Chapter 10: Intermolecular Forces- The Uniqueness of Water

supercritical fluid

A substance at conditions above its critical temperature and pressure, where the liquid and vapor phases are indistinguishable and have some characteristic of both a liquid and a gas. Can penetrate materials like a gas but also dissolve substances in those materials like a liquid.

volatile

A word that describes substances if they evaporate readily at normal temperatures and pressures

Cgas = kH x Pgas Cgas= concentration of gas in a particular solvent kH= Henry's law constant for the gas in that solvent Pgas= the partial pressure of the gas in the environment surrounding the solvent

Equation for Henry's Law

In water molecules, the partial negative charges on oxygen atoms and the partial positive charges on hydrogen atoms result in attractions between a hydrogen atom in one molecule and an oxygen atom of another.

Explain dipole-dipole interactions in water molecules

Solubility in water is due to the balance between hydrophilic and hydrophobic interactions. As the hydrophobic portion of the molecule increases in the size, the entire molecule becomes more hydrophobic and solubility in water decreases.

Explain solubility in water for molecules that contain both polar and nonpolar groups.

boiling point

The greater the amount of energy required to separate the particles, the higher the

Intramolecular forces are stronger: strength is 100-1000 kJ/mol Intermolecular forces are weaker and typically act over longer distances: strength is 1-100 kJ/mol

Which are stronger: intermolecular forces or intramolecular forces?

(a) solid (b) gas

Which phase a substance is most likely to be the stable phase at (a) low temperatures and high pressures (b) high temperatures and low pressures?

dispersion forces: All substance have "clouds" that can be temporarily distorted.

Which type of intermolecular force exists in all substances?

Coulomb's law states that as charge increases, the attraction of two oppositely charged species for each other increases. Because of the full positive or negative charge on the ion, the ion dipole interaction is stronger than the dipole-dipole interaction.

Why are dipole-dipole interactions generally weaker than ion-dipole interactions?

Its strength is noticeably higher than other dipole-dipole interactions

Why are hydrogen bonds considered a special class of dipole-dipole interactions?

Because the atomic nuclei and the clouds of electrons shared by groups of atoms within a molecule interact with the electrons and nuclei in neighboring atoms or molecules

Why are the partial charges in temporary dipoles in molecules more likely to be distributed over greater distances than in single atoms?

Dipole-dipole interactions only involve partial charges, caused by unequal sharing of electrons within a molecule. In contrast, the ion involved in an ion-dipole interaction has lost or gained one or more electrons, and it has a full positive or negative charge,

Why aren't dipole-dipole interactions as strong as ion-dipole interactions?

When a salt dissolves in water, the ion-dipole interactions must overcome the electrostatic interactions between the ions themselves. As an ion is pulled away from its solid-state neighbors, it becomes surrounded by water molecules forming a sphere of hydration.

Why can a salt such as NaCl dissolve in water?

At high pressures, the volume of the gas becomes very small and free volume does not equal actual volume. At low temperatures, the kinetic energy of the gas molecules is reduced enough that the particles stick together (interact) via there intermolecular interactions.

Why do gases behave nonideally at high pressures and low temperatures?

High pressures push gas molecules closer together; low temperatures reduce their average speed. Both of these conditions favor intermolecular interactions because molecules that are closer together collide and interact with each other more frequently, and low temperatures slow molecules so they have more time to interact. Hence, both temperature and pressure influence the degree to which molecules interact.

Why do temperature and pressure have an effect on the phase of a substance?

concave meniscus (water)

adhesive > cohesive (water/glass > water/water)

ion-dipole interactions

an attractive force between an ion and a molecule that has a permanent dipole moment

dipole-dipole interactions

an attractive force between polar molecules

dispersion force

an intermolecular force between nonpolar molecules caused by the presence of temporary dipoles in the molecules

convex meniscus (mercury)

cohesive > adhesive (mercury/mercury > mercury/glass)

intramolecular forces

interaction between atoms within a molecule

cohesive forces

interactions between like particles

intermolecular forces

interactions between molecules/atoms

adhesive forces

interactions between unlike particles

meniscus

the concave or convex surface of a liquid in a small-diameter tube

surface tension

the energy needed to separate the molecules at the surface of a liquid

vapor pressure

the pressure exerted by a gas at a given temperature in equilibrium with its liquid phase

