Thermochemistry
Thermochemistry
(required homework notes are sections 7.4, 7.5, 19.5, and 19.7)
7.4 The Nature of Energy: Key Definitions
Work: the result of a force acting through a distance
Energy: capacity to do work
When ball has energy due to motion, when it collides with another stationary ball, it does WORK on it, resulting in transfer of energy
Heat: the flow of energy caused by a temperature difference
Types of energy:
Kinetic energy: energy associated with motion of an object
Thermal energy: energy associated with the temperature of an object
type of kinetic energy bc it arises from the motions of atoms or molecules w/in a substance
Potential energy: energy associated with the position or composition of an object
Chemical energy: energy associated with the relative positions of electrons and nuclei in atoms and molecules, is also a form of potential energy
Some chemical compounds like methane in natural gas or the iron in a chemical hand warmer have potential energy, and a chem reaction can release that potential energy
Energy Conservation and Energy Transfer
Law of conservation of energy: states that energy can’t be made or destroyed, but can be transferred from one object to another, and assume different forms
Ex.) if you raise a ball, some potential energy become kinetic energy as ball faces ground
Ex.) if you release a compressed spring, potential energy becomes kinetic as spring expands out
Ex.) When iron reacts with oxygen within a chemical hand warmer, chemical energy of the iron and oxygen becomes thermal energy that increases the temp of hand and glove
System: a collection of objects that can be identified (example the chemicals in a beaker
Surroundings: everything with which the system can exchange energy
In an energy exchange, energy transfers between the system and the surroundings, if the system loses energy, the surroundings gain exact amount
Units of Energy
We can find the units of energy from the definition of kinetic energy
An object of mass m, moving at velocity v, has a kinetic energy KE given by the equation
7.3:
Internal energy (e): sum of the kinetic and potential energies of all of the particles that compose the system
State function: internal energy is this, meaning that its value depends on the state of the system, not on how the system arrived at that state
Mountain analogy: it doesn’t matter what path, twisty, curvy, straight, just on the initial and final values of height, in the case of systems, energy.
Internal energy change: delta E= E(final)- E(initial)
Difference in internal energy between products and reactants… use same formula from above but use the product energy my reactant energy
Summarizing:
If reactants have higher internal energy than the products, delta E system is negative and energy flows out of the system into the surroundings
If the reactants have a lower internal energy than the products, delta E system is positive and energy flows into the system from the surroundings
Change in internal energy (delta E) is the sum of the heat transferred (q) and the work done (w)
If work is done by the system (lose heat), then w will be negative; if work is done ON the system (heat added), then w will be positive
If the system releases heat to its surroundings, then q will be negative; if a system has heat added to it, then q will be positive
7.6: Enthalpy
Enthalpy (H): sum of a system’s internal energy and the product of its pressure and volume
DIFFERENCE between H and E…
Delta E is a measure of all the energy (heat and work) exchanged with the surroundings
Delta H is a measure of only the heat exchanged under conditions of constant pressure
Endothermic reaction: absorbs heat from its surrounding
Has a positive delta H (enthalpy) because energy flows into the system
Feels cold to the touch
Exothermic: gives off heat to its surroundings
Negative delta H (enthalpy)
Feels warm to the touch
Exothermic and Endothermic processes: a molecular view
In and exothermic reaction, some bonds break and new ones are formed, nuclei and electrons are reorganized into an arrangement with lower PE
As atoms rearrange, PE converts into thermal energy, so heat is emitted
Bond breaking absorbs energy, in exothermic, break then form releasing energy
Endothermic reaction: strong bonds break and weak ones form
Nuclei and electrons reorganize into an arrangement with higher potential energy, absorbing thermal energy in the process
7.5 Measuring Delta E for Chemical Reactions: Constant-Volume Calorimetry
Calorimetry: measure the thermal energy exchanged between the reaction (defined as system) and the surroundings by observing the change in temperature of the surroundings
Bomb calorimeter: a piece of equipment that measures delta E for combustion reactions
In a bomb calorimeter, the reaction happens in a sealed container called a bomb which makes sure that the reaction happens at constant volume
19.5 Heat Transfer and Changes in the Entropy of the Surroundings
Spontaneity criterion is an increase in the entropy of the universe
When water freezes, entropy decreases,but the process is spontaneous
When water vapor in air condenses into fog on a cold night, entropy of water decreases, but why spontaneous?
For any spontaneous process, the entropy of the universe increases (delta spontaneity universe>0)
Although the entropy of water decreases during freezing and condensation, entropy of universe must increase in order for these processes to be spontaneous
Second law of thermodynamics: the entropy of the universe must increase for a process to be spontaneous
Entropy of a system can decrease as long as the entropy of the surrounding increases by a greater amount so the overall entropy of the universe undergoes a net increase
For liquid water freezing or water water vapor condensing, the change in entropy for the system is negative
so for the entropy of the universe to be positive, the entropy of the surroundings must be positive and greater in absolute value or magnitude than the entropy of the system
But why does freezing ice or condensation of water increase the entropy of surroundings?
