Thermochemistry Flashcards
What is thermochemistry ?
The branch of chemistry concerned with the heat effects that accompany chemical reactions
What is system?
The substances involved in the chemical and physical ch/anges under investigation
*reactants and products
What is surroundings?
Everything else in the universe
*outside the beaker where the chemicals are
Potential Energy (PE)
- Energy due to condition, position, or composition
- Associated with forces of attraction or repulsion between objects
Kinetic Energy
- Energy due to the motion of an object
- Depends on the mass of the object (m) and its velocity (v)
- KE = 1/2 mv^2
- Energy can change from potential to kinetic
Electrical Energy
Kinetic energy associated with the flow of electrical charge
Heat or Thermal Energy
kinetic energy associated with molecular motion
Nuclear Energy
potential energy in the nucleus of atoms
Light or radiant Energy
kinetic energy associated with energy transitions in an atom
Chemical Energy
Potential energy in the attachment of atoms or because of their position
Kinetic Energy
ek = 1/2mv^2
[ek] = kg (m/s)^2 = J
Work
w = F x d
= m x a x d
[w] = kg (m/s^2) m = J
a = acceleration
Law of Conservation of Energy
Energy cannot be created or destroyed
First Law of Thermodynamics
Energy can be transferred between objects
Energy can be transformed from one form to another
Heat → light → sound
Energy of the universe is constant
Units of Energy
joule (J): the amount of energy needed to move 1 kg mass a distance of 1 meter
1 J = 1 N x m = 1 kg m^2/s^2
calorie (cal): the amount of energy needed to raise one gram of water by 1 degree Celsius
kcal = energy needed to raise 1000 g of water 1 degree celsius
Food Calories = kcals
Energy Conversion Factors
1 calorie (cal) = 4.184 joules (J) (exact)
1 Calorie (Cal) = 1000 calories (cal)
1 kilowatt-hour (kWh) = 3.60 x 10^6 joules (J)
Methods of Transferring Energy
Heat (Thermal Energy): transfer of energy between two objects due to a temperature difference
- Temperature is a property that reflects random motion of particles in a substance
- An intensive property
Work: force acting over a distance
- Heat and Work
- q and w
- Energy changes
Pathway
Specific conditions that define the path by which energy is transferred
Energy change is independent of the pathway
Work and heat are dependent on the pathway
Internal Energy
The internal energy is the total amount of kinetic and potential energy a system possesses
- PE + KE = Internal energy (E or U)
The change in the internal energy of a system only depends on the amount of energy in the system at the beginning and end
- A state function is a mathematical function whose result only depends on the initial and final conditions, not on the process used
Directionality of Heat Transfer
internal energy is transferred between a system and its surroundings as a result of a temperature difference
- Temperature is a measure of the average KE of particles in solids, liquids, and gasses
Energy is transferred as the initially faster particles decrease in motion and the initially slower molecules increase in motion
- Particles will have the same average KE and thus the same temperature
Heat “flows” from hotter to colder
- Temperature may change
- Phase may change (an isothermal process)
Quantity of Heat Energy Absorbed Heat Capacity
When a system absorbs heat, its temperature increase
The increase in temperature is directly proportional to the amount of heat absorbed
The proportionality constant is called the heat capacity, C
- Units of C are J/ C or J/K q = C x ΔT
The heat capacity of an object depends on its mass
The heat capacity of an object depends on the type of material
Specific Heat Capacity
Measure of a substance’s intrinsic ability to absorb heat
The specific heat capacity is the amount of heat energy required to raise the temperature of one gram of a substance 1 degree Celsius
q = m x Cs x ΔT
Units are J/g C
The molar heat capacity is the amount of heat energy required to raise the temperature of one mole of a substance 1 degree Celsius
The rather high specific heat of water allows it to absorb a lot of heat energy without large increases in temperature
- keeping ocean shore communities and beaches cool in the summer
- allows it to be used as an effective coolant to absorb heat
Heat Capacity
The quantity of heat required to change the temperature of a system by one degree
Molar heat capacity
System is one mole of substance
Specific heat capacity, c q = mcΔT
System is one gram of substance
Heat capacity q = CΔT
(mass of system) x specific heat
Law of conservation of energy
In interaction between a system and its surroundings the total energy remains constant-energy is neither created nor destroyed
qsystem + qsurroundings = 0
qsystem = - qsurroundings
Heat of Reaction and Calorimetry
Chemical energy
Contributes to the internal energy of a system
Heat of reaction, qrxn
The quantity of heat exchanged between a system and its surroundings when a chemical reaction occurs within the system, at constant temperature
Conservation of Energy in a Chemical Reaction
The energy of the system, reactants and products increases, while the energy of the surroundings decreases
Reactant + Energy –> Product
The energy of the reactant and products decreases, while the energy of the surroundings increases
Reactant –> Product + Energy
An Exothermic Reaction
Exothermic process: any process that gives off heat transfers thermal energy from the system to the surroundings
An Endothermic Reaction
Endothermic process: any process in which heat has to be supplied to the system from the surroundings
