Chemistry test 1: part 2

transitions from solid to liquid, or solid to gas can be triggered by

a change in temp or a change in pressure

fusion

At low temperatures, the molecules in a liquid sample lose the required kinetic energy to rearrange by breaking and reforming intermolecular interactions with their neighbors

• This results in the formation of a solid, typically a crystal

• The reverse of this process, where molecules in a solid gaining enough kinetic energy to break the strongly favorable interactions with their neighbors and forming a liquid, is called melting

what is the enthalpy of fusion

• The change in energy associated with the transition between the liquid and solid phases is called the enthalpy of fusion

• When a solid melts, the bonds between adjacent molecules are broken; this requires energy, making this an endothermic process

• When a liquid freezes, the system is able to lower its energy through the formation of more favorable intermolecular bonds; this is an exothermic process

• This quantity is not as commonly used as the enthalpy of vaporization because it depends on the crystal structure formed, and many molecules can form multiple

sublimation

at certain pressures and temps, a sample of matter in the solid phase can transition directly to the gas phase in a process called __

the reverse of this process is called deposition, when gas molecules transition directly into a solid

is sublimation endothermic or exothermic

molecules go from having their nearest neighbors at optimized positions in the solid to randomly positioned around them

it is endothermic because there is a penalty to losing those optimized interactions

what is the enthalpy of sublimation

is the energy change associated with the transition between the solid and gas phases

• When the solid sublimes, this breaks intermolecular interactions and increases the systems energy (endothermic)

• When the gas deposits, this forms intermolecular interactions (exothermic)

• The enthalpy of sublimation is always positive (endothermic)

what is a heating curve

represents the amount of energy needed to change the temp of a sample

includes single phase and multi phase regions

this includes the consideration of the heat capacity of different phases (using the 𝑞 = 𝑚 ∗ 𝑐 ∗ ∆𝑇 relationship), forming the sloped regions, as well as flat regions corresponding to the enthalpy associated with phase changes (𝑞 = 𝑛 ∗ ∆𝐻)

how can we interpret heating curve

The number of single-phase (sloped) regions tells us the number of different phases that occur for a sample in a given temperature range

• The number of multi-phase (plateau) regions tells us the number of phase transitions present

The temperatures associated with transitions (locations of plateaus)

• The enthalpies associated with phase transitions (width of plateaus)

• The heat capacities of different phases (slopes of single-phase regions)

phase diagram

plots that show the range of temps and pressures that different phases are stable

also show the lines that correspond to the transitions between these phases, or the sets of conditions where more than one phase will exist within a sample at equilibrium

what can we learn from a phase diagram

triple point and critical point

if we take a vertical slice, that is a constant temp system where we are varying the pressure

if we take a horizontal slice, that is a constant pressure system where we are varying the temp

many materials can form more than one type of crystal and thus have multiple temperature-pressure regions corresponding to those different crystal polymorphs

triple point

the joint between the three primary phases; all three phases coexist within a sample under those conditions

critical point

the endpoint of a vaporization curve; above that temperature and pressure, the matter exists in a state that is neither liquid nor gas

• this is called the supercritical fluid, which has intermediate and continuous density between the two main phases

solution

a uniform, homogenous mixture of two or more substances, which can be in any state

the solvent is usually water

the solute are

less abundant

solvent is

most abundant

why does a solution form

higher entropy states are thermodynamically favored

a mixed system will always have a higher entropy, so mixing will always be favored by entrophy

however, the favorability of the intermolecular forces (enthalpy) in the mixture compared to the separate phases will determine whether the formation of the solution is overall favored

how can we predict whether mixing will be favored or not

consider the IMF involved in the pure substances and in the mixture

if the forces in the solution are similar or more favorable than those in the pure state, mixing will be enthalpically favored

• can be summed up in the common aphorism “like dissolves like”

when mixing is enthalpically unfavorable,

whether mixing occurs depends on the relative magnitude of the entropic favorability and enthalpic unfavorability

we describe solutions as either miscible or immiscible, depending on whether the overall mixing is favored or not

miscible

liquids completely mix together

immiscible

do not mix together

how do solutions form on a microscopic level

When the solution is formed by two gases, the ballistic motion of the molecules cause the components to randomly intermingle and form a new solution

• There are no enthalpic costs or benefits to this mixing because there are no intermolecular interactions between the gas molecules

• For any mixture, there is an entropic benefit compared to the unmixed state

why does a solution form

entropy is commonly defined in popular culture as chaos, although this is not a very scientific definition

entropy is a measure of how many different ways its possible to have effectively the same macroscopic state of the system with different microscopic arrangements of particles

higher entropy states are more favorable

a mixed system will always have a higher entrophy

how do solutions form

For mixtures of gases, there are no intermolecular interactions

• For liquid-liquid or liquid-solid solutions, the intermolecular forces are different in the mixture compared to the unmixed state

• We need to compare the energies between the mixed and unmixed states

When a solution forms, some but not all intermolecular bonds with other molecules of the same type will be broken; these energies can be approximated using the enthalpy of vaporization (or sublimation) for each pure substance

