Week 9.2 - Thermodynamics and Equilibria

7/28/2019 Week 9.2 - Thermodynamics and Equilibria

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Prepared by:Mrs Faraziehan Senusi

PA-A11-7CPhysical Transformation of Pure Substances

Chemical Equilibrium

Chapter 4

Thermodynamic and Equilibria

First Law of Thermodynamics

Reference: Chemistry: the Molecular Nature of Matter and Change,6th ed, 2011, Martin S. Silberberg, McGraw-Hill

Second Law of Thermodynamics

Simple Mixtures



Simple Mixtures



• A solution is a homogeneous mixture of two or moresubstances.

• The solute is the substance present in a smaller

amount, and the solvent is the substance present in a

larger amount.

• The solvent is the medium in which the solutes are

dissolved.

The Solution Process



Energy Changes and Solution Formation

• There are three energy steps in forming a solution: – separation of solvent molecules ( H 1),

– separation of solute molecules ( H 2),

– formation of solute-solvent interactions ( H 3

).



• We define the enthalpy change in the solution

process as

H soln = H 1 + H 2 + H 3

• H soln can either be positive or negative depending

on the intermolecular forces.

• If the solute-solvent attraction is stronger than the

solvent-solvent attraction and solute-solute attraction,

the solution process is favorable, or exothermic

(ΔHsoln

< 0).

• If the solute-solvent interaction is weaker than the

solvent-solvent and solute-solute interactions, then

the solution process is endothermic (ΔHsoln > 0).



• Breaking attractive intermolecular forces is always

endothermic.

• Forming attractive intermolecular forces is always

exothermic.



Solution Formation, Spontaneity, and

Disorder

• A spontaneous process occurs without outside

intervention.• When energy of the system decreases (e.g.

dropping a book and allowing it to fall to a lower

potential energy), the process is spontaneous.

• Some spontaneous processes do not involve thesystem moving to a lower energy state.



• If the process leads to a greater state of disorder,

then the process is spontaneous.

• Example: a mixture of CCl4 and C6H14 is less

ordered than the two separate liquids. Therefore,

they spontaneously mix even though H soln is very

close to zero.

• There are solutions that form by physical processes

and those by chemical processes.



Solution Formation and Chemical

Reactions

• Consider:

Ni( s) + 2HCl(aq) NiCl2(aq) + H2( g )

• Note the chemical form of the substance being

dissolved has changed (Ni NiCl2).

• When all the water is removed from the solution, no Ni is found only NiCl2·6H2O. Therefore, Ni

dissolution in HCl is a chemical process.



• Example:

NaCl( s) + H2O (l ) Na+(aq) + Cl-(aq)

• When the water is removed from the solution,

NaCl is found. Therefore, NaCl dissolution is a physical process.



Saturated Solutions

and Solubility

• Dissolve: solute + solvent solution.

• Crystallization: solution solute + solvent.

• Saturation: crystallization and dissolution are in

equilibrium.

• Solubility: amount of solute required to form a

saturated solution.

• Supersaturated: a solution formed when moresolute is dissolved than in a saturated solution.



Factors Affecting Solubility

Solute-Solvent Interaction• Polar liquids tend to dissolve in polar solvents.

• Miscible liquids: mix in any proportions.

• Immiscible liquids: do not mix.• Intermolecular forces are important: water and

ethanol are miscible because the broken hydrogen

bonds in both pure liquids are re-established in the

mixture.

• The number of carbon atoms in a chain affect

solubility: the more C atoms the less soluble in

water.



• The number of -OH groups within a moleculeincreases solubility in water.

• Generalization: “like dissolves like”.

• The more polar bonds in the molecule, the better it

dissolves in a polar solvent.

• The less polar the molecule, the less it dissolves in

a polar solvent and the better is dissolves in a non-

polar solvent.





Pressure Effects

• Solubility of a gas in a liquid is a function of the

pressure of the gas.

• The higher the pressure, the more molecules of gas

are close to the solvent and the greater the chance

of a gas molecule striking the surface and enteringthe solution.

– Therefore, the higher the pressure,

the greater the solubility.