Bomb Calorimetry
qrxn = -q cal
qcal = q bomb + q water + q wires + …
The “Coffee-Cup” Calorimeter
A simple calorimeter
well insulated and therefore insulated
measure temperature change
q rxn = - q cal
q cal = m x Cs x ΔT
Work
In addition to heat effects chemical reactions may also do work
Gas formed pushes against the atmosphere
The volume changes
Pressure -volume work
Pressure-Volume Work
w = -P ext ΔV
The First Law of Thermodynamics
Internal Energy, U
Total energy (potential and kinetic) in a system
Translational kinetic energy
Molecular rotation
Bond vibrations
Intermolecular attractions
Chemical bonds
Electrons
The first law of thermodynamics
A system contains only internal energy
A system does not contain heat or work
These only occur during a change in the system
ΔU = q + w
An isolated system is unable to exchange either heat or work with its surroundings, so that ΔU isolated system = 0, and we can say:
The energy of an isolated system is constant
q > 0
Heat is transferred from the surroundings to the system
q < 0
Heat is transferred from the system to the surroundings
w > 0
Work is done by the surroundings on the system
w < 0
Work is done by the system on the surroundings
Sign of ΔU = q + w
q > 0 and w > 0 ΔU > 0
q > 0 and w < 0: the sign of ΔU depends on the magnitudes of q and w
q < 0 and w > 0: the sign of ΔU depends on the magnitudes of q and w
q < 0 and w < 0 ΔU < 0
Functions of State
Any property that has a unique value for a specified state of a system is said to be a function of state or a state function
Obtain three different samples of water:
Purified by extensive distillation of groundwater;
Synthesized by burning pure H2(g) in pure O2(g) and
Prepared by driving off the water of hydration from CuSO4 x 5H2O and condensing the gaseous water to a liquid
Water at 293.15 K and 1.00 atm is in a specified state
d = 0.99820 g/mL
This density is a unique function of the state
It does not matter how the state was established
State functions?
properties that are determined by the state of the system, regardless of how that condition was achieved
- energy, pressure, volume, temperature
Functions of …
U is a function of state
- Not easily measured
ΔU has a unique value between two states
- Is easily measured
Imagine a system that changes from state 1 to state 2 and back to state 1
Path Dependent Functions
Changes in heat and work functions are not functions of state
Heats of Reaction: ΔU and ΔH
Reactants –> Products
ΔU = Uf - Ui
ΔU = q rxn + w
In a system at constant volume (bomb calorimeter):
ΔU = q rxn + 0 = q rxn = q v
Indirect Determination of ΔH: Hess’s Law
ΔH is an extensive property
Enthalpy change is directly proportional to the amount of substance in a system
ΔH changes sign when a process is reversed
Standards States and Standard Enthalpy Changes
Define a particular state as a standard state
Standard enthalpy of reaction, ΔH*
The enthalpy change of a reaction in which all reactants and products are in their standard states
Standard state
The pure element or compound at a pressure of 1 bar and at the temperature of interest
Standard Enthalpies of Formation ΔHf*
The enthalpy change that occurs in the formation of one mole of a substance in the standard state from the reference forms of the elements in their standard states
The standard enthalpy of formation of a pure element in its reference state is 0.
Standard Enthalpies of Formation
Standard states:
Compound
For a gas, pressure is exactly 1 bar
For a solution, concentration is exactly 1 molar
Pure substance (liquid or solid)
Element
The form [(N2(g), K(s)] in which it exist at 1 bar and 25 degree celsius
Bond Order, Length, and Energy
Bond Order
Single bond, bond order = 1
Double bond, bond order =2
Triple bond, bond order = 3
Bond length
The distance between the centers of two atoms joined by a covalent bond
Bond energy
The stability of a chemical bond
The length of the covalent bond between two atoms can be approximated as the sum of the covalent radii of the two atoms
What affects bond length?
- The smaller the principle quantum numbers of the valence orbitals, the shorter the bond
- The higher the bond multiplicity, the shorter the bond
- The higher the effective nuclear charge of the bonded atoms, the shorter the bond
- The larger the electronegativity difference, the shorter the bond
Bond Enthalpy and Bond Length
We can also measure an average bond length for different bond types
As the number of bonds between two atoms increases, the bond length decreases
Bond Energy
- Bond strength increases as more electrons are shared between the atoms
- Bond strength increases as the electronegativity difference (Δχ) between bonded atoms increases
- Bond strength decreases as bonds become longer
Covalent Bond Strength
Most simply, the strength of a bond is measured by determining how much energy is required to break the bond
This is the bond enthalpy
The bond enthalpy for a Cl-Cl bond, D(Cl-Cl), is measured to be 242 kJ/mol
Average Bond Enthalpies
This table lists the average bond enthalpies for many different types of bonds
Average bond enthalpies are positive, because bond breaking is an endothermic process
Enthalpies of Reaction
Yet another way to estimate ΔH for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed
ΔH rxn = E (bond enthalpies of bonds broken) - E (bond enthalpies of bonds formed)
Bond Energy and Enthalpy
Bond energy values can be used to calculate approximate energies for reactions:
ΔH: E D(bonds broken) - D(bonds formed)
E = sum of terms
D is the molar bond energy (always positive)
