Chemistry II-AP Notes
Thermochemistry & Thermodynamics
Topics: Important Equations & Constants:
I. Types of energies
II. Definition of state function - examples
III. System vs. surroundings - significance
IV. Laws of Thermodynamics
V. Heat vs. work DE = q - w
VI. Energy systems of importance in chemistry for gases, w = PDV or DnRT
1. expansion & compression of gases
2. electrical work
VII. Enthalpy - definition; function; equation @ constant P, QP = DH
VIII. Calorimetry & calorimeters then DE = DH - PDV
1. constant volume or DE = DH - DnRT
2. constant pressure if no gases are involved or Dn = 0
then DE = DH
IX. Heat of Reactions
1. heat of formation; definition of DHf for elements DHRxn = S DHproducts - S DHreactants
2. heat of synthesis / decomposition
3. heat of combustion
4. heat of neutralization
X. Hess's Law of Summation
XI. Sources of energy / problems with current sources
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XII. Entropy - definition; 2nd Law DSRxn = S DSproducts - S DSreactants
XIII. Free energy (Gibb's) - definition DGRxn = S DGproducts - S DGreactants
XIV. Gibb's-Helmholtz equation - significance of sign notations DGo = DHo - TDSo
XV. Phase changes & equilibrium if DG = 0, then DH = TDS
XVI. DG vs. DGo T = DH / DS
XVII. Reactions - solving for DS, DG
XVIII. Effect of temperature / pressure on free energy DG = DGo + RTlnQ
XIX. Relationship of DG to K DGo = -RTlnK
XX. Relationship of DG to Eo DGo = -nFEo
Introduction:
For every chemical reactions we have studied, there is always an energy change that accompanies them.
Consider the two following reactions:
1. the combustion of methane to produce carbon dioxide and water
2. the photosynthesis process in which carbon dioxide and water are combined to produce glucose as well as the by-product oxygen.
In the first reaction, a small amount of energy is added to initiate the reaction, but a large amount of energy is released as the reaction proceeds.
In the second, a large amount of energy goes into the reactants, with very little energy being released.
We are going to deal with heat flow associated with a reaction
The first reaction is EXOTHERMIC - THAT IS, ENERGY IS RELEASED FROM THE REACTION.
The second reaction in ENDOTHERMIC - THAT IS, ENERGY IS COMING INTO THE REACTION.
Most common chemical reactions that we know of are exothermic, but one familiar endothermic process (not reaction) is the melting of ice.
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Discuss types of energy forms
1. potential vs. kinetic
2. types of energy - chemical, mechanical, heat, radiant, light (EM), nuclear, electrical, (others?)
Discussion of a "state" function
- most important - a property of a system that does not depend on its origin!
- examples - specific heat, energy; note: heat and work are not state functions!
Distinguish between the system and the surroundings - point of emphasis for chemists vs. engineers
1. SYSTEM: any substance or substances under consideration; ex. 1 mole of a gas at STP
2. SURROUNDINGS: everything outside the system being studied.
-generally, though, this means the immediate area, but the system can include the universe.
Types of systems:
A. Open - system and surroundings are in contact - most common
- may exchange matter and/or energy
- lab room conditions where the room temperature and room pressure are affecting the lab reaction
B. Closed - doesn't exchange matter with its surroundings, but may exchange energy.
- calorimeters are a good example
C. Isolated - doesn't exchange either matter or energy to the surroundings
- would have to be something done in a vacuum
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IN GENERAL, THE STATE OF A SYSTEM IS IDENTIFIED IF WE KNOW:
1. all substances composing the system;
2. the quantity and physical state of each substance (gas, liq, solid[crystalline form])
3. temperature and pressure of the system.
With these state properties, we can determine the other state properties
For example, consider 1.00 mole of liquid water 25oC and 1 atm. pressure
Then we also know : Mass...... 18.0 g Density...... 1.0 g/ml Volume...... 18.0 ml Vapor pressure..... 17.5 mm Hg Heat capacity...... 18.0 cal/oC **
NOTE: All systems are specified at equilibrium
[State the 3 Laws of Thermodynamics]
We will start our discussion with reactions that occur at constant pressure. This includes the normal lab conditions - the atmospheric pressure in a lab doesn't change with us doing a lab (HOPEFULLY!)
What we will measure is the energy change (DE) as reactants become products.
