Unit 2: Solutions

The rate of vaporization increases with:

-Increasing temperature -Increasing surface area -Decreasing strength of intermolecular forces (liquids that vaporize easily are volatile; others nonvolatile)

Relationship between torr and atm

1 atm = 760 torr (1 atm = 760 mm Hg = 760 torr)

Explains the heating curve and its calculations

1.) as solid approaches melting point 2.) actual melting of the solid 3.) liquid approaching boiling point 4.) actual boiling point 5.) heating vapor (gas) to a higher temperature beyond boiling point. Melting and boiling are constants. 1,3,5 calculations: mCΔT = #J (convert to kJ) (m: mass of substance (g); C: constant based on state; T: TempF - TempI) 2: ΔHfus multiplied by # moles = # kJ 4: ΔHvap mult. By # moles = # kJ Regions: 1: Cs; 3: Cl; 5: Cg

Solution equilibrium

A dynamic equilibrium is reached where the rates of dissolution (breaking into ions) and deposition (forming salt) become equal.

Solution

A homogeneous mixture of two or more substances Not always liquid Example: seawater Has at least 2 components: • Solvent: major component (water) • Solute: minor component (salt)

The critical point

A specific temperature and pressure where the gas and liquid phases of a substance commingle and are indistinguishable. There's no definite separation between liquid and gas (before there was gas above liquid). Now it's like a blob. You form supercritical fluids with the vaporization of a sealed liquid in heating.

Heat of vaporization value

Always + because it's an endothermic process (heat is absorbed)

Solubility

Amount of substance that will dissolve in a given amount of solvent. Depends on intermolecular forces

Why are intermolecular forces much weaker than bonding forces?

Because of large distances between molecules compared to the length of a bond

Boiling point

Boiling point: the temperature at which its vapor pressure equals the external pressure •An increase in the temperature of a liquid increases its vapor pressure (more molecules are able to vaporize) •Enough thermal energy for the interior of the liquid (not just the surface) to become gas •Normal boiling point: the temperature at which its vapor pressure equals 1 atm •The temperature of the liquid NEVER goes above its boiling point

Adhesive forces > cohesive forces

Capillary action will occur. the liquid will draw up while cohesive pull along those molecules not in direct contact with the wall (this happens until gravity balances)

Solids are _____ or ______

Crystalline (arrangement in a well ordered 3D array) or amorphous (no long range order; ex: rubber)

Intermolecular forces in nonpolar molecules

Dispersion forces

Intermolecular forces in polar molecules

Dispersion forces Dipole-dipole forces

Ice melting in a warm drink is an ________ process

Endothermic (melting ice absorbs heat from the liquid)

Spontaneous mixing

Energetically favorable to form uniform mixtures

Hydrogen bonding forces

Exhibited by polar molecules containing Hydrogen atoms bonded directly to F, O, or N atoms. A "super" dipole-dipole force (much stronger) Why: the large EN difference between H and these elements mean H will have a large partial + charge. These atoms are also quite small, so they can approach one another very closely. The strong attraction occurs between the H in each molecule and the F, O, or N on its neighbors

dipole-dipole forces

Exists in molecules that are polar. Polar molecules have permanent dipoles that interact with the permanent dipoles of neighboring molecules...these interactions involve the positive end of one polar molecule being attracted to the negative end of another polar molecule.

Where do intermolecular forces originate from?

From the interactions between charges, partial charges, and temporary charges on molecules/atoms/ions

Hydrogen bonding vs chemical bond

H bond is 2-5% as strong as a chemical bond but is still a strong intermolecular force.

ΔHfus

Heat of fusion The amount of heat required to melt 1 mole of a solid Value is positive (endothermic)

ΔHsub

Heat of sublimation The sum of the vaporization and fusion values (ΔHsub = ΔHvap + ΔHfus)

Increase in intermolecular forces means

Increase in melting and boiling points

Spectator ions

Ions that appear unchanged on both sides of the equation and do not participate. Omitted to form the net ionic equation

Dissociated (aqueous) substances are present as _______

Ions. Better to show the dissociated nature of compounds.

Relationship between kelvin and Celsius

K = C + 273

What results in a larger dispersion force?

Larger electron cloud (electrons held less tightly by nucleus & polarize more easily). All other variables constant, dispersion force generally increases with molar mass (more electrons dispersed over greater volume). Shape can play a role as well (linear shape allows for more interactions with neighbors than spherical shapes)

Unsaturated solution

Less than equilibrium amount (of solvent)

Give an example of pressure causing a phase change.