Henry's Law

the principle that the concentration of a sparingly soluble, chemically unreactive gas in a liquid is proportional to the partial pressure of the gas

polarizability

the relative ease with which the electron cloud in a molecule , ion, or atom can be distorted, inducing temporary dipole.

viscosity

the resistance to flow of a liquid

capillary action

the rise of a liquid in a narrow tube (a capillary tube) as a result of adhesive forces between the liquid and the tube and cohesive forces within the liquid. In a narrow tube, adhesion to the tube wall draws the outer ring of water molecules upward. At the same time, cohesive forces between the outer ring molecules and those adjacent to them draw the molecules upward. If the tube is narrow enough, this combination of adhesion and cohesion draws a column of water up.

hydrogen bond

the strongest dipole-dipole interaction. It occurs between a hydrogen atom bonded to a small, highly electronegative element (O, N, F) and an atom of oxygen and nitrogen in another molecule. Also includes molecules of HF

triple point

the temperature and pressure at which all three phases of a substance coexist. Under these conditions, freezing and melting, boiling and condensation, and sublimation and deposition all proceed at the same rate

normal boiling point

the temperature at which the vapor pressure of a liquid equals 1 atm

0.010 degrees celsius 0.0060 atm

triple point of water

increases

As temperature increases, vapor pressure

vapor pressure

At a given temperature, a more volatile substance has a higher ------- than a less volatile substance

the solvent-solute interaction are weaker than the solute-solute interactions that keep the solute together and the solvent-solvent interactions that keep solvent molecules together

Nonpolar solutes tend not to dissolve in polar solvents because

As temperature increases, the surface tension decreases because the molecular "film" on the surface of the liquid has fewer molecules held together by tight intermolecular forces. Likewise, the viscosity decreases as the temperature increases because molecules have more energy to readily break the intermolecular forces to enable them to slide past each other more freely.

Describe how the surface tension and viscosity of a liquid are affected by increasing temperature.

Surface tension is the resistance of a liquid to increase its surface area by moving the molecules of a liquid apart. The greater the intermolecular forces between the molecules in the liquid, the greater the surface tension.

Describe the origin of surface tension at the molecular level.

If a glass of water is covered, molecules at the surface still evaporate, but they are confined to the space above the water. Some of them condense at the liquid surface and return to the liquid phase. In a short time, the rates of the two processes equalize, meaning that as many molecules leave the surface as reenter it. At this point of dynamic equilibrium, no further changes takes place in the level of the liquid in the gas, even though molecules continue to evaporate and condense constantly. In such a situation at constant temperature, the pressure exerted by the gas in equilibrium with its liquid is called the vapor pressure of the liquid.

Describe the state of equilibrium involved with vapor pressure.

The trend in the melting/freezing equilibrium line for water shows a decrease in the melting point of ice as pressure increases. This trend in the melting/freezing points for water is opposite the trend for almost all other substances. The reason for water's unusual behavior is that water expands when it freezes. Most other substances are denser in the solid state than the liquid state because they contract when they freeze. Water expands when it freezes because hydrogen bonding between molecules of water in the solid phase creates a structure that is more open than the structure in liquid water, making ice less dense than liquid water. Applying enough pressure to ice forces it into a physical state (liquid water) in which it takes up less volume.

Explain the slope of the equilibrium line between solid and liquid on the phase diagram of water compared to most other substances.

A needle floats on water because it has a high surface tension due to its hydrogen bonding. The only intermolecular forces between methanol are dispersion forces.

Explain why a needle floats on the surface of water but sinks in a container of methanol.

Above the triple point, the solid phase must change to the liquid phase to enter the gaseous phase. Below the triple point, changing the temperature at a given pressure will sublime solid water into the gas phase.

Freeze drying is used to preserve food at low temp with minimal loss of flavor. Freeze drying works by freezing the food, then lowering the pressure with a vacuum pump to sublime the ice. Must the pressure be lower than the pressure at the triple point of H2O?

Within a sphere of hydration, the water molecules closest to the ion are oriented so that their oxygen atoms (negative poles) are directed towards a cation or their hydrogen atoms (positive poles) are directed toward an anion. The number of water molecules oriented this way depends on the size of the ion. The water molecules further from the ions are more randomly oriented.

