Thermodynamics: Energy and Work - Chapter 5

• Chapter 5 is about thermodynamics.

Thermodynamics

• Thermodynamics is the study of energy and its transformations
• Thermochemistry studies the relationships between chemical reactions and energy changes
• Energy is the ability to do work or transfer heat
• Work is energy used to cause an object that has mass to move
• Heat is energy used to cause the temperature of an object to rise

Kinds of Energy

• Kinetic energy is the energy an object possesses by virtue of its motion, Ek=1/2mv^2
• Potential energy is energy in relation to the position to other objects
• Potential energy is considered at rest or stored energy
• It is expressed by the formula Ep = mgh
• Example: an object raised to above the surface of the Earth demonstrates potential energy
• Forces other than gravity can lead to potential energy
• Example: electrostatic forces between charged particles in chemistry
• An electron has potential energy when it is near a proton

Energy and Work

• Energy at the atomic or molecular level examined
• Example: examining how foods store energy that is released to be used as energy
• Thermal energy and how it is associated with the kinetic energy of molecules in a substance is examined
• Energy is measured in two units.
• Joule is the SI unit for energy with 1 J equaling 1 kg•m²/s²
• Because a joule is not a large amount of energy, kilojoules (kJ) are generally used
• 1000 Joules = 1 kiloJoule
• Calorie is another unit of energy
• 1 cal = 4.184 J and 1000 cal = 1 kcal=1 Cal

System and Surroundings

• The system includes the molecules of interest
• For example, hydrogen and oxygen molecules.
• The surroundings are everything else
• For example, the cylinder and piston.

Definitions of Energy

• Energy used to move an object over some distance is work
• w=Fd, where w is work, F is the force, and d is the distance over which the force is exerted
• Energy can also be transferred as heat
• Heat flows from warmer objects to cooler objects.
• Energy can be converted from one type to another
• A cyclist at the top of a hill, has potential energy
• As the cyclist coasts down the hill, potential energy is converted to kinetic energy
• At the bottom of the hill, potential energy is now kinetic energy.

The First Law of Thermodynamics

• Energy is neither created nor destroyed
• The total energy of the universe is a constant
• If the system loses energy, it must be gained by the surroundings

Internal Energy

• Internal energy is the sum of all kinetic and potential energy of all components of the system
• E = Efinal - Einitial
• If E > 0, then Efinal > Einitial, the system absorbed energy from the surroundings and this is called endergonic
• If E < 0, then Efinal < Einitial, the system released energy to the surroundings and this is called exergonic
• In a chemical reaction, the initial state of the system refers to the reactants and the final state refers to the products
• Energy gained or lost in a system can be analyzed by examining the processes that cause the changes to the system, heat and work
• When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w)
• E = q + w.

Conventions of values for q and w

• 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
• When heat is absorbed by the system from the surroundings, the process is endothermic
• When heat is released by the system into the surroundings, the process is exothermic.

State functions

• The internal energy of a system is a state function
• A state function is a property of a system that is determined by specifying its condition.
• The value of a state function does not depend on the history of the sample, only its present condition
• Change in energy E is a state function because it could have resulted from changes in work or heat
• Work (w) and heat (q) individually are not state functions because they are specific in their route of change

Enthalpy

• If a process takes place at constant pressure, the only work done is this pressure-volume work
• Heat flow during the process can be accounted for by measuring the enthalpy of the system
• Enthalpy is the internal energy plus the product of pressure and volume: H = E + PV
• When the system changes at constant pressure, the change in enthalpy, H, is
• H = (E + PV) which is written as H= E+PV
• Since E = q + w and w = −P V, we can substitute these into the enthalpy expression.
• H= E+PV yields H = (q + w) – w and H = q
• At constant pressure, the change in enthalpy is the heat gained or lost.

Endothermicity and hermicity defined

• A process is endothermic when H is positive
• A process is exothermic when H is negative

Enthalpy of Reaction

• The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants:
• Formula: H = Hproducts - Hreactants
• This quantity, H, is called the enthalpy of reaction, or the heat of reaction

Enthalpy explained

• Enthalpy is an extensive property
• H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction
• H for a reaction depends on the state of the products and the state of the reactants

Calorimetry

• Since the exact enthalpy of the reactants and products cannot be known, H is measured through calorimetry, which is the measurement of heat flow
• The amount of energy required to raise the temperature of a substance by 1 K (1 C) is its heat capacity
• Specific heat capacity, or simply specific heat, is defined as the amount of energy required to raise the temperature of 1 g of a substance by 1 K (or 1 C)
• specific heat = heat transferred/mass * temperature change
• s = q/mT

Constant Pressure Calorimetry

• By carrying out a reaction in aqueous solution in a simple calorimeter, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter
• Because the specific heat for water is well known (4.184 J/g-K), H for the reaction can be measured with this equation:
• q=m s T

Bomb Calorimetry

• Reactions can be carried out in a sealed "bomb"
• The heat absorbed (or released) by the water is a very good approximation of the enthalpy change for the reaction
• Because the volume in the bomb calorimeter is constant, what is measured is really the change in internal energy, E, not H
• For most reactions, the difference is very small

Hess's Law

• H is well known for many reactions, and it is inconvenient to measure H for every reaction
• However, H can be estimated using published H values and the properties of enthalpy
• Hess's law states that “if a reaction is carried out in a series of steps, H for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps.”
• Because H is a state function, the total enthalpy change depends only on the initial state of the reactants and the final state of the products

Enthalpies of Formation

• An enthalpy of formation, Hf, is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms
• Standard enthalpies of formation, H°, are measured under standard conditions (25 °C and 1.00 atm pressure)
• Hess's law can be used: H = nHf,products - m Hf,reactants, where n and m are the stoichiometric coefficients

Energy in foods and fuels

• Most of the fuel in the food comes from carbohydrates and fats
• The vast majority of the energy consumed in this country comes from fossil fuels