Because both processes are exothermic, and give off heat to surroundings
When thinking of entropy as the dispersal or randomization of energy, the release of heat energy by the system disperses that energy into the surroundings, increasing the entropy of the surroundings
Freezing water below zero celsius and the condensation of water vapor on a cold night increases entropy of the universe because the heat given to the surroundings increases entropy of surroundings to a sufficient degree to overcome the entropy decrease in water
Exothermic increases entropy of surroundings; endothermic process decreases entropy of surroundings
Creating bonds releases energy; breaking bonds absorbs energy
The temperature dependence of delta entropy (s) surroundings
Freezing water increases entropy of surroundings by dispersing heat energy into surroundings
Freezing water is not spontaneous at all temperature
Freezing water nonspontaneous above freezing point (0 celsius)
Magnitude of increase in entropy of surroundings is due to the dispersal of energy into the surroundings is TEMPERATURE DEPENDENT
Higher the temp, smaller the increase in entropy for a given amount of energy dispersed into the surroundings
If you give 1000 J of energy into surroundings that are hot, the entropy increase is small because the surroundings already have a lot of energy
If you disperse the same 1000 J of energy into surroundings that are cold, the entropy increase is large because the impact of 1000 J is great on surroundings that have little energy
Impact of heat released to surroundings by freezing water depends on temp of surroundings, higher temp = smaller impact
Decrease in entropy of the system is not overcome by the increase in entropy of the surroundings bc the magnitude of the positive entropy of surroundings is smaller at higher temps, resulting in negative entropy universe
Freezing water is not spontaneous at high temp
Units of entropy are J/K (energy units divided by temperature units)
At low temperatures, the decrease in entropy of a system is overcome by large increase in the entropy of the surroundings, which results in a positive change in universe entropy and spontaneous process
Quantifying Entropy Changes in the Surroundings
When a system exchanges heat with surroundings, it changes the entropy of the surroundings
Use same equation to quantify entropy changes in the surroundings
Change in entropy of the surroundings depends on:
The amount of heat transferred into or out of the surroundings
The temperature of the surroundings
Any heat leaving the system goes into the surroundings and vice versa (q system= -q surroundings)
Delta S (entropy) = (-change in energy of system)/Temperature
A process that emits heat into the surroundings (entropy of the system is negative) increases the entropy of the surroundings
Process that absorbs heat from the surroundings (change in entropy positive) decreases the entropy of the surroundings (negative change in entropy of surroundings)
Magnitude of change in entropy of the surroundings i proportional to the magnitude of the change in entropy of the system
Exothermic qualities become less of a determining factor for spontaneity as temp increases
19.6 Gibbs Free Energy
Change in free energy (gibbs free energy)= (change in enthalpy) - temperature(change in entropy)
Gibbs free energy is a criterion for spontaneity
Also called chemical potential because it is analogous to mechanical potential energy
Chemical systems tend toward lower gibbs free energy (lower chem potential)
Gibbs free energy is proportional to negative of change in entropy of the universe
Decrease in gibbs (negative) means spontaneous process
Increase in gibbs free energy corresponds to a nonspontaneous process
Effects of Enthalpy, Entropy, and temperature on spontaneity
Negative Enthalpy, positive entropy
Exothermic as enthalpy is less than zero, and entropy is positive…
Change in free energy is negative at all temperatures and reaction is spontaneous at all temperatures
(negative enthalpy means heat emission)
Entropy of system and surroundings increase, entropy of univ must also increase, so reaction is always spontaneous
Positive enthalpy, negative entropy
Endothermic (positive enthalpy), negative entropy
Change in free energy is positive at all temps→ non spontaneous reaction always
Heat absorbed
Entropy of system and surroundings decrease, so entropy of universe must also decrease = nonspontaneous reaction at all temps
Negative enthalpy, negative entropy
Exothermic, change in entropy negative, change in free energy depends on temperature
Spontaneous reaction at low temperatures but non spontaneous at high temperature
Because an increase in heat, means more negative entropy, making the spontaneity lower at higher temperatures
Positive enthalpy, positive entropy
Endothermic, change in entropy positive
Sign of change in free energy depends on temperature
Nonspontaneous at low temps, spontaneous at high temperatures
19.7 Entropy Changes in Chemical Reactions: Calculating delta S rxn
Defining Standard States and Standard Entropy Changes
Standard enthalpy change for a reaction: change in enthalpy for a process in which all reactants and products are in their standard states:
For a gas: pure gas at exactly 1 atm
For a liquid or solid: pure substance in most stable form at pressure of 1 atm and temp of interest
For a substance in solution: concentration of 1 M
Standard entropy change for a reaction: change in entropy for a process in which all reactants and products are in their standard states
Standard entropy change for a reaction = Entropy change of products - entropy change of reactants
Standard Molar Entropies and the Third Law of Thermodynamics
Standard molar entropies:
Standard value of zero to standard enthalpy of formation for an element in its standard state
Absolute values of enthalpy can’t be determined
Third law of thermodynamics: the entropy of a perfect crystal at absolute zero (0 K) is zero
Relative Standard Entropies: Gases, Liquids, Solids
Entropy of gas> entropy of liquid> entropy of solid
The larger the gas is, the greater its entropy
The more closely spaced energy states allow for greater dispersal of energy at a given temp, therefore a greater entropy
Relative Standard Entropies: Allotropes
Allotropes: elements that can exist in two or more forms in the same state of matter (example: allotropes of carbon…. Diamond and graphite)
Relative Standard Entropies: Molecular Complexity
Entropy generally increases w/ increasing molecular complexity
Molecules have more places to put energy than atoms
Gaseous argon less entropy than NO
More complex molecules have more entropy than simpler ones
Greater energy dispersal means greater energy
Calculating Standard Entropy Change for a Reaction:
To calculate this, subtract the standard entropies of the reactants multiplied by their stoichiometric coefficients from the standard entropies of the products multiplied by their stoichiometric coefficients in the form of an equation