Chem 155 Exam 4 Study Guide Chapter 6 – Thermochemistry ...

Chapter 6 – Thermochemistry
Chem 155 Exam 4 Study Guide
- Exothermic reactions = heat produced / given off; q is negative
Endothermic reactions = heat required / taken in; q is positive
- The First Law of Thermodynamics says that energy can be converted from one
form to another, but cannot be created or destroyed. The important implications for
thermochemistry are the following:
Energy lost by one object must be gained by another object.
∆E system + ∆E surroundings = 0, or ∆E system = -∆E surroundings
The change in energy is the sum of the heat and the work.
∆E = q + w
(Be sure q and w have the same units – J or kJ – when you add them!!!)
- In chemistry, we are usually concerned with pressure-volume work.
w = -P∆V, where ∆V = V final - V initial
If the volume increases, V final > V initial . Work is done on the surroundings by the
system (the system has to “push out” against the surroundings to expand) and w is
negative.
If the volume decreases, V initial > V final . Work is done by the surroundings on the
system (the surroundings “push in” as the system compresses) and w is positive.
- The change in enthalpy for a reaction is equivalent to the heat given off or taken
in by the reaction at constant pressure. It is usually given in J/mol or kJ/mol. The “per
mole” does not necessarily include the proper stoichiometric coefficient! In other words,
CO 2 (g) + 2H 2 O(l) → CH 4 (g) + 2O 2 (g) ∆H = 890.4 kJ/mol
says that one mole of CO 2 reacting completely with H 2 O requires 890.4 kJ, but it also
says that two moles of H 2 O reacting completely with CO 2 requires 890.4 kJ. (What
would be the enthalpy change if only one mole of H 2 O reacted?)
- For gases, the change in internal energy (∆E) equals ∆H - RT∆n, where ∆n is the
change in the number of moles of gas, i.e.
∆n = # moles of product gases - # moles of reactant gases
The type of gas does not matter, just the amount!
- Hess’s Law says we can add known enthalpies of reactions to determine an
unknown enthalpy for another reaction. When you add reactions, substances that appear
on both sides of the reaction arrow cancel. When you reverse a reaction, change the sign
of ∆H. When you multiply a reaction by a constant, multiply ∆H by the same constant.
- The fundamental equations for calorimetry are q = ms∆T (or q = C∆T) and, in
accordance with the First Law of Thermodynamics, q 1 = -q 2 (the heat lost by one

substance must be gained by another substance, since we assume the calorimeter is an
isolated system from which no energy can escape). This could be q reaction = -q calorimeter or
q metal1 = -q metal2 or q reaction = -(q calorimeter + q water ) – you must be able to identify what
component of the system is losing heat and what other component(s) will gain the heat.
Chapter 7 – Quantum Theory
- Know the meaning of the wavelength (λ) and frequency (ν). Know how the two
are related: λν = a constant, which is c = 3.00x10 8 m/s in the case of light waves. Know
the units associated with frequency (Hz, cycles/second or simply s -1 ) and wavelength (m
or nm or some other prefix in front of m).
- Know the relation between energy and frequency: E = hν where h is Planck’s
constant.
- Be able to describe the photoelectric effect and its importance in the
development of quantum theory (i.e. to explain the photoelectric effect, Einstein
postulated that light has both a wavelike and a particle-like nature, an idea previously
unheard-of).
- In fact, everything has a dual wave/particle nature, including matter. The
wavelength associated with a quantity of matter is called the de Broglie wavelength. It is
very, very small for macroscopic objects (which is why we don’t notice it in day-to-day
life) but is significant for particles like electrons, which are “smeared out” in space.
(What does this have to do with the Heisenberg Uncertainty Principle?)
- Know what atomic emission lines/spectra are, and what causes them
- Know your quantum numbers
- Know how to fill an orbital energy level diagram, and how to write the shorthand
notation for electron configuration (e.g. 1s 2 2s 2 2p 6 …)
- Diamagnetic v. paramagnetic
- Bohr’s model of the atom was able to explain the atomic spectrum of hydrogen,
but failed for multi-electron atoms. Why? (Compare Figure 7.22 on page 291 to Figure
7.23 on page 292. Note that, in 7.22, the energy levels are arranged very regularly and
predictably, corresponding to “orbit radii” in the Bohr model. In 7.23, for multi-electron
atoms, splitting of the orbital energy levels occurs and does not correspond to anything
predicted by the Bohr model. What is responsible for this splitting? What does this have
to do with orbital shape? How does this disprove the Bohr model’s theory about the
shapes of orbitals?)

Chapter 8 – Periodic Relationships
- Know the meaning of and trends for: atomic radius, ionic radius, electron
affinity, ionization energy
- Be able to write the electron configuration for ions (and remember that electrons
are first removed from the furthest out s orbital when forming transition metal cations,
not from the outermost d orbital)
- Be able to identify when two atoms/ions are isoelectronic
- Know the meaning of the term “valence electrons”