CHM1205 Lecture Notes

CHM1205: Introduction to Physical Chemistry

Course Requirements

• Contact hours:

  • 2 hours lecture weekly

  • 1 hour tutorial weekly

• Co-requisites: CHM1207

• Pre-requisites:

  • CHM1102, CHM1103

  • Algebra*

• Methods of Assessment:

  • Bi-weekly quizzes (15%)

  • In-class participation (10%)

  • 2 Tests (30%) - online

  • 2 Assignments (15%)

  • Final Exam (30%) in person

• Alternative assessment policy.

Schedule

• Lectures:

  • Thursday 8:15 – 10:10 am OR

  • Friday 7:15 – 9:10 pm

  • Coordinated by Mr. Stennard George & Ms. Rhea Benn

  • Office hours: TBA

• Tutorials:

  • One hour per week, based on programme

  • Worksheets must be done before tutorial session

Objectives of CHM 1205

• To enable the student to have some working familiarity with the physical concepts of chemical reactions.

• To enable the student to function in technical fields of endeavour with this familiarity.

• To provide the student with a foundation for further study of Physical Chemistry.

What is Physical Chemistry?

• Physical Chemistry is the study of underlying principles that govern the behavior of chemical systems.

• It is the branch of chemistry that deals with the physical properties of chemical substances.

Importance of Physical Chemistry

• Helps us understand the reason for change.

• Helps us understand why certain processes have a natural tendency to occur.

• Allows us to manipulate conditions to promote a desirable change.

• Enables us to measure and predict the rate of change and extent of change.

• Therefore, we realize the power of affecting the reason for, rate of, and extent of change.

Definition of Thermodynamics

• Thermodynamics is derived from two Greek words: thermos (heat) and dynamis (power).

• It is the physics of heat, work, enthalpy, and entropy changes in relation to the spontaneity of processes.

Zeroth Law of Thermodynamics

• If two systems, A and B, are in thermal equilibrium with a third system, C, then A and B are in thermal equilibrium with each other.

The First Law of Thermodynamics

• Energy cannot be created or destroyed, but is converted from one form to another.

• The total energy of the universe is constant!

  • $Energy<em>{universe} = Energy</em>{system} + Energy_{surroundings}$

First Law of Thermodynamics Equation

• $\Delta U = q + w$

  • Where:

    • U is internal energy

    • q is heat exchange in the system

    • w is work done on/by the system

• $w = -p\Delta V$

Sign Conventions for q and w

• The sign of q and w can change depending on the process.

  • Work done by the system: Negative (-)

  • Work done on the system: Positive (+)

  • Heat absorbed by the system (endothermic): Positive (+)

  • Heat lost by the system (exothermic): Negative (-)

Definition of Enthalpy

• Enthalpy is a measure of the heat content of a system under constant pressure.

• The symbol for enthalpy is H.

• $H = U + pV$

• We usually only quantify heat exchange between the system and surroundings, which is given by:

  • $\Delta H = \Delta U + p\Delta V$

Definition of Standard Enthalpy of Reaction

• It is the heat change that occurs at constant pressure when reactants at 298K and 1 atm are transformed to products at the same temperature and pressure.

• $\Delta H°_{reaction} = \sum m\Delta H°(products) - \sum n\Delta H°(reactants)$

• Since we can’t measure absolute enthalpy values, we use the molar enthalpies of formation.

• $\Delta H°<em>{reaction} = \sum m\Delta H°</em>f(products) - \sum n\Delta H°_f(reactants)$

The Standard Molar Enthalpy of Formation

• The standard molar enthalpy of formation is the enthalpy change when 1 mole of a compound is formed from its constituent elements at 1 atm and 298K.

• E.g., the standard molar enthalpy of formation of CO2 may be represented by:

  • $C(graphite) + O<em>2(g) \rightarrow CO</em>2(g) \quad \Delta H^0_{reaction} = -393.5 \frac{kJ}{mol}$

Standard Molar Enthalpy of Formation of Elements

• $\Delta H^0_f$ of substances are summarized in many data tables.

• By convention, $\Delta H^0_f$ of elements in their most stable allotropic forms are assigned a value of zero.

  • E.g., $\Delta H^0<em>f {O</em>2} = 0$

• This is because the equation which represents the standard molar enthalpy of formation of O2(g) looks like this:

  • $O<em>2(g) \rightarrow O</em>2(g)$

  • Since the initial and final states are the same, there is no change in enthalpy.

