Chem Notes
Chapter 11:
Solutions – homogenous systems that contain two or more substances
Solutions are not only solids in liquids but can also be solids in solids
Ex: gold in silver
Solvent – major component of the solution
Regarded as the carrier for the solute, which can participate in chemical reactions
in the solution or leave the solution through precipitation or evaporation
Solute – minor component of the solution
Composition – amount of solute in the solution
Formed by mixing two or more pure substances whose molecules interact directly
in the mixed state
Chemical reactions are usually carried out in solution and their description requires
extensions of the rules of stoichiometry
Like pure substances, solutions can be in phase equilibrium with gases, solids, and liquids
Show interesting effects that depend on the molecular weight of the solute
Mass percentage: percentage by mass of a given substance
Most commonly represented form is by mole fraction
Mole fraction – number of moles of that substance divided by the total number of moles
present
When containing n moles of species 1 and n moles of species 2 mole fractions
are:
X1= N1/N1+N2
X2= N2/N1+N2 = 1X1
MOLE FRACTION MUST TOTAL 1!
Terms involving species 1 are assigned to the solvents while terms involving species 2
are assigned to the solutes
Concentration – number of moles per unit volume
Use molarity moles solute/ liters of solution =mol L^1 = M (molar)
If a solution is heated or cooled, molarity changes as the volume of the substance
changes when heated or cooled
Molality – ratio of two masses and does not depend on temperature
= number of moles of solute per kg of solvent
moles solute/kg solvent = mol kg^1
Concentration – moles of solute/final solution volume = (molarity*initial solution
volume)/final solution volume = MVi/Vf
In the formation of a solution the attractions among the particles in the original phases
(solvent to solvent) and (solute to solute) attractions are broken up and replaced by
solvent to solute attractions
A solution has its components present in variable proportions and cannot be represented
by a chemical formula
Equations for simple dissolution reactions do not include the solvent as a
reactant
They indicate the original state of the solute in parentheses on the left side of the equation
and identify the solvent in parentheses on the right side Solute and solvent can be any combination of solid, liquid, gas – liquid water is
indisputably the best known and most important solvent
Molecular substances that have polar molecules are readily dissolved by water
Examples are the sugars (have formula C (H m) 2 n
Despite their general formula, the sugars do not contain water molecules but
they do include polar OH groups bonded to carbon atoms, which provide sites for
hydrogen bonding interactions with water molecules
replace the solute – solute interactions and the individual aquated sugar
molecules move off into the solution
many other molecular substances follow the same pattern provided they are
sufficiently polar
non polar substances such as carbon tetrachloride, octane, and the common
oils and waxes – don’t dissolve significantly in water
Solubility – the maximum mass that can be dissolved in 1 L at 25C
The dissolution of ionic species occurs through the iondipole forces
Each positive ion in solution is surrounded by water molecules oriented with the
negative end of their dipole moments toward the positive ion
When a halide (such as KCl) is dissolved, the anion forms a hydrogen bond with one of
the H atoms in a water molecule that places the atoms 0—H—Cl nearly in a straight line
Each ion dissolved in water and its surrounding solvation shell of water molecules
constitute an entity held together by ion dipole forces or by hydrogen bonds
These solvated ions can move as intact entities when an electric field is applied
Since the resulting solution is a conductor of electricity, ionic species such as
K SO are called ELECTROLYTES
2 4
Ions of both charges must be present to maintain overall charge neutrality
Precipitation reaction – ionic exchange where two anions exchange places
Spectator ions – ensure charge neutrality but do not take part directly in the chemical
reaction
Omitting these spectator ions from the balanced chemical equation results in the net ionic
equation
Net ionic equation – includes only the ions and molecules that actually take part in the
reaction
In dissolution and precipitation reactions, ions retain identities and in particular, oxidation
states do not change
Ions simply exchange the positions they had in a solid surrounded by other ions for new
positions in solution surrounded by solvent molecules
Undergo the reverse process in precipitation
Acid – substance that when dissolved in water, increases the number of hydronium ions
over the number present in pure water
Base – a substance that when dissolved increases the number of hydroxide ions over the
number present in pure water
When an acidic solution is mixed with a basic solution, a neutralization reaction
occurs
In most acid base reactions, there is no sharp color change at the end point
Ex: A volume of 50 mL is measured out and titrated with a solution of 1.306 M NaOH.
