Презентация на тему Heat flow and the first law of thermodynamics. Kind of thermodynamic process. Adiabatic processes

thermodynamics.
Kind of thermodynamic process. Adiabatic processes.

contact with a neighboring system changes, we say that

there has been a heat flow into or out of the system.
An energy unit related to thermal processes is the calorie (cal), which is defined as the amount of energy transfer necessary to raise the temperature of 1 gram of water by 1 degree (from 14.5°C to 15.5°C).

in the presence of nonconservative forces. It transforms into

internal energy. For example, friction produces heating
1 cal = 4.186 J

particular sample of a substance is defined as the

amount of energy needed to raise the temperature of that sample by 1 °C.
C=Q/ΔΤ
The specific heat capacity c of a substance is the heat capacity per unit mass.
c=C/m=Q/(mΔΤ)
Specific heat is essentially a measure of how thermally insensitive a substance is to the addition of energy. The greater a material’s specific heat, the more energy must be added to a given mass of the material to cause a particular temperature change.

we can relate the energy Q transferred between a

sample of mass m and specific heat capacity c of a material and its surroundings to a temperature change ΔT as
Q=mc ΔT

varies with temperature. For example, the specific heat of

water varies by only about 1% from 0 c °C to 100 °C at atmospheric pressure. Usually such variations are negligible.

pressure
Measured values of specific heats are found to depend

on the conditions of the experiment. In general, measurements made in a constant pressure process are different from those made in a constant volume process. For solids and liquids, the difference between the two values is usually no greater than a few percent and is often neglected.

does not result in a change in emperature. This

is the case when the physical characteristics of the substance change from one form to another; such a change is called a phase change. Two common phase changes:
melting: from solid to liquid
boiling: from liquid to gas
change in the crystalline structure of a solid
All such phase changes involve a change in internal energy but no change in temperature.
The increase in internal energy in boiling, for example, is represented by the breaking of bonds between molecules in the liquid state; this bond breaking allows the molecules to move farther apart in the gaseous state, with a corresponding increase in intermolecular potential energy.

heat L:
Q=±mL
Latent heat of fusion Lf is the term

used when the phase change is from solid to liquid,
Latent heat of vaporization Lv is the term used when the phase change is from liquid to gas (the liquid “vaporizes vaporizes”).

We describe the state of a system using such

variables as pressure, volume, temperature, and internal energy. These quantities are called state variables. Macroscopic state of a system can be specified only if the system is in thermal equilibrium. When we regard a thermodynamic process we imply that all its state variables change quasi-statically, that is, slowly enough to allow the system to remain essentially in thermal equilibrium at all times.

done by the gas as its volume changes from

Vi to Vf is




The work done by a gas in a quasi-static process equals the area under the curve on a PV diagram, evaluated between the initial and final states. It depends on the path between the initial and final states.

Pf(Vf-Vi)
1) Wa< Wb as Pf < Pi
2) Wa

Wb as the coloured area in (b) case is large then the area in (a) case

in which energy can be transferred between a system

and its surroundings:
One way is work done by the system, which requires that there be a macroscopic displacement of the point of application of a force.
The other is heat, which occurs on a molecular level whenever a temperature difference exists across the boundary of the system.
Both mechanisms result in a change in the internal energy of the system and therefore usually result in measurable changes in the macroscopic variables of the system, such as the pressure, temperature, and volume of a gas.

energy ΔU of the system is equal to the

heat Q put into a system minus the work W done by the system.

ΔU= Q - W

Note: here W is with the minus sign as the work is done by the system.

case of the law of conservation of energy that

encompasses changes in internal energy and energy transfer by heat and work. It provides a connection between the microscopic and macroscopic approaches.

the system,
ΔQ - heat flow into the system.
Isobaric

(constant pressure):
W=PΔV
dQ = CpdT
Isochoric (constant volume):
ΔW = 0
ΔQ = ΔU
dQ = CVdT
Cp, CV are specific heat capacities, Cp = CV + nR, n is the number of moles.
Isothermal (constant temperature):
ΔU = 0
ΔQ = ΔW

of adiabatic process is described by formula:
PVγ = const
TVγ−1

= const
γ=CP/CV

cyclic process, the change in the internal energy must

be zero. Therefore the energy Q added to the system must equal the negative of the work W done by the system during the cycle:

ΔU = 0,
Q = W

On a PV diagram, a cyclic process appears as a closed curve. In a cyclic process, the net work done by the system per cycle, equals the area enclosed by the path representing the process on a PV diagram.