How are the molecules in a sphere of hydration arranged?

The stronger the attractive forces in a solid, the greater the amount of energy needed to overcome those forces to cause melting or vaporization. Thus, a substance made of particles that interact relatively strongly has high melting and boiling points, which means the substance is likely to be a solid at room temperature and normal pressure. Under the same conditions, a substance with somewhat weaker particle particle interaction is likely to be a liquid. One with very weak particle particle interaction is likely to be a gas.

How do the attractive forces of a molecule influence its physical properties?

A molecule with a permanent dipole can create a dipole-induced dipole interaction when inducing a temporary dipole in a nonpolar molecule by perturbing the electron distribution in the nonpolar molecule. The intermolecular interaction in this case is weaker than the dipole-dipole force between two polar molecules and is of the same order of magnitude as the dispersion forces between temporary dipoles in nonpolar molecules

How does a dipole induced dipole interaction work?

In nonpolar compounds, viscosity increases with increasing molar mass. Larger molar masses mean stronger dispersion forces between molecules. With stronger interactions, larger/longer molecules do not slide past each other as easily as short molecules.

How does viscosity vary in nonpolar compounds?

The viscosities of polar compounds are influenced by both dispersion forces and dipole-dipole interactions. The stronger the interactions, the greater the viscosity.

How does viscosity vary in polar compounds?

As water is cooled to 4 degrees Celsius, its density increases, a pattern observed for most substances. However, as water is cooled from 4 to 0 degrees Celsius, it expands and its density decreases as the pattern of its hydrogen bonds changes with temperature. As water freezes at 0, its density drops even more, causing ice to float on water. This unusual behavior is caused by the formation of a network of hydrogen bonds in ice. The molecules form an extensive and open hexagonal network in the solid phase. Because of the space between the molecules in the network, the same number of molecules occupies more volume in ice than water. When ice melts, some of the hydrogen bonds in the rigid array break, allowing the molecules in the liquid to be arranged more compactly, up to 4 degrees where the density begins to decrease once more.

How does water's density change with temperature?

Solid CO2 does not melt into liquid at normal temperatures and pressures. Rather, it sublimes directly into gas.

In the phase diagram of carbon dioxide, the blue region representing liquid CO2 does not extend below 5.1 atm. What does this mean?

Triple point: solid, liquid and gas Critical point: liquid and gas

What phases of a substance are present at its triple point and critical point?

intensive

Is vapor pressure an intensive or extensive property of a liquid?

high melting points

Larger lattice energies typically correspond to

equilibrium lines

Lines separating the regions in a phase diagram. The two states bordering them are at equilibrium at the temperature and pressure combinations along such a line

sparingly soluble; soluble

Nonpolar molecules are very ----- in polar solvents but can be quite ------ in nonpolar solvents as a result of compatible solute-solvent interactions

dipole-dipole interactions between solute and solvent molecules

Polar solutes tend to dissolve in polar solutes because of

interactions

Polar substances with stronger ------------ tend to have higher boiling points than substances with similar molar masses but weaker -------------

critical point

The place where the boiling/condensation equilibrium line ends. A specific temperature and pressure at which the liquid and gas phases of a substance have the same density and are indistinguishable from each other. This point is reached because thermal expansion at this high temperature causes the liquid to become less dense, while the high pressure compresses the gas into a small volume, increasing its density.

induced dipole

The separation of charge produced in an atom or molecule by a momentary uneven distribution of electrons. This occurs when one atom's positive nucleus is attracted to the other atoms negative electrons, and vice versa, even as their electron clouds repel each other.

Relatively weak dipole-induced dipole interactions between water molecules and the nonpolar molecules of oxygen

What accounts for oxygen's (or any other sparingly soluble gas's) solubility in water?

1. Temperature: the higher the temperature, the greater the number of molecules with sufficient kinetic energy to break the attractive forces that hold them together in the liquid in the liquid and to enter the gas phase. 2. Surface area: the larger the surface area of the liquid, the greater the number of molecules on the surface in a position to enter the gas phase. 3. Intermolecular forces: the stronger the intermolecular forces, the greater the kinetic energy needed for a molecule to escape the surface, and the smaller the number of molecules in the population that have this energy.