Liquid propane (pressure release causes change into gas)

Molecules with dipole-dipole forces have _____ melting and boiling points compared to molecules with hydrogen bonding

Lower

Molar mass of a gas

M = dRT/P

Polar molecules are ______ with other polar molecules but are _______ with other nonpolar molecules

Miscible Not miscible Why water and oil don't mix

intermolecular forces are ________ than bonding forces

Much weaker

Cohesive forces > adhesive forces

No capillary action

Solubility of ionic compounds

Not all ionic compounds dissolve in water. A compound is soluble if it dissolves in water and insoluble if it does not Soluble compounds are strong electrolytes Polyatomic ions dissolve as intact units (ex: Nitrate stays as Nitrate)

Ion-dipole forces

Occurs when an ionic compound is mixed with a polar compound (especially important in aqueous solutions). The positively charged end of a polar molecule is attracted to negative ions and the negatively charged end of the molecule is attracted to positive ions. Strongest intermolecular force

Boyle's Law

P1V1 = P2V2 Pressure and volume are inversely related (pressing on a ballon: increase pressure, decrease volume) How pressure changes according to volume. (n and T are constant)

Ideal Gas Law

PV = nRT P: pressure (atm) V: volume (L) n: moles R: universal (ideal) gas constant (0.08206 (L*atm)/(K*mol) T: temperature (K)

Compare the boiling and melting points of polar and nonpolar molecules.

Polar molecules have higher melting and boiling points than nonpolar molecules of similar mass. Due to stronger intermolecular forces

What type of molecules exhibit capillary action?

Polar molecules; nonpolar really don't

Slope of most fusion curves and why

Positive slope because increasing pressure favors the denser phase.... which most of the time is solid (water has negative slope... liquid more dense because water expands upon freezing)

Dispersion (London) forces

Present in ALL molecules and atoms. It's the result of fluctuations in the electron distribution within molecules or atoms (EN). Electrons may become asymmetrically arranged... this fleeting charge separation is called an instantaneous (or temporary) dipole (not stable). An instantaneous dipole in 1 atom induces instantaneous dipoles on neighboring atoms which then attract one another.

Condensation value in relation to heat of vaporization

Same amount but it's released (- value). Condensation is gas to liquid

Net ionic equation

Shows only the species that actually change change during the reaction

Molecular equation for aqueous solutions

Shows the complete neutral formation of each compound in the reaction as if they existed as molecules

Effect of temperature on the solubility of gases in water

Solubility decreases with increasing temperature (bubbles on bottom of a pot of boiling water)

Relationship between surface tension and intermolecular forces

Surface tension decreases with decreasing IFs

What are three properties affected by intermolecular forces?

Surface tension, viscosity, and capillary action

Phases of matter can be changed due to

Temperature, pressure, or both

Capillary action

The ability of a liquid to flow against gravity up a narrow tube •Combination of two forces: -Attraction between molecules in a liquid (cohesive) -Attraction between these molecules and the surface of the tube (adhesive) •Adhesive forces cause spreading out, while cohesive forces cause staying together •Capillary action will occur if adhesive forces are greater...causing the liquid to draw up while cohesive pull along those molecules not in direct contact with the wall (this happens until gravity balances) •The thinner the tube, the higher the rise •If cohesive forces dominate, no capillary action •Polar liquids tend to exhibit; nonpolar really don't

Why do liquids assume the shape of their containers?

The atoms or molecules that compose them are free to move around, but they're not easily compressed because the molecules or atoms are in close contact already

Why is water less dense as a solid than as a liquid?

The density of ice being less than "liquid water" can be attributed to hydrogen bonding. This can be related to the shape of snowflakes (no 2 are the same) Each molecule forms hydrogen bonds to 4 other molecules resulting in open cavities. As ice melts, motion of the molecules collapses the structure, resulting in a more random hydrogen bonding but strong enough to keep molecules close.

Why do solids have a definite shape and volume?

The molecules and atoms are fixed in place and have definite volume

Sublimation

The phase transition from solid to gas. Molecules (although fewer than liquid) do have enough energy to break free. The pressure of a gas in dynamic equilibrium with its solid is the vapor pressure of the solid

What's important in determining the miscibility of liquids?

