si l icon has four valence electrons—this is important in understanding two types of conductors made from si l icon. The conductivity of pure si l icon is too low for it to be used for many applications, so impurities of other elements are added to si l icon to increase its conductivity. The process of adding these impurities is cal led doping. n-type semiconductors are made by doping silicon (or germanium) with an

element with five valence electrons (P, As, Sb are common). This provides the semiconductor with extra electrons that can carry an electric current.

p-type semiconductors are made by doping silicon (or germanium) with an electron with three valence electrons (B is the most common, but Al, Ga, and In are also used). This creates electron vacancies, or holes, that can be fi lled with electrons. The holes travel in the opposite direction of the flow of electricity.

One example of a phase change occurs when a liquid evaporates from an open container. This is evidence that the molecules of a liquid can escape the liquid’s surface and form a gas, a process called vaporization, or evaporation. Vaporization is endothermic because energy is

required to overcome to relatively strong intermolecular forces in the liquid. The energy required to vaporize 1 mole of a liquid at a pressure of 1 atm is called the heat of vaporization, or the enthalpy of vaporization, ΔHvap.

VAPOR PRESSURE AND CHANGES OF STATE

When a liquid is placed in a closed container, the amount of liquid at first decreases but eventually becomes constant. The decrease occurs because there is an initial net transfer of molecules from the liquid to the vapor phase. This evaporation process occurs at a constant rate at a given temperature. As the number of vapor molecules increases, so does the rate of return of these molecules to the liquid. The process by which vapor molecules re-form a liquid is called condensation. Eventually, enough vapor molecules are present above the liquid so that the rate of condensation equals the rate of evaporation. At this point no further net change occurs in the amount of liquid or vapor because the two opposite processes exactly balance each other; the system is at equilibrium. Molecules are still constantly escaping from and entering the liquid at a high rate, but the two processes balance each other. Something else we can take from this…what is the vapor pressure of

water at 100°C?

The pressure of the vapor present at equilibrium is called the equilibrium vapor pressure (or just the vapor pressure) of the liquid. Liquids with high vapor pressures are said to be volatile (they

evaporate rapidly from an open dish). The vapor pressure of a liquid is principally determined

by the size of its IMFs. In general, substances with large molar masses have relatively low vapor pressures, mainly because of large London dispersion forces. Large/strong IMFs = low vapor pressure Small/weak IMFs = high vapor pressure

Vapor pressure increases significantly with temperature. Solids also have vapor pressures. Some solids sublime (ex.

iodine, dry ice).

There is one last thing for us to discuss about the diff erent phases of matter—heating curves! These are pretty simple—a heating curve shows temperature as a function of time as a solid substance is heated. What do each of the following segments represent?

 A-B = _____________________________  B-C = _____________________________  C-D =_____________________________  D-E = _____________________________  E-F = _____________________________  The melting point and boiling point can also be found on

heating curves. Where do you see them?  Where would we fi nd each of the following?

Heat (enthalpy) of vaporization Heat (enthalpy) of fusion