What are the factors affecting evaporation rate?

three regions corresponding to the three phases of matter plus a fourth region called a supercritical region

What are the regions of a phase diagram?

its polarity

What influences water's physical properties (including its boiling point and melting point)?

Lattice energy is related to how much energy it takes to separate the ions in an ionic solid. Larger (more negative) energy values correspond to a stronger interaction between ions.

What is lattice energy?

Two substances are said to be miscible when they dissolve completely into each other. A substance is insoluble when it doesn't dissolve in a solvent.

What is the difference between the word miscible and the word insoluble?

Because all atoms and molecules have electrons, all atoms and molecules experience London forces to some degree. The larger the cloud of electrons surrounding the nucleus in an atom or several nuclei in a molecule, the more likely those electrons are to be distributed unevenly or polarized. Electrons in larger atoms are held less tightly by the nucleus because of both their greater average distance from the nucleus and the screening of the nuclear charge by electrons in lower energy orbitals. Consequently, they are more easily polarized than electrons in smaller atoms or molecules. Greater polarizability leads to stronger temporary dipoles and stronger intermolecular interactions, so London dispersion forces become stronger as atoms and molecules become larger.

Why do the strength of dispersion forces increase as the number of electrons in atoms and molecules increase?

The meniscus is the result of two competing forces: adhesive and cohesive. The cohesive forces are hydrogen bonds between water molecules and the adhesive forces are dipole-dipole interactions between the water molecules and the polar Si-O-Si groups on the surface of the glass. The adhesive forces are strong enough to cause the water to climb upward on the glass, creating the concave meniscus. The adhesive forces in this case are greater than the cohesive forces because surface water molecules next to the glass have less contact with other water molecules and therefore experience smaller cohesive forces.

Why does the meniscus form?

As temperature increases, the kinetic energy of the molecules of water and gas increases, and more energy is available to disrupt intermolecular attractions, reducing solubility. More gas molecules have the necessary kinetic energy to overcome the dipole-induced dipole interactions between the gas and the water molecules, and the dissolved gas escapes from the solution.

Why does the solubility of a gas decrease with increasing temperature?

The amount of gas that dissolves in a liquid depends on the frequency and number of collisions that the gas molecules have with the surface of the liquid: the more collisions, the more gas molecules that may form intermolecular attractions with the solvent . Conversely, when the pressure above a solution decreases, the solubility of the gas decreases.

Why does the solubility of a gas increase with increasing pressure?

At low temperatures, the solubility of a gas increases because fewer gas molecules dissolved in the solvent have sufficient kinetic energy to overcome the intermolecular forces between the solvent molecules. At low temperatures, the gas molecules are "trapped" in the solvent and cannot escape into the gas phase.

Why does the solubility of most gases in most liquids increase with decreasing temperatures?

When the average kinetic energy of of the liquid molecules increases, more of the molecules can escape the liquid phase and enter the gas phase. More molecules in the gas phase increase the vapor pressure.

Why does the vapor pressure of a liquid increase as temp increases?

This arises from the observation that most chemical reactions in nature, in our bodies, and in the laboratory take place at or near 1 atm

Why is it called the "normal" boiling point?

These molecules have more surface area to interact with each other, which gives more of a chance for interactions (dispersion forces) between them

Why is it that larger molecules (more spread out molecules) have higher boiling points?

SO3 is nonpolar and would therefore dissolve better in nonpolar solvents

Would SO2 or SO3 be more soluble in nonpolar solvents?

hydrophobic

a "water-fearing" or repulsive interaction between a solute and water that diminishes water solubility

hydrophilic

a "water-loving" or attractive interaction between a solute and water that promotes water solubility

phase diagram

a graphical representation of how the stabilities of the physical states of a substance depend on temperature and pressure

Clausius-Clapeyron equation

a relationship between the vapor pressure of a substance at two temperatures and its heat of vaporization. We can use this equation to calculate heat of vaporization if the vapor pressures at two temperatures are known.

dipole-induced dipole interaction

a temporary/induced dipole created in a nonpolar molecule when a polar molecule is in close proximity to it