31.66mL of that titrant is required to reach the dye end point. Calculate the concentration
of the acid. Mol/L
Number of moles = V of NaOH * Concentration (.03166 * 1.306M= 4.135 x 10^2 mol
NaOH
Since 1 mol reacts with 1 mol of OH the number of moles originally present must also
have been 4.135 x 10^2 mol. Concentration would then be:
4.135 x 10^2/.05L =.827 M
Oxidationreduction (redox) reactions – electrons are transferred between reacting
species as they combine to form products
Exchange is described as a change in the oxidation number of the reactants
Oxidation number of the species giving up electrons increases whereas that or the
species accepting the electrons decreases
A prototype redox reaction is that of magnesium with oxygen, when completed,
result is magnesium oxide
2Mg(s) + O2(g) ▯2MgO(s)
Magnesium is oxidized as it gives up electrons as its oxidation number increases from 0
(in elemental Mg to +2 in Mg)
Oxygen which accepts the electrons is said to be reduced as its oxidation number
decreases from 0 to 2
Transfer of electrons:
2Mg + O2 ▯ 2MgO
(2 x 2e) 2+ 2
| | /\
\/ | |
ARROWS POINT AWAY FROM THE SPECIES GIVING UP ELECTRONS AND
TOWARDS THE SPECIES ACCEPTING ELECTRONS
Oxidation – describes the process in which the oxidation number of a species increases
even if oxygen is not involved in the reaction
Calcium Chloride
Ca(s) + Cl2(g) ▯CaCl2
(2e) (2*1e) +2 1(2)
| | | |
\/ \/
Oxidation reactions are among the most common in chemistry – seen in combustion of
coal, natural gas, and gasoline for heat and power
Phase Equilibrium: Nonvolatile Solutes
Solution made by dissolving a nonvolatile solute in a solvent
Nonvolatile – vapor pressure of the solute above the solution is negligible
Ex: solution of sucrose in water in which the vapor pressure of sucrose above the
solution is 0
The solvent vapor pressure is not zero and changes with the composition of the solution
at a fixed temperature **If the mole fraction of solvent (X1) is 1, then the vapor pressure is P1(standard), the
vapor pressure of the pure solvent at the temperature of the experiment**
When X1 approaches 0 (giving pure solute) the vapor pressure P1 of the solution must
also go to 0 because solvent is no longer present.
