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# Thermodynamics - Heat and Temperature

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### Thermodynamics - Heat and Temperature

1. 1. HEAT AND TEMPERATUREThermodynamics
2. 2. 1.0 - Introduction Heat is a form of energy. It is measured like other forms of energy in J (Joules). The above statement should have led you to realise that heat and temperature are two different things as temperature is measured in K (Kelvin) or more commonly in °C (Degree Celsius or Centigrade) or °F (Degree Fahrenheit). An object may contain various different types of energy.
3. 3. 1.1 - Introduction Cont.  Take for example a student throwing a jar containing bees. The object has the potential energy caused by the gravitational field of the earth and kinetic energy as it is moving. However, the individual bees also have their own kinetic energy - we will call this the random kinetic energy of the jar.  If we said that the jar is now a metal ball and the bees in the jar are the atoms of the metal ball, the average random kinetic energy caused by the vibration of the atoms is the temperature of the metal ball.
4. 4. 1.2 - Introduction Cont. In a substance, Kinetic Energy is present due to the masses and velocities of its particles being vibrated, rotated or translated, and Potential Energy is present due to the attractive forces between each of the particles as bonds AND between separate particles. The sum of the kinetic and potential energies of all the particles is called the internal or thermal energy of the substance. The term heat is used to describe the internal energy of a substance. The study of the transfers of this energy is called thermodynamics.
5. 5. 2.0 - Thermal Equilibrium  Thermal Equilibrium is the state at which two objects in an isolated environment gain the same temperature after the process of heat transfer from the body containing more heat (TB) to the other (TA).  Note that in an isolated environment, there is no heat lost to the surroundings – therefore, the heat lost by the hotter object (TB) is equal to the heat gained by the less hotter object (TA).  Note that this is theoretical and in practise, heat is lost through radiation (even if the experiment is conducted in space) – all objects that have temperatures above absolute zero, radiate energy in the form of electro-magnetic radiation. And when conducted on earth heat is also lost through conduction and convection (refer to Slide 5.0 – Calorimetry).
6. 6. 3.0 - Measuring Temperature There are many different scales used to measure temperature. Below are the three scales that are mainly used in the present day;  The Fahrenheit Scale  Developed by German physicist, Gabriel Fahrenheit (1686 – 1736)  In this scale, the freezing point of a salt solution is 0°F, the freezing point of pure water is 32°F, and the boiling point of pure water is 212°F.  This scale is mainly used in the US, UK and Canada.
7. 7. 3.1 - Measuring Temperature Cont.  The Celsius Scale  Developed by Andres Celsius (1701 – 1744)  In this scale, the freezing point of pure water is 0°C and the boiling point of pure water is 100°C.  The Kelvin/Absolute Scale  Developed by Lord Kelvin (1824 – 1907)  In this scale, 0 K is the absolute zero temperature – this means that at this temperature, there is absolutely no particle motion.  Note; (0 K = -273.15°C).
8. 8. 4.0 - