# CHM135H1 Study Guide - Final Guide: Magnesium Hydroxide, Pipette, Dont

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School
Department
Course
Professor Sublimation/deposition (solid to gas/gas to solid)
Vapouriation/Evaporation (liquid to gas/gas to liquid)
Melting/Freezing (solid to liquid/liquid to solid)
Essential Ideas
5.1 The enthalpy change of a reaction can be calculated from their effect on the temperature of their surroundings
5.2 First Law of Thermodynamics: Energy cannot be created nor destroyed
5.3 Energy is absorbed (endothermic) when bonds are broken and released (exothermic) when bonds are formed
15.1 The concept of energy change in a single step being equivalent to the summation of smaller steps can be applied to
changes involving ionic compounds (Hess’ Law)
15.2 Second Law of Thermodynamics: A reaction is spontaneous if the overall transformation leads to an increase of
total entropy (system + surroundings). The direction of spontaneous change always increases the total entropy of the
universe, at the expense energy available to do useful work
Measuring Energy Changes
Energy: A measure of the ability to do work, that is to move an object against an opposing force
Heat: A measure of the total energy (in J) and is dependent on the mass of substance present
Temperature: A measure of average KE (in or K) of the particles present (independent of mass)
Absolute Zero: -273 → when the movement of all particles has stopped
Specific heat capacity: The heat energy required to raise the temperature of 1g of substance by 1
System: The area of interest
Surroundings: Theoretically, everything else in the universe
Open system: Able to exchange energy and matter with the surroundings
Closed system: Only able to exchange energy, not matter
Isolated system: No energy or matter is able to be exchanged
Enthalpy: The heat content of a system (stored in chemical bonds and intermolecular forces in the form of potential energy)
Standard conditions: 100 kPa pressure, concentration of one mol/d , and allm3
substances are in their standard state (temperature is usually given at 298 K)
Enthalpy of formation: The change in the heat content of a system when one mole of a substance is formed from its elements in
standard conditions
Standard enthalpy change of combustion: the change in the heat content of a system for the complete combustion of one mole
of a substance in excess oxygen under standard conditions
Exothermic: Combustion and neutralization
Endothermic:
Calorimetry: The process of measuring the transfer of heat energy in a system
1. No heat is transferred between the calorimeter and the outside environment (isolated)
2. Any heat absorbed by the calorimeter (the equipment) is negligible
3. Any dilute aqueous solution is assumed to have the same physical properties of water: density (1g/mL) and specific
heat capacity (4.184 J/g)
Bond dissociation energy: The energy required to break one mole of any bond, from a gaseous molecule, under standard
conditions and pressure, to form individual atoms
One of the largest source of error is that heat is lost to the environment. For example, in a reaction between zinc and aqueous
copper sulfate, heat is lost as soon as the system’s temperature is higher than the surroundings. To determine molar enthalpy, the
limiting reagent needs to be identified.
11. Explain the meaning of the term change in enthalpy and describe how it is measured.
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Def: The enthalpy change for a reaction that is carried out in a series of steps is equal to the sum of the enthalpy change of the
individual steps, given that the starting conditions, final conditions, reactants, and products are the same.
Bond Enthalpies
Average bond enthalpy: the energy needed to break one mole of a bond in a gaseous molecule averaged over similar compounds
May be inaccurate because they do not take into consideration intermolecular forces
Ozone Depletion
Both Oxygen and Ozone play a role in protecting life against harmful UV radiation
Oxygen: one double bond
Ozone: one double bond and one single bond
Specific wavelengths of UV radiation breaks the bonds. Since oxygens double bond is stronger than ozones 1.5 bond,
radiation of higher energy and shorter wavelength is needed
vEphoton =h=λ
hc
The bond energy of ozone is 363 kJ/mol. Calculate the wavelength of UV radiation needed to break the bond.
E63 kJ 363 000 JL photon = 3 =
JEphoton =363 000
6.02 × 1023
λ = hc
Ephoton
.63 10 Js .00 0 ms J= 6 × −34 × 3 × 1 8−1 ×363000
6.02×1023 −1
30 nm= 3
High energy UV radiation breaks the double bond of oxygen. The oxygen atoms have unpaired electrons (reactive free radicals)
and react with an oxygen molecule to form ozone (creating ozone is an exothermic reaction). Lower energy radiation breaks the
bonds of ozone. Oxygen atoms react with ozone to create two molecules of oxygen.
This cycle is significant because it absorbs the dangerous UV light and the stratosphere has become warmer
First ionization energy: The minimum energy required to remove one mole of electrons from one mole of gaseous ions
(endothermic → positive value)
First electron affinity: The enthalpy change when one mole of gaseous electrons is added to one mole of gaseous atoms
(exothermic → negative value)
Born-Haber cycle (based on Hess’ Law)
Consider the formation of NaCl(s): a(s)Cl (g) → N aCl(s)N+2
1
2
1. Na is Atomized to form one mole of gaseous ions
2. One mole of chlorine atoms is formed when ½ mol of Cl-Cl bonds break (Enthalpy of Atomization of Chlorine: AKA
bond dissociation energy)
3. Electron is removed from Na (Ionization energy of Na)
4. Electron is added to Cl (Electron affinity of Cl)
5. The gaseous ions come together to form one mole of solid NaCl (Lattice enthalpy of NaCl)
H(NaCl)H(Na)H(Na) ΔH(Cl l)H(Cl)H(NaCl)Δ Θ
Lat = Δ Θ
atom + Δ i
Θ+2
1Θ
BDE C+ Δ e
Θ− Δ f
Θ
Lattice Enthalpy
Def: The enthalpy change that occurs when one mole of a solid ionic compound is separated into gaseous ions under
standard conditions
Can be calculated by assuming the crystal is made up from perfectly spherical ions
This ionic model assumes ESFA is the only interaction between the ions
Energy needed to separate the ions depends on the product of ionic charges and sum of ionic radii
As ionic radii of one ion increases, the ionic attraction decreases
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**This is a trend, not a law. As the difference in electronegativity decreases, the covalent character of the bond
increases. Lattice enthalpy assumes only ESFA is present. The additional contribution of covalent bonding accounts for
higher than expected lattice enthalpy.
HΔΘ
lat =Knm
R+R
mn+Xn
K is a constant that depends on the geometry of the lattice and “n” and “m” are magnitudes of the charges
The theoretical lattice enthalpy is lower than the experimental one because the theoretical value assumes that only electrostatic
forces of attraction are present between the ions (ionic model). As the anion becomes larger, it becomes polarized by the cation,
and the bond gets some covalent character.
Enthalpy Change of Solution (exothermic)
Def: The enthalpy change when one mole of a solute is dissolved in a solvent to infinite dilution under standard
conditions (298 K) and pressure (100 kPa)
Ions separated by water molecules are said to be Hydrated
Hydration Enthalpy: the strength of interaction depends on the attraction between the ion and polar water molecules
Def: The enthalpy change that occurs when one mole of its constituent gaseous ions is dissolved to form an
infinitely dilute solution
As ionic radius increases, the ESFA between the ions and water molecules decreases with the increasing distance
, where A is a constantHΔΘ
hyd A
Rionic
The hydration enthalpies become more exothermic as ionic charge increases and the ionic radius decreases → hasAl3+
the most exothermic hydration enthalpy because it has the highest charge and smallest radius
, where B is a constant and n is the chargeHΔΘ
hyd Bn
Rionic
Enthalpy of solution is related to both lattice enthalpy and hydration enthalpies
First the solid is sublimed into gaseous ions. Then these gaseous ions are dissolved in water
H(NaCl)H(NaCl)Δ Θ
sol = Δ Θ
lattice +H(N a )H(Cl )Δ Θ
hyd ++ Δ Θ
hyd
+ 90 kJ/mol 24 kJ/mol 59 kJ/mol = 7 − 4 3
+ kJ/mol = 7
**Compare this value to the theoretical value (+3.88 kJ/mol). This shows the general problem when a small numerical value is
calculated from the difference of two large numerical values.
Entropy (S) and Spontaneity
Refers to the distribution of available energy among the particles. The more ways the energy can be distributed, the
higher the entropy
Gibbs free energy (G) relates the energy that can be obtained from a chemical reaction to the change in enthalpy,
change in entropy, and change in temperature
This is a convenient way to take into account both the direct entropy change resulting from the
transformation of the chemicals, and the indirect entropy change of the surroundings as a result of the
gain/loss of heat energy
Def: A thermodynamic potential that can be used to calculate the maximum of reversible work that may be
performed by a thermodynamic system at a constant temperature and pressure
Skills
1. Predict whether a change will result in increase or decrease in entropy by considering the states of the reactants and
products
2. Calculate change in entropy from given entropy values
3. Application of change in Gibbs free energy is equal to the change in enthalpy minus the change in temperature
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