As the mole fraction X1 changes from 1 to 0, P1 drops from P1(standard) to 0
For solutions that conform to the straight line relationship between mole fraction and
pressure follows the equation P1=X1P1(standard) – Raoults law
Such solutions are known as Ideal solutions and other solutions that deviate from this
equation are known as nonideal solutions
They may show positive deviations (with vapor pressure higher than those
predicted by Raoults Law) or negative deviations (with lower vapor pressures)
On a molecular level, negative deviations arise when the solute attracts solvent
molecules which reduces their tendency to escape in the vapor phase
Positive deviations arise in the opposite case when solvent and solute molecules
are not strongly attracted to each other. Even nonideal solutions with nondissociating
solutes approach Raoults law as X1 approaches 1, just as all real gases obey the ideal gas
law at low densities
Raoults law forms the basis for four properties of dilute solutions, called colligative
properties – depend on the collective effect of the number of dissolved particles rather
than on the nature of the particular particles involved
1. Lowering of the vapor pressure of a solution relative to pure solvent
2. elevation of the boiling point of solution relative to the pure solvent
3. depression of the freezing point of a solution relative to the pure solvent
4. phenomenon of osmotic pressure
Vapor pressure lowering
X1= 1 – X2
Difference in vapor pressure of the pure solvent and the solution is proportional to the
mole fraction of the solute
Negative sign implies vapor pressure lowering – vapor pressure is always less above
a dilute solution than it is above the pure solvent
Boiling point elevation
ΔT = Km(molality = moles solute/mass solvent)
Freezing Point depression
ΔT=Km(molality = moles solute/mass solvent)
Osmotic Pressure
OP=pgh
P= density
G= gravity (9.807)
H= height (in meters)
Osmotic Pressure and Concentration
OP= cRT
C= Concentration
R= gas constant (.08206)
T= absolute temperature (in Kelvin)
OP(V)=nRT V= volume
N= number of moles
R= gas constant (.08026)
T= temperature in K
Chapter 12:
Systems part of the universe of immediate interest in a particular experiment or study
Closed system – boundaries prevent the flow of matter into or out of it (impermeable)
Ex: heating and cooling a metal object, has diathermal and nonrigid walls
Open system – boundaries permit flow of matter into or out of it
Isolated System – exchanges neither matter nor energy with the rest of the universe
Rigid walls – prevent the system from gaining energy by mechanical process such as
compression and deformation
Non rigid walls – prevent mechanical energy transfer
Adiabatic walls – prevent the system from gaining or losing thermal energy
Diathermal walls – permit thermal energy transfer
Most chemical reactions are modeled as open system (matter is exchanged) with
diathermal (thermal energy is transferred) and non rigid walls(density of the matter
may change during the reaction
System part of the universe that is left to exchange energy with the rest of the system
provide the external forces that cause changes in the properties of the system during a
process
Two types of properties – extensive and intensive
Extensive property – can be written as the sum of the corresponding property in the two
subsystems
Ex: volume, mass, and energy
Intensive property – same as the corresponding property of each of the subsystems
Ex: temperature, pressure (if the temperature of a system is 298K and you break it
in half, the two halves will both have a temperature of 298K)
Thermodynamic state – a macroscopic condition of a system in which the properties of
the system are held at selected fixed values independent of time
Equilibrium – after the system has been prepared by establishing a set of constraints in
the surroundings, after all disturbances caused by the preparation cease and none of its
properties change with time, the system is said to be in equilibrium
Thermodynamic process – changes the thermodynamic state of a system
may be physical such as changing the pressure of a gaseous system or boiling
a liquid
may be chemical such as it involves a chemical reaction such as the
decomposition of solid CaCO3 at 900K and 1 atm to give solid CaO and
gaseous CO2 at the same temperature and pressure
BECAUSE A PROCESS CHANGES THE STATE OF EQUILIBRIUM,
PROCESS MUST START WITH THE SYSTEM IN A PARTICULAR
EQUILIBRIUM STATE AND MUST ALSO END WITH THE SYSTEM
IN A PARTICULAR EQUILIBRIUM STATE Irreversible state – cannot be represented as a path on a thermodynamic surface because
the intermediate stages are not thermodynamic equilibrium states and thus do not
correspond to points on the equation of state surface
Ex: gas confined by a piston in which the piston is then pulled out to increase the
volume to V2 and chaotic gas currents arise as the molecules begin to move in to the
larger volume. Intermediate stages are not thermodynamic states as such properties as
density and temperature are changed rapidly in time. Eventually the currents cease and