Standard Enthalpy of Reaction

• Consider the following reaction:

  • $aA + bB \rightarrow cC + dD$

• $\Delta H°<em>{reaction} = {c\Delta H°</em>f(C) + d\Delta H°<em>f(D)} - {a\Delta H°</em>f(A) + b\Delta H°_f(B)}$

Determining Standard Enthalpy of Reaction from $\Delta H_f$ data

• The Direct method

  • $C(graphite) + O<em>2(g) \rightarrow CO</em>2(g) \quad \Delta H^0<em>f = \Delta H^0</em>{reaction}$

  • $\Delta H^0<em>{reaction} = \Delta H^0</em>f(CO<em>2) - [\Delta H^0</em>f(O<em>2) - \Delta H^0</em>f(graphite)] = -393.5 \frac{kJ}{mol} - 0 = -393.5 \frac{kJ}{mol}$

Spontaneous Processes

• The second law of thermodynamics explains why chemical processes tend to favor one direction.

• A reaction that does occur under the specified set of conditions is said to be spontaneous and vice versa.

• Examples of spontaneous processes that we observe each day.

• Processes that occur spontaneously in one direction cannot occur spontaneously in the opposite direction under the same conditions.

• How can thermodynamics help us to predict whether a process will occur spontaneously?

Predictor of Spontaneity

• We might assume that spontaneous processes occur to decrease the energy of a system; i.e., exothermic reactions are spontaneous ($\Delta H$ = -ve).

• Indeed, a large number of exothermic reactions are spontaneous.

  • E.g., $CH<em>4(g) + 2O</em>2(g) \rightarrow CO<em>2(g) + 2H</em>2O(l) \quad \Delta H^0 = -890.4 \frac{kJ}{mol}$

• However, a number of spontaneous reactions are endothermic.

  • E.g., $H<em>2O(s) \rightarrow H</em>2O (l) \quad \Delta H^0 = 6.01kJ$

Predictor of Spontaneity (cont.)

• A better assumption, therefore, is that exothermicity favors the spontaneity of a reaction but does not guarantee it.

• In order to predict the spontaneity of a process, we need to know the changes in enthalpy AND entropy of the system.

Entropy (S)

• Entropy (represented by the symbol S) is a direct measure of the randomness/disorder of a system.

• S{solid} < S{liquid} << S_{gas}

• An ordered state has a low probability of occurring and, hence, a small entropy.

• However, a disordered state will have a high probability of occurring and, hence, a high entropy.

Disorder and Probability

• Disorder is more probable than order!

Second Law of Thermodynamics

• The Second Law of Thermodynamics is an expression of the universal law of increasing entropy.

• It states that the entropy of an isolated system which is not in equilibrium will tend to increase in a spontaneous process over time, approaching a maximum value at equilibrium.

An Equation for The Second Law of Thermodynamics

• $\Delta S<em>{universe} = \Delta S</em>{system} + \Delta S_{surroundings}$

• For a spontaneous process to occur:

  • \Delta S{universe} = \Delta S{system} + \Delta S_{surroundings} > 0

• And at equilibrium:

  • $\Delta S<em>{universe} = \Delta S</em>{system} + \Delta S_{surroundings} = 0$

Spontaneous Change

• Spontaneous change ALWAYS moves toward higher entropy!

• Hence:

  • \Delta S{universe} = \Delta S{system} + \Delta S_{surroundings} > 0

Air Conditioner Example

• It cools the air in the room, thus decreasing the entropy of the air.

• However, the heat used in operating the air conditioner always makes a bigger contribution to the entropy of the surroundings than the decrease in entropy of the air.

• Thus, the TOTAL entropy of the universe INCREASES!

Standard Entropy Change of Reaction

• $\Delta S°<em>{reaction} = \sum \Delta S°</em>{products} - \sum \Delta S°_{reactants}$

• Consider the following reaction:

  • $aA + bB \rightarrow cC + dD$

• Hence:

  • $\Delta S°<em>{reaction} = (cS°</em>C + dS°<em>D) – (aS°</em>A + bS°_B)$

Example

• $N<em>2(g) + 3H</em>2(g) \rightarrow 2NH_3(g)$

• $S°<em>{H</em>2} = 131 \frac{J}{K \cdot mol}$

• $S°<em>{NH</em>3} = 193 \frac{J}{K \cdot mol}$

• $S°<em>{N</em>2} = 192 \frac{J}{K \cdot mol}$

• What is the $\Delta S°_{reaction}$?

• $\Delta S°_{reaction} = [2 \times 193] – [192 + (3 \times 131)] = -199 \frac{J}{K \cdot mol}$

Third Law of Thermodynamics

• The entropy of a substance approaches zero as its temperature approaches absolute zero.

• The entropy of a perfect crystalline substance is zero at the absolute zero of temperature.

How is this possible?

• Absolute zero is 0K (-273.15 °C).

• At absolute zero, the particles in a substance have minimum motion.

• Hence, the particles are in a PERFECTLY ORDERED state.

• Thus, if there is NO disorder, then the entropy is ZERO.