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.
Kinetic Molecular Collision Theory
- Reactant particles must collide with each other.
- The collisions must have a minimum amount of energy associated with them.
- The particles must collide with the proper orientation for the rearrangement of atoms to take place.
Reaction rate
-How fast a reaction will occur can be determined by the amount of activation energy that must be supplied to the system vs. its temperature.
-At the same temperature, reactions with lower activation energies will usually happen faster than reactions with higher activation energies.
-The rate of a reaction, or the speed at which a reaction will occur, is expressed as either the amount of product that is formed per unit time, or the amount of reactant consumed per unit time.
-At the beginning of a chemical reaction there are only reactants in the reaction mixture. Over time, as reactants are consumed, their concentration decreases and the product concentration increases -Reaction rates are generally given in units of concentration per unit of time.
Factor affecting reaction rates
-How fast a reaction will occur depends on the concentration of the reactants at the start of the reaction.
-the reaction rate increases as the temperature of the reaction mixture increases’
-Adding a catalyst to a reaction will also speed up a reaction
Equilibrium
-the rate of forming products equals the rate of forming reactants , equilibrium is established and the concentration of reactants and products remains constant.
-Although the concentrations of both product and reactant are fluctuating as reactants are forming products and vice versa, the ratio of their concentration remains constant (dynamic equilibrium).
Equilibrium Constant Kc
-Kc is equal to the concentrations of the products raised to their coefficients from a balanced chemical reaction divided by the concentrations of the reactants raised to their coefficients from a balanced chemical reaction, as in the example below.
- The concentrations can change for aqueous solutions and gases, but not for solids and pure liquids. The concentration of pure solids and liquids is reflected by their density. The density of a pure solid or liquid is a characteristic of that pure substance and cannot be changed even during a chemical reaction
-When considering spontaneity, reactions that have a Kc greater than one will generally be spontaneous (exergonic) in the forward direction and non-spontaneous (endergonic) in the reverse direction when the reaction takes place at room temperature (25 °C and 1 atm pressure).
Catalysts do NOT affect the position of anequilibrium, butspeed up the rate of a reaction.
True
Le’Chatelier’s Principle
-Equilibrium is reestablished when the ratio of the concentration of reactants to products equals the Kc value for a given reaction at a specific temperature.
-The reaction will shift so that the ratio of the concentrations of reactants and products maintains the equilibrium constant or Kc.
Le’Chatelier’s Principle
Stressors
- Increasing or decreasing the reactant concentration
- Increasing or decreasing the product concentration
- Increasing or decreasing the temperature of the reaction mixture
- Increasing or decreasing the volume of the reaction vessel for gaseous reactions
-Endothermic reactions– Energy is considered a reactant in endothermic reactions, so increases in the reaction temperature would have the same effect of increasing the reactant concentration. The reaction would shift to the right resulting in an increase of product concentrations. If the temperature is decreased, the energy supplied to the reaction also decreases causing the equilibrium to shift to replace the missing energy resulting in increased reactant concentrations.
-Exothermic reactions– The opposite is true for exothermic reactions. Energy is a product for exothermic reactions, thus if the reaction temperature is increased, the reaction will shift to remove the excess energy towards the reactants. The concentration of reactants will increase. Decreasing the reaction temperature is analogous to removing products, so the reaction will shift to form more energy; the final result being an increase in product concentrations.
- Changes in volume of gaseous reactions. Due to Boyle’s Law (Unit 5) any increases in the pressure exerted on a gas result in a decrase in gas volume and vice versa (decreased pressure results in an increase in volume). The shift in equilibrium depends on the moles of gas on either side of the reaction arrow. When the volume of the reaction vessel is increased (decreased pressure) for a gaseous reaction, the equilibrium will shift to the side of the reaction where there are more moles of gas. When the volume of the reaction vessel decreases (pressure increases), then the reaction will shift to the side where there are fewer moles of gas.
Dissolution constants (Kd)
-Kd values are temperature dependent as all equilibrium constants are temperature dependent.
-Solutes that dissolve readily in a particular solvent will have high Kd
-those that are insoluble have a low Kd.
- Non-polar solutes will have a Kd = 0 in a polar solvent, but a high Kd in a non-polar solvent
-Polar solutes have high Kd values in polar solvents, yet the Kd is equal to 0 in non-polar solvents.
-Ionic compounds that are very soluble in water have high dissociation constants at a given temperature whereas those compounds that are slightly soluble or insoluble will have low to very low (≈ 0) dissociation constants.
-Weak electrolytes that only partially dissociate and ionize (become an ion) in solution will display a low Kd.
-Strong electrolytes that completely dissolve and ionize display a high Kd.
-The Kd of a non-electrolyte depends on whether the solute is soluble in water or not (polar or non-polar, see previous paragraph).