the system approach a new equilibrium thermodynamic state. Conditions between cannot
be described by only a few macroscopic variables and is no thermodynamic then
Reversible state – proceeds through a continuous series of thermodynamic states, and
thus can be shown as a path on the equation of state surface
Ex: if a gas is expanded slowly by slowly pulling out a piston, only a tiny change
in the force exerted from the outside is required to change the direction of motion of the
piston and begins to compress the gas. Since the final equilibrium state would only be
reached after a finite time, such a process could never occur in a finite number of times
State functions – certain properties of a system that are uniquely determined by the
thermodynamic state of the system
Δ means change – usually used in terms of changes in state functions in a thermodynamic
process
ΔV= Vfinal – Vinitial change in volume between initial and final
ΔU= Ufinal – Uinitial change in internal energy between initial and final states
Since U, V, T are state functions, Δu, Δv, and Δt depend only on the initial and final states
Work
First law of thermodynamics relates the energy change in a thermodynamic process to the
amount of work done on the system and the amount of heat transferred to the system
Work – the product of the external force on a body times the distance through which the
force acts
If a body moves in a straight line from point Ri to Rf with a constant force F, the work
done on the body is W= F(RfRi)
Since F=MA, we can rewrite Work function as W=MA(RfRi)
When acceleration is constant, it is equal to the change in velocity (VfVi)/t
W=MgΔh
This is the change in potential energy of the object showing once again that the
mechanical work done in moving a body is equal to the change in energy of the body
Pressurevolume work – results when a system is compressed or expanded under the
influence of an outside pressure
Ex: Imagine that a gas has pressure Pi and is confined in a cylinder by a
frictionless piston of cross sectional area A and negligible mass. The force exerted on the
inside face of the piston by the is F=PA because pressure is defined as force divided by
area
If there is a gas on the outer side of the piston with pressure Pext then if Pext=Pi
then the piston will experience no net force
If Pext is increased the gas will be compressed and if it is decreased the gas will
expand
W=PextAΔH The product ΔH is the volume change of the system, ΔV, so the work is:
W= PextΔV
For an expansion, ΔV>0 causing w<0 and the system does work (pushes back on the
surroundings)
For a compression, (by making Pext>Pi) work is done on the system and pushed back on
the surroundings but ΔV<0 causing w>0
If NO VOLUME CHANGE, ΔV=0, and no pressure volume work is done
Internal Energy
Potential energy between molecules includes the lattice energy of solids and the attractive
and repulsive interactions between molecules in gases and liquids.
Kinetic energy appears in the translation and the internal motions of individual molecules
Gas molecules are in a constant state of motion even when no overall gas flow is taking
place in the container; the same is true of molecules in liquids and solids
Heat
Heating gas causes it to expand which enables it to move things and do work on the
surroundings
Heat – a mean of increasing the internal energy of a system without mechanical
interaction
Thermal energy (heat) – the amount of energy transferred between two objects initially at
different temperatures
Ex: when a hot body is brought into contact with a colder body the two
temperatures change until they become equal
Calorimetry – how to measure the amount of energy transferred as heat
Use an ice calorimeter
Specific heat capacity – the amount of heat required to increase the temperature of a 1g
mass by 1 degree C
If twice as much heat is transferred than the resulting temperature will change
twice as much
Q= McΔT
Q is the heat transferred to a body of mass M with specific heat capacity C to cause a
temperature change of ΔT
Since heat, like work, is energy in the process of being transferred, the appropriate unit
for it is joules
One calorie was defined as the amount of heat required to increase the temperature of 1 g
of water from 14.5c to 15.5c
1 calorie = 4.184 joules
If the energy change is caused by mechanical contact of the system with its
surroundings work is done while if it is caused by thermal contact heat is
transferred
In many cases, both heat and energy cross the boundary of a system and the change in
internal energy ΔU is the sum of the two combinations
ΔU= q +w
In any process, the heat added to the system is removed from the surroundings meaning:
Qsystem= Qsurrounding
The same thing occurs for work: Wsystem= Wsurrounding
By invoking the first law: ΔUsystem = ΔUsurrounding Heat Capacity:
Heat capacity C is defined as he amount of energy that must be added to the system to
increase its temperature by 1K
Q=C ΔT
Two different heat capacity functions:
Cp and Cv
Cp = the heat capacity at constant pressure
Cv = the heat capacity at
More
Less