The polarity of molecules composing liquids

Vapor pressure

The pressure of a gas in dynamic equilibrium. Depends on intermolecular forces and temperature: • weak IFs / volatile liquid / high vapor pressure : easily overcame by thermal energy • strong IFs / nonvolatile liquid / low vapor pressure : harder for thermal energy to overcome

Mole fraction

The ratio of the moles of solute in solution to the total number of moles of both solvent and solute x = n1/ n(total) Only works if you know the total pressure; can use this to find the partial pressures

Surface tension

The resistance of a liquid to increase its surface area. Liquids tend to minimize their surface area. A molecule at the surface has fewer neighbors to interact with. Molecules at the surface are less stable (have a higher potential energy). In order to increase the surface area, some molecules from the interior have to be moved to the surface...which is a process that requires energy •The energy required to increase the surface area by a unit amount is expressed as J/m2 and is dependent on intermolecular forces •Items "float" on the surface (even if more dense than the liquid) since the surface has a higher potential energy resulting in a minimization of surface area (an impenetrable skin) *Surface tension decreases with decreasing intermolecular forces*

What determines the phase of a substance at a given temperature?

The strength of the intermolecular forces between the molecules or atoms that compose a substance

How does the tube affect capillary action?

Thinner the tube, the higher the rise

Properties of gases

Uniformly fills any container, is easily compressed, mixes completely with any other gas, and exerts pressure on its surroundings

Charles' Law

V1/T1 = V2/T2 (P and n are constant)

Avogadro's Law

V1/n1 = V2/n2 (P and T are constant) (n = # of moles)

In a gas, the attractive energy between molecules is

Very low; allows for expansion and filling of a container

Give an example of temperature causing a phase change.

Water

Viscosity

a liquid's resistance to flow (SI unit: kg/(ms) Greater in substances with stronger intermolecular forces (more strongly attracted...less chance to move)...increase in molar mass •Also depends on molecular shape (increases in longer molecules that can interact over a greater area and become entangled) •Also depends on temperature (thermal energy partially overcomes intermolecular forces...allowing movement) •Viscosity decreases with increasing temperature

Phase diagram

a map of the phase of a substance as a function of pressure (y-axis) and temperature (x-axis). Regions: solid, liquid, gas Lines: sublimation, fusion, vaporization Triple point: all 3 states exist Critical point: end of vaporization line (supercritical fluid exists)

Entropy

a measure of energy randomization or energy dispersal in a system (disorder) •Two gases confined to their own space have a dispersal of kinetic energy when the barrier is removed (spontaneous mixing...greater entropy) •The tendency for all kinds of energy to disperse whenever it is not restrained from doing so is the reason that two ideal gases mix •Think about an iron rod being hot vs. cold

opposite of sublimation (and comparison)

deposition. Sublimation happens more often (escape to atmosphere) compared to deposition.

List the intermolecular forces from weakest to strongest.

dispersion, dipole-dipole, hydrogen bonding, ion-dipole

Freezing is an _______ process.

exothermic (heat is released when a liquid becomes a solid)

Opposite of melting is

freezing

What does the magnitude of the dispersion force depend on?

how easily the electrons can move or polarize (form a dipole moment) in response to an instantaneous dipole...which depends on the size (or volume) of the electron cloud

nonvolatile liquids

liquids that do not evaporate readily

volatile liquids

liquids that evaporate readily

fusion is also known as

melting point

Recrystallization

solid put into hot solvent...the saturated solution is cooled leading to supersaturation where excess solid comes out -Rock candy

Aqueous solutions

solutions with water as the solvent

Dalton's Law of Partial Pressures

states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture

Gas collection over water

subtract water vapor pressure from total pressure to get the gas pressure

Miscibility

the ability to mix without separating into two phases

Cohesive forces

the attraction between molecules in a liquid (within the liquid itself; to stay together)

Adhesive forces

the attraction between the liquid molecules and the surface of the tube. (Causes spreading out)

Saturated solution

the dissolved solute is in dynamic equilibrium with the solid (or undissolved) solute -Adding additional solute will not dissolve

Supersaturated solution

under certain circumstances, a solution can contain more than the equilibrium amount -Unstable and excess normally precipitates out

ΔHfus vs ΔHvap

ΔHfus is significantly less than the ΔHvap. Reason: Melting requires a smaller overcoming of IFs than does evaporation

Equation for overall energy change upon solution formation

ΔHsoln = ΔHsolute (+) + ΔHsolvent(+) + ΔHmix (-)

Precipitation reactions

•A solid forms upon the mixing of two solutions •Two aqueous (soluble) ionic compounds (reactants) yield another aqueous (soluble) ionic compound and a solid (insoluble) ionic compound (precipitate) •A precipitate does not always form when two aqueous solutions are mixed