The heat flow associated with any chemical reaction is related to an important property of the substances involved.
This property is called "heat content" or ENTHALPY - given the symbol H
DH = H(products) – H(reactants) = QP Note: it is “heat flow”
Dicsuss units of heat: We will mix calories and joules
In the English system: 1.0 calorie is the amount of heat needed to raise 1.0 gram of water 1oC.
For the metric, we use the Joule: 1.0 cal = 4.184 Joule
Consider what happens when we stick a lighted splint into an oxygen - hydrogen gas mixture. The elements explosively react to produce water and give off heat.
At 25oC and 1 atm. pressure, 68.3 kcal of heat are evolved to the surroundings per mole of water produced.
THUS THE ENTHALPY OF WATER IS 68.3 KCAL LESS THAN THE COMBINED ENTHALPY OF THE REACTANTS.
In an endothermic process, like the decomposition of mercuric oxide:
2 HgO(s) à 2 Hg(liq) + O2(g) DH = +21.7 kcal
The products are at a higher energy level than the reactant.
MEASUREMENTS OF HEAT FLOW IN REACTIONS LEAD ONLY TO VALUES OF DH AND NOT THE ACTUAL ENTHALPY OF THE SUBSTANCE.
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WE BELIEVE THAT ENTHALPY OR "HEAT CONTENT" OF A FIXED AMOUNT OF A PURE SUBSTANCE AT SET CONDITIONS IS A CHARACTERISTIC PROPERTY OF THAT SUBSTANCE, MUCH LIKE VOLUME.
Ex. IF YOU HAVE 1.00 GRAM OF LIQUID WATER AT 25oC and 1 atm. pressure, YOU KNOW YOU HAVE 1.00 ML.
Quantities, such a volume and enthalpy, whose values are fixed when one specifies the temperature, pressure, and the state of matter, are called State Properties.
Their magnitude depends only on the "state" of the substance - not its history.
Thus 1.00 GRAM OF LIQUID WATER AT 25oC and 1 atm. pressure has the same enthalpy whether it was formed from the reaction above, if an ice cube melted, if steam was condensed, or by some other process.
The enthalpy changes (at least slightly) when the temperature changes.
It is important to study energy changes for at least two reasons:
1. gives some insight into the structural aspects of a process;
2. it shows that the work done studying a process may explain why some processes are spontaneous and other aren't.
Thermodynamics: Gk. meaning “movement of heat”
Thermodynamics - deals with macroscopic measures (those that are visible)
This means our theories don't depend on atomic structures or forces that we cannot observe. For example, it doesn't matter whether our theory about ionic bonding is correct on the ionic level. We are measuring heat flow between measuring quantities of matter.
Because we have measurable quantities, we believe these laws to be universal.
FIRST LAW OF THERMODYNAMICS: In physical and chemical processes, energy is neither created nor destroyed.
This implies that energy can be transformed from one form to another.
We distinguish between two types of energy:
One definition of energy is the ability to do work, so we classify energy as either heat or useful work.
USEFUL WORK, THEN, INCLUDES ALL OTHER FORMS OF ENERGY WHICH ARE COMPLETELY INTERCONVERTIBLE.
INTERESTING NOTE: The book mentions that only 10 - 15% of the energy released in the combustion reaction in our engines is for movement.
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If we mix two or more bodies at different temperatures (regardless of size), they will attain an intermediate temperature. If we consider a system with its surroundings, (and we can assume the surroundings will have an infinite capacity to gain or lose heat) then the system will gradually attain the temperature of the surroundings.
WHEN A SYSTEM TAKES IN HEAT (Q) FROM THE SURROUNDINGS, WE CONSIDER THIS QUANTITY TO BE POSITIVE (increases heat content); IF THE SYSTEM GIVES UP HEAT TO THE SURROUNDINGS, THE SIGN OF Q IS NEGATIVE.
REMEMBER: THE SIGN VALUE DEALS WITH HEAT GAIN/LOSS FOR THE SYSTEM!
Q = + IF SYSTEM GAINS HEAT (ICE CUBE MELTING)
Q = - IF SYSTEM LOSES HEAT (MOST COMBUSTION REACTIONS)
NOW WE DEAL WITH OTHER ENERGY CHANGE: (WORK)
WHEN A SYSTEM DOES WORK, THE AMOUNT OF WORK IS ASSIGNED A POSITIVE VALUE.