– The lower the pressure,

the fewer molecules of gas

are close to the solvent and

the lower the solubility.



• If S g is the solubility of a gas, k is a constant, and P g

is the partial pressure of a gas, then Henry’s Lawgives:

• Carbonated beverages are bottled with a partial

pressure of CO2 > 1 atm.

• As the bottle is opened, the partial pressure of CO2

decreases and the solubility of CO2 decreases.• Therefore, bubbles of CO2 escape from solution.

g g kP S



Temperature Effects

• Experience tells us that sugar dissolves better in

warm water than cold.

• As temperature increases, solubility of solidsgenerally increases.

• Sometimes, solubility decreases as temperature

increases (e.g. Ce2(SO4)3).









Mass Percentage, ppm, and ppb

• All methods involve quantifying amount of solute

per amount of solvent (or solution).

• Generally amounts or measures are masses, molesor liters.

• Qualitatively solutions are dilute or concentrated.

• Definitions:

Ways of Expressing

Concentration

100solutionof masstotal

solutionincomponentof masscomponentof %mass



• Parts per million (ppm) can be expressed as 1 mg

of solute per kilogram of solution.

– If the density of the solution is 1g/mL, then 1 ppm = 1

mg solute per liter of solution.

• Parts per billion (ppb) are 1 g of solute per

kilogram of solution.


solutionincomponentof masscomponentof ppm


solutionincomponentof masscomponentof ppb



Mole Fraction, Molarity, and Molality

• Recall mass can be converted to moles using the

molar mass.

• We define

• Converting between molarity ( M ) and molality (m)

requires density.

solutionof molestotal

solutionincomponentof moles

componentof fractionMole

solutionof liters

solutemolesMolarity

solventof kg

solutemoles Molality, m





• Colligative means “tied together”.• Colligative properties depend on quantity of solute

molecules.

•

There are four important colligative properties of asolution that are directly proportional to the number of

solute particles present.

• They are

(1) vapor pressure lowering,

(2) boiling point elevation,

(3) freezing point depression,

(4) osmotic pressure.

Colligative Propertiescolligative properties are properties of

solutions that depend upon the ratio of the

number of solute particles to the number of solvent molecules in a solution



Lowering Vapor Pressure• A solution containing a nonvolatile liquid or a solid as a

solute always has a lower vapor pressure than the puresolvent

• Non-volatile solvents reduce the ability of the surface

solvent molecules to escape the liquid. Therefore, vapor

pressure is lowered.

• The amount of vapor pressure lowering depends on the

amount of solute.



• The vapor pressure of a liquid depends on the ease

with which the molecules are able to escape from the

surface of the liquid.

• When a solute is dissolved in a liquid, some of the

total volume of the solution is occupied by solute

molecules, and so there are fewer solvent molecules

per unit area at the surface.

• As a result, solvent molecules vaporize at a slower

rate than if no solute were present.

• The lowering of the vapor pressure of a solvent due

to the presence of nonvolatile, nonionizing solutes is

summarized by Raoult’s Law.



• Raoult’s Law: Xsolvent represents the mole fraction

of the solvent in a solution, P°solvent is the vapor

pressure of the pure solvent, and Psolvent is thevapor pressure of the solvent in the solution.

• If the solute is nonvolatile, the vapor pressure of the solution is entirely due to the vapor pressure of

the solvent, Psolution = Psolvent.

• Ideal solution: one that obeys Raoult’s law.• Raoult’s law breaks down when the solvent-solvent

and solute-solute intermolecular forces are greater

than solute-solvent intermolecular forces.





• Consider an ideal solution of two volatile

components, A and B.• The vapor pressure of each component above the

solution is proportional to its mole fraction in the

solution.

• The total vapor pressure of the solution is, by

Dalton’s Law of Partial Pressures, equal to the sum

of the vapor pressures of the two components.





• The elevation of the boiling point of a solvent caused

by the presence of a nonvolatile, nonionized solute is

proportional to the number of moles of solutedissolved in a given mass of solvent.

