Unit 8 Flashcards
Thermodynamics
-the study of energy, work and heat. -The first law states that energy can neither be created nor destroyed but converted from one form to another.
-A substance with a higher temperature will have greater kinetic energy than a substance at a lower temperature.
The attractions between ions or bonds between atoms is a form of
potential energy
Heat transfer
the flow of energy from an object with more kinetic energy (high temperature) to one with less kinetic energy (low temperature)
Q or q
Heat energy
System vs surroundings
-In a chemical reaction, the reactants and the products are the system while -the reaction solvent (for reactions involving solutions), the container in which the reaction is taking place, and the air are all a part of the surroundings.
Potential energy// bond dissociation energy
-Potential energy is stored within chemical bonds.
-The energy that holds atoms and ions together via attractive forces is a form of potential energy.
-In order to break bonds, enough heat energy must be put into the system to overcome the attractions that hold the atoms or ions together in order to separate ions or atoms from each other.
-The amount of energy required to break bonds is called “bond dissociation energy.”
-If breaking bonds requires an input of heat energy from the surroundings, the opposite process of forming bonds will release heat energy to the surroundings.
-Thus, bond breaking is an endothermic process while bond formation is an exothermic process.
Bond breaking
Endothermic
Bond forming
Exothermic
Energy when breaking and forming bonds
-Energy (E) is supplied by the surroundings to break bonds during a reaction (the system)
-Energy is released to the surroundings, from the system, during bond formation.
-If more energy is released to the surroundings from the process of bond formation than during bond breaking, the overall chemical reaction is exothermic.
-If less energy is released during bond formation than during bond breaking, the overall chemical reaction is endothermic.
If more energy is released to the surroundings from the process of bond formation than during bond breaking
the overall chemical reaction is exothermic
If less energy is released during bond formation than during bond breaking
the overall chemical reaction is endothermic.
Endothermic reaction
more E in than E out
Exothermic reaction
more E out than E in
enthalpy of reaction or ΔH, the change in the heat of a reaction
-The difference in potential energy between the reactants and products
-The ΔH of the reaction is simply the potential energy of the products minus the potential energy of the reactants
-The overall process is exothermic
(-ΔH) if the reactants have more potential energy than the products. -The reactants have high “potential” to react and will have lower bond dissociation energies than the products.
-A reaction is endothermic (+ΔH) if the products have more potential energy than the reactants.
[Enthalpy]
exothermic (-ΔH)
if the reactants have more potential energy than the products.
-In general, exothermic reactions are “favorable” because energy is released to the surroundings, and can be harnessed to do work. In contrast, endothermic reactions are “unfavorable” since a relatively large amount of energy must be added to these reactions for them to take place.
[Enthalpy]
endothermic (+ΔH)
if the products have more potential energy than the reactants.
-The reactants have low potential energy, and thus, low “potential” to react with one another; reactant compounds will have higher bond dissociation energies than the product compounds.
Hess’s Law
states that regardless of the pathway or number of steps to obtain a specific set of products from a given set of reactants the enthalpy of the overall reaction remains constant.
Bond dissociation energy (BDE)
E required to break bonds within a compound in the gas phase
1 mol, 0K, 1 atm
-strong bonds, HIGH BDE, low PE (not likely to react)
-weak bonds, LOW BDE, high PE (likely to react)
Heat of Formation (Hf)
Energy change during formation of 1 mole of a compound from its constituent forms
25^C, 1atm
-energy that goes out when forming bonds
Delta H = H products - H reactants
Delta H = Ebreak bonds + Eform bonds
Favorable reactions
Spontaneous
Exothermic
Endo is NOT spontaneous (except ice melting)
The process of melting is spontaneous at temperatures above the melting point of a pure solid.
-There are two factors that play a role in whether a reaction or process will be spontaneous or not:
-change in enthalpy (ΔH) and the change in entropy (ΔS).
-A spontaneous reaction is one that, once the activation energy is supplied, continues without any additional input of energy.
-A non-spontaneous reaction is one that by definition cannot undergo a net reaction.
The 2nd Law of Thermodynamics
-states that the entropy of the universe increases over time.
-In other words, the entropy of the universe never decreases.
There are two major ways to determine the change in entropy for a given process:
- phase changes—look for the phases of the reactants compared to the phases of the products.
-When solid reactants become liquid or gaseous products, the entropy in the system increases.
- Any time the phases change from solid –> liquid –> gas, the entropy of the system increases.
-The opposite is true when the phases of the reactants change from gases to become liquids and solids as product substance. - change in the number of moles from reactants to products—in the balanced equation, the side with the fewer total number of moles has less entropy than the side of the equation with the greatest total number of moles.
Thus, combination or synthesis reactions generally display a decrease in entropy. Consequently, decomposition reactions will display an increase in entropy—simply reverse the reaction above to prove the statement is true.
Gibbs free energy or DeltaG
-change in Gibbs free energy, or ΔG, of a system is a measure of spontaneity.
-When free energy is released to do work, a process is spontaneous; the value is negative since energy is being released into the universe. -When free energy is absorbed by a system, the process is not spontaneous and the value of ΔG is positive.
-Spontaneous reactions with a – ΔG are considered exergonic while non-spontaneous reactions with a + ΔG are considered endergonic reactions.
-spontaneity is a function of both the enthalpy and entropy of a reaction.
-A reaction will always be spontaneous (exergonic) when the enthalpy and entropy are both favorable (-ΔH, +ΔS).
-A reaction will always be non-spontaneous (endergonic) when both enthalpy and entropy are unfavorable (+ΔH, -ΔS). For all other reactions, the Gibbs Free energy equation can be used to determine spontaneity:
ΔG = ΔH – TΔS
Gibbs free energy equation
ΔG = ΔH – TΔS
ΔG is free energy of process
ΔH is the enthalpy change
T is Kelvin temperature
ΔS is the entropy change for the system
Exothermic reactions (-ΔH) having a decrease in entropy (-ΔS) will only be spontaneous at lower temperatures, see Table 8.3 below.