WHEN A SYSTEM HAS WORK DONE ON IT, THE AMOUNT OF WORK IS ASSIGNED A NEGATIVE VALUE.
IT IS MOST IMPORTANT TO GET THE SIGN VALUES STRAIGHT:
Q = “+” WHEN SYSTEM ABSORBS HEAT (GAINS ENERGY)
Q = “-” " " LOSES " (LOSES " )
w = “+” " " DOES WORK (SYSTEM LOSES ENERGY)
w = “-” " " HAS WORK DONE ON IT (GAINS ENERGY)
FOR ANY SYSTEM, WE ARE DEALING WITH THE TOTAL AMOUNT OF ENERGY WITHIN A SUBSTANCE. THEN : DEtotal = DEfinal - DEinitial or DE = Q - w
The energy of chemical reactions results from the making and breaking of chemical bonds.
BOND ENERGY: defined as the energy required to make (or break) a chemical bond. This explains why some reactions are exothermic while others are endothermic. The enthalpy (heat content) deals with how much energy the products have with respect to the energy of the reactants.
Energy of substances may be the kinetic energy of molecules in motion (like gases), or a vibrational motion energy, or a translational motion energy.
The topic of entropy will be covered later. Entropy (DS) is the measure of disorder, or randomness, of a system.
THE SECOND LAW OF THERMODYNAMICS STATES THAT ALL SYSTEMS TEND TO MOVE TOWARD MAXIMUM ENTROPY.
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REASON for discussing entropy, though, is to make you see that bonds restrict this move towards disorder. They restrict movement. If, in a chemical reaction bonds can be broken, the energy that was stored up as potential energy is lost. Thus most reactions are exothermic because the substances tend to want to break their bonds and move freely. Go back to the equation: DEtotal = DEfinal - DEinitial or DE = Q - W
In chemistry the most common forms of work deal with expansion or compression of gases and electrical work. Let's discuss gas expansion.
Reminder:
The Standard Heat of Formation of the stable form of any element at 25oC and 1.0 atm. is ZERO. This makes sense if you think about it. By definition, the heat of formation is the energy involved when the elements that make up the substance are reacted to produce the substance, and since there is no reaction to produce an element, there is no energy change.
When we talk about the "heat" of formation we are actually talking about the "enthalpy" of formation.
Another review: We said that for any reaction, the change in the enthalpy is equal to the difference between the enthalpies of the products and the enthalpies of the reactants.
DH = S(DHproducts) - S (DHreactants)
Make sure to note order. Since most reactions that go naturally are exothermic, energy has a negative value for the overall reaction. This agrees with our study of matter - we know it tries to achieve the lowest energy level possible.
From the general equation above, we can talk about many different types of enthalpy changes.
Heat of reaction - most general - will cover any reaction Heat of formation
Heat of combustion (when a substance is reacted with oxygen)
Heat of neutralization (results from an acid-base reaction)
Heat of decomposition (usually the opposite of the heat of formation)
These are just a few types of enthalpy changes.
TWO IMPORTANT NOTES:
1. The values of the enthalpy of formation for a substance are experimental values and they may vary slightly from reference to reference.
2. The values in these tables are MOLAR values; if the enthalpy is calculated for a substance in a balanced equation, and the substance has a coefficient greater than one. then the amount of heat must be made proportional.
For example, in the synthesis of ammonia N2(g) + 3 H2(g) à 2 NH3(g)
(Heat of Formation reaction) DHf = 2 DHammonia - [ DHnitrogen gas + 3 DHhydrogen gas ]
DHf = 2 (-46.19 kJ/mole) - [ 0 + 3 (0) ]
DHf = -92.38 kJ for the reaction
Note: for the decomposition of ammonia, it takes 46.19 kJ for each mole of ammonia decomposed. (Opposite reaction ---> opposite sign)
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Problem: Calculate the heat of reaction for the decomposition reaction of calcium carbonate with heating.
CaCO3(s) CaO(s) + CO2(g)
D
DH = [(- 635.5 kJ) + (-393.5 kJ)] - (-1207.1 kJ) = +178.1 kJ (endothermic); requires heat to go
It is also nonspontaneous - Lucky for us or we wouldn't have the "WHITE CLIFFS" OF DOVER!
Work with the Calorimeter:
Closed system: used to measure energy changes since no matter is lost.