Chapter 6 Flashcards
Thermodynamics
The transformation of energy from one form to another.
*pg122
Zeroth Law of thermodynamics
If 2 systems are both in thermal equilibrium with a third system, then the two initial systems are in thermal equilibrium with one another. This law defines temperature and establishes the link between heat and temp.
*pg122
Thermal Equilibrium
When systems are in thermal equilibrium their temp must be the same. When 2 bodies with different temperatures are in contact with one another, the hear will flow from the body with the higher temperature to the lower temp to achieve equilibrium.
*pg122
First law of thermodynamics
States that the total energy of the universe is constant. Energy can be transferred but can’t be created or destroyed. An important result of this is that an isolated system has a constant energy, but when substances are in contact the energy flows. This law also establishes that work can be put into a system to increase its overall energy.
*pg 122
Conventions used in thermodynamics
to designate a starting and finishing line, we use 3 distinct designations to describe energy flow: the system, the surroundings, and the thermodynamic universe (or just universe)
*pg122
The system
What we’re looking at: a melting ice cube, a solid dissolving into water, a beating heart, etc. We define everything in terms of the system. so energy flowing in(entering) to the heat is + and flowing out(leaving) is -
*pg123
The surrounding
Everything else except for the thing we’re looking at: the table the ice is sitting on, the beaker that holds the solid and water, the chest in which the heart is. The signs are opposite from the system, so a energy entering the surroundings is - and exiting is + (this is because energy entering into surroundings is technically exiting the system, and vice versa)
*pg123
Thermodynamic universe
Combination of system and surroundings
*pg123
Enthalpy
Is a measure of the heat energy that is released or absorbed when bonds are broken and formed during a reaction that’s run at constant pressure. The symbol is H. follows:
– when a bond is formed, energy is released. ΔH < 0
– Energy must be put into a bond in order to break it. ΔH > 0
This can also be viewed as the energy stored in the chemical bonds of a compound
*pg123
Heat of reaction
Enthalpy change, ΔH
ΔH = H_prod - H_react
Exothermic vs endothermic
Exothermic: energy is released from the system, and the products are in a lower energy state than the reactants. ΔH = neg
Endothermic: Energy is absorbed into the system, and the products are in a higher energy state than the reactants. ΔH = pos
3 ways to Calculate ΔH_rxn
1) Hess’s law of heat summation
2) Standard heats of formation (ΔHº_f)
3) summation of avg bond enthalpies.
* pg125
Standard conditions
the temp is 298 K (25ºC)
Pressure is 1 atm
All liquids and solids are assumed to be pure
all solutions are considered to be of 1 M
These conditions are normally given by a º sign (ΔHº)
different from STP (standard temp and pressure)
*pg125
The standard Heat of formation (ΔHº_f)
The amount of energy required to make one mole of a compound from its constituent elements in their natural or standard state (the state in which the element originally exists)
For example: the ΔHº_f for O2 is 0 but for O is 249 Kj/mol
formula:
ΔHº =[∑Δn x ΔHº_f,(products)]− [∑Δn x ΔHº_f,(reactants)]
*pg126
Hess’s law of heat summation
States that if a reaction occurs in several steps, then the sum of the energies absorbed or given off in all the steps will be the same as that for the overall reaction. This is because enthalpy is a state function, which means that changes are independent of pathway of the reaction. therefore ΔH is independent of the pathway of the reaction.
basically: add the steps together to get the final result. to get that you can multiply the step or reverse it to cancel the intermediates out
* pg127
2 rules of Hess’s law
- If a reaction is reversed, the sign of ΔH is reversed too
- If an equation is multiplied by a coeff, then ΔH must be multiplied by that same value.
* pg127
Summation of average bond enthalpies
Because enthalpy is the energy in the bonds and we know that energy is needed to break a bond and released to form a bond we can use this info to calculate ΔH_rxn by:
ΔH_rxn = ∑ (BDE bonds broken) - ∑ (BDE bonds formed)
*pg128
The second Law of Thermodynamics
The disorder of the universe increase in a spontaneous process. For example: A bouncing ball will spontaneously come to rest but a resting ball won’t start bouncing. Can also be defined as all processes tend to run in a direction that leads to a maximum disorder.
*pg 130
Entropy
The measurement of disorder or randomness. The greater the disorder of a system, the greater its entropy. symbol: S or for change: ΔS.
*pg 130
Formula for change in entropy
ΔS = S_prod - S_react
*pg130
some predictable cases for entropy
- Liquids have more entropy than solids
- Gases have more entropy than both solids and liquids
- particles in solutions have more than undissolved sokids
- 2 moles of a substance have more than one mole
- the value of ΔS for a reverse reaction has the same magnitude as that of the forward reaction but with opposite sign: ΔS_ reverse = - ΔS_forward
- pg130
Third Law of Thermodynamics
Defines absolute zero to be a state of zero-entropy. It describes the least thermally energetic state, and therefore the lowest achievable temperature (which described by Kelvin is 0 K)
*PG 131
Absolute zero
Thermal energy is absent and only the least energetic thermodynamic state is available to the system in question. If only one state is possible, then there is no randomness to the system and S=0.
*pg131
Gibb’s free energy
The magnitude of the change in Gibbs free energy, ΔG, is the energy that is available (free) to do useful work from a chemical reaction. Uses both ΔH and ΔS.
Formula:
ΔG = ΔH - TΔS
– T is the absolute temp in K
– ΔG < 0 = spontaneous in the forward reaction
– ΔG = 0 reaction at equilibrium
– ΔG < non spontaneous in the forward reaction
– opposite in reverse reaction
*pg132
ΔG and Temperature
At low temperatures, the entropy doesn’t have much influence on the free energy, and ΔH is the dominant factor for spontaneity. But as temp inc the entropy term becomes more dominant. In general, the universe tends towards increasing disorder (positive ΔS) and stable bonds (ΔH)
*PG132 (Look at table)
Reaction energy diagrams
Plots the free energy diagram plots the free energy of the total reactions vs the conversion of reactants to products
*pg134(check book for diagrams)
Activation energy (E_a)
The extra energy the reactants required to overcome the activation barrier, and determines the kinetics of the reaction. The higher the barrier the slower the reaction will proceed towards equilibrium, and the lower the faster.
But E_a does not determine equilibrium. An eternally slow reaction (very big E_a) can have a very favourable (large) K_a
*pg134
Thermodynamics vs Kinetics
Thermodynamics predicts the spontaneity (and equilibrium) while kinetics predicts the rates. If you had a starting and finishing line, thermodynamics will tell you far you will go and kinetics will tell you how quickly you’ll get there.
*pg134
Reversibility
Reactions follow the principle of microscopic reversibility: the reverse reaction has the same magnitude for all thermodynamics value (ΔG, ΔH, and ΔS) but of the opposite sign, and the same reaction pathway but in reverse. which means you can draw the diagram for it by using the mirror image of the forward reaction. The E_a is diff for the reverse reaction though
*pg135