Some intermolecular bonds with molecules of the other type will be formed during mixing; the strength of these interactions can be approximated by the enthalpy of vaporization for the solutio

how can we teste whether mixing is enthalpically favored or not

We can use the enthalpies associated with each of these steps to predict the enthalpic favorability of mixing using Hess’ Law

• If the mixture is more energetically favorable than the pure substances, its formation will ‘free up’ and release energy - This energy will be released as heat (exothermic)

• If the mixture is less energetically favorable than the pure substances, its formation will require energy The solution will take up heat from its surroundings (endothermic

how can we describe limits on solubility

• For any solute and solvent at a given temperature, there is a limit called the point of saturation

• Solutions can be nominally above this concentration (called supersaturated solutions), but the solute will precipitate out at the first opportunity to return to the saturation limit

how can we predict relative solubilities

• As mentioned previously, entropy always favors mixing, and the favorability of the change in enthalpy determines whether a solution forms

• We can predict relative solubilities based on the relative strengths of intermolecular interactions (favorabilities of enthalpies of mixing)

what impact does temp have on solubility

when a solid is dissolved in a liquid, the temp can impact how favorable the mixing is and thus the extent of solubility (how much solute will dissolve)

in general, the solubility of a solid will increase with temp (the direction of the trend with temp depends on the enthalpy of dissolution)

recrystallization - formed at high temps and cooled to lower temps, The lower solubility at low temperatures causes the solute to crash out into crystals

what factors influence the solubility of a gas in a liquid

At higher temperatures, gases become less soluble in liquids (more gas molecules evaporate and less dissolve back into the solution)

• At higher partial pressures of the solute gas, gases become more soluble in liquids (more gas molecules collide with the surface and dissolve into the solution)

what is henrys law

a quantitative description of the relationship between a gas partial pressure and its solubility in a solution

how do we describe solutions

by its concentration

this concentration can be expressed in Molarity, mole fraction, molality, mass %, parts per million/billion

molarity (M)

moles solute/liters of solute

mole fraction

moles of solute/moles of solution

molality (m)

moles of solute/kg of solvent

mass %

mass of solute/mass of solution x 100%

parts per million/billion

mass solute/mass of solution (scaled)

electrolytes

when some compounds dissolve in water and dissociate into ionic species

not all ionic compounds will dissociate even if they do dissolve in water

there are also some molecular compounds that also dissociate into electrolytes

the degree of dissociation is important because there is a class of properties that depend on the concentration of solute in a solution

strong electrolyte

a compound that fully dissociates into ions when dissolved

weak electrolyte

a compound that partially dissociates into ions when dissolved

nonelectrolytes

compounds that do not dissociate when dissolved

common electrolytes

acids and bases

soluble ionic compounds

van’t hoff factor

when one formula unit of an electrolyte dissociates in solution, it produces on average more than one dissolved solute particle

we can represent the average nymber of solute particles produced per formula unit of solute dissolved in solution using i

what properties of solutions are different than of the pure solvent

Because the intermolecular forces within a solution are different than those of the pure solvent, many properties are different for solutions than for the pure solvent

• For relatively dilute solutions, we can often treat the solution as pure solvent

• For more concentrated solutions, we need to consider the presence of the solute

what are colligative properties

class of properties of a solution that is affected by the presence of the solute relative to the pure solvent

the change due to the presence of solute is only determined by the amount of solute present; the specific identity of the solute does not matter

four colligative properties

vapor pressure reduction

boiling point elevation

freezing point depression

osmotic pressure (unique to solutions)

vapor pressure is __ by the presence of solute in a solvent

reduced

in pure solvents, all molecules at the surface are solvent, and can evaporate freely

in solutions, some solute molecules are at the surface instead, meaning

there are fewer chances for the solvent to evaporate at the surface

since solvent molecules in the gas pahse can still freely deposit back into the liquid, this reduces the equilibrium vapor pressure

this applies to all components in the mixture

raoults law

This is a rescaling of the pure material’s vapor pressure by the fraction of molecules at the surface of that component

• Raoult’s Law applies to all components in the mixture (solvent and all volatile, non-ionic solutes

how is boiling point affected by presence of solute

because vapor pressure at each temp is reduced, the boiling points will be increased for each pressure

how is melting point affected by the presence of solute

Freezing points are reduced, because there’s an additional penalty to displace the solute molecules to form a crystal from a solution, and we need the crystal to be more favorable (lower temperature) to overcome that barrier

osmotic pressure

property of a solution

it is entropically favorable for the solution to remain homogenous, meaning that there are not regions more or less concentrated

When there is a semi-permeable membrane (only solvent can pass through), the solvent will try to distribute itself so that the solution has as close to the same concentration as possible on either side of the membrane

how can we use colligative properties to help identify unknown materials

Because colligative properties of solutions depend only on how much solute is present, we can use measurements of a solution to identify how many moles of the solute were added to the solution

• If we also know the mass of solute used to make the solution, this tells us the molar mass of the compound