Heat of vaporization (ΔHvap)

•Amount of heat required to vaporize one mole of a liquid to gas •Always + because it's an endothermic process (heat is absorbed) •Temperature has an effect on the value •Condensation from gas to liquid involves the same amount of heat but it is released

Explain dynamic equilibrium in a sealed water container

•As the concentration (partial pressure) of gaseous water molecules increases, the rate of condensation also increases •As long as the water remains at a constant temperature, the rate of evaporation is constant •A dynamic equilibrium is reached in that the rates of condensation and vaporization are the same •When a system in dynamic equilibrium is disturbed (pressure), the system responds so as to minimize the disturbance and return to a state of equilibrium

Fusion

•Back to solids, as the temperature is raised, the water molecules vibrate faster and faster •At the melting point (fusion), the molecules have enough thermal energy to overcome the intermolecular forces holding them in the solid framework and a liquid forms •The opposite of melting (liquid to solid) is freezing •Once the melting point is reached, additional heat only causes more rapid melting (the solid temperature does not increase)

Explain the effect of intermolecular forces on solution formation

•Depending on the forces in the solute and the solvent, those forces may promote the formation of a solution or prevent it •Three types of interactions: -Solvent-solute -Solvent-solvent -Solute-solute •Use the rule of thumb "like dissolves like" (both polar or nonpolar vs. polar and nonpolar)

Energetics of solution formation

•Energy changes (exothermic or endothermic) associated with solution formation can be understood in three steps 1.) Separating the solute into its constituent particles (always endothermic...ΔH > 0) - Energy required to overcome forces holding solute together 2.) Separating the solvent particles from each other to make room for the solute particles (always endothermic) - Energy required to overcome intermolecular forces of solvent 3.) Mixing (always exothermic... ΔH < 0) - Energy released as solute interacts with solvent •The overall energy change upon solution formation is the sum of all three steps •The overall sign for ΔHsoln depends on magnitudes

Effect of ΔHhydration in energy change of solution formation equation

•In terms of magnitudes: -If solute < hydration, solution is negative (exothermic) -If solute > hydration, solution is positive (endothermic)...if a solution forms at all -If solute is roughly equal to hydration, the solution process is neither exothermic nor endothermic

Factors affecting solubility

•Solubility of solids in water can be highly dependent on temperature -Sugar in hot vs. cold tea •The solubility of MOST solids in water increases with increasing temperature •Recrystallization: solid put into hot solvent...the saturated solution is cooled leading to supersaturation where excess solid comes out -Rock candy

Explain the various potential results from the energy change of a solution formation equation

•Sum of endo terms about equal to exo term - ΔHsoln is about 0...not much change in energy -Increasing entropy upon mixing drives the solution process •Sum of endo terms < exo term (a) - ΔHsoln is negative...exothermic process (T increase) - Tendency toward lower energy and greater entropy drives the solution process •Sum of endo terms > exo term (b) - ΔHsoln is positive...endothermic process (T decrease) - As long as ΔHsoln is not too large, the tendency toward greater entropy will still drive the solution process - If it is too large, a solution will not form

Energetics of melting and freezing

•The ice melting in a warm drink is an endothermic process (melting ice absorbs heat from the liquid) •The amount of heat required to melt 1 mol of a solid is called the heat of fusion (ΔHfus) -Value is positive (endothermic) •Freezing is exothermic (heat is released when a liquid becomes a solid) •In general, the heat of fusion is significantly less than the heat of vaporization -Melting requires a smaller overcoming of intermolecular forces than does evaporation

Vaporization

•The process by which thermal energy can overcome intermolecular forces and produce a phase change from liquid to gas •Both evaporation and condensation occur in an open beaker of water, but evaporation happens at a greater rate because the molecules escape and never find their way back. •Difference between evaporation and vaporization: vaporization is much quicker and we force it.

Explain the energetics of vaporization

•Vaporization is an endothermic process (it takes energy to vaporize the molecules in a liquid) -Remember that vaporization requires overcoming the intermolecular forces that hold liquids together •Condensation is exothermic (steam burn) -Gets "colder" in the desert than in a coastal area due to lack of humidity...less condensation so lower temperature

heat of hydration (ΔHhydration)

•When ionic compounds are dissolved in water, ΔHsolvent and ΔHmix combine to form the heat of hydration term •The enthalpy change that occurs when 1 mol of gaseous solute ions are dissolved in water •Always largely negative for ionic compounds (ion-dipole forces are strong)