• It expressed as:

where ΔTb = T b(soln) – T b(solvent), K b = Molal boiling-

point-elevation constant (°C /m) , m = molality of the

solute.

m K T bb



• Based on the phase diagram,

• At 1 atm (normal boiling point of pure liquid), there is a lower

vapor pressure of the solution. Therefore, a higher

temperature is required to teach a vapor pressure of 1 atm for the solution (T b).

• Therefore the triple point - critical point curve is lowered.



Freezing Point Depression

•

The freezing point of a liquid is the temperature at which theforces of attraction among molecules are just great enough

to overcome their kinetic energies and thus cause a phase

change from the liquid to the solid state.

•

The freezing (melting) point of a substance is thetemperature at which the liquid and solid phases are in

equilibrium.

• When a solution freezes, almost pure solvent is formed first.

– Therefore, the sublimation curve for the pure solvent isthe same as for the solution.

– Therefore, the triple point occurs at a lower temperature

because of the lower vapor pressure for the solution.



• The melting-point (freezing-point) curve is a vertical line

from the triple point.

• The solution freezes at a lower temperature (T f ) than the

pure solvent.

• Decrease in freezing point (Tf = Tf (solvent) – Tf (soln)) is

directly proportional to molality ( K f is the molal freezing-

point-depression constant):

m K T f f





Osmosis• Osmosis is the spontaneous process by which the

solvent molecules pass through a semipermeablemembrane from a solution of lower concentration of

solute into a solution of higher concentration of solute.

• A semipermeable membrane (e.g., cellophane) separates

two solutions.

• Solvent molecules may pass through the membrane in

either direction, but the rate at which they pass into the

more concentrated solution is found to be greater than

the rate in the opposite direction.

• The initial difference between the two rates is directly

proportional to the difference in concentration between

the two solutions.





• Osmotic pressure, , is the pressure required to

stop osmosis:

M = molarity

• Isotonic solutions: two solutions with the same

separated by a semipermeable membrane.

• Hypotonic solutions: a solution of lower than a

hypertonic solution.

MRT

RT V

n

nRT V





• Red blood cells are surrounded by semipermeable

membranes.

• Crenation: – red blood cells placed in hypertonic solution

(relative to intracellular solution);

– there is a lower solute concentration in the cellthan the surrounding tissue;

– osmosis occurs and water passes through the

membrane out of the cell.

– The cell shrivels up.



• Hemolysis:

– red blood cells placed in a hypotonic solution;

– there is a higher solute concentration in the cell;

– osmosis occurs and water moves into the cell.

– The cell bursts.

• To prevent crenation or hemolysis,IV (intravenous) solutions must be isotonic.







• Tyndall effect: ability of a Colloid to scatter light. The

beam of light can be seen through the colloid.



Hydrophilic and Hydrophobic Colloids

• Focus on colloids in water.

• “Water loving” colloids: hydrophilic.• “Water hating” colloids: hydrophobic.

• Molecules arrange themselves so that hydrophobic portions

are oriented towards each other.

• If a large hydrophobic macromolecule (giant molecule) needsto exist in water (e.g. in a cell), hydrophobic molecules embed

themselves into the macromolecule leaving the hydrophilic

ends to interact with water .



• Typical hydrophilic groups are polar (containing C-O,

O-H, N-H bonds) or charged.

• Hydrophobic colloids need to be stabilized in water.• Adsorption: when something sticks to a surface we

say that it is adsorbed.

• If ions are adsorbed onto the surface of a colloid, the

colloids appears hydrophilic and is stabilized in water.



• Consider a small drop of oil in water.

• Solid soaps are usually sodium salts of long-chain organic

acids called fatty acids. They have a polar “head” and a

nonpolar “hydrocarbon tail.”

• Sodium stearate, a typical soap has a long hydrophobic tail

(CH3(CH2)16 -) and a small hydrophobic head (-CO2- Na+).

• The hydrophobic tail can be absorbed into the oil drop,

leaving the hydrophilic head on the surface.

• The hydrophilic heads then interact with the water and the

oil drop is stabilized in water.

• Most dirt stains on people and clothing are oil-based.

Soaps are molecules with long hydrophobic tails and

hydrophilic heads that remove dirt by stabilizing the

colloid in water.





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Week 9.2 - Thermodynamics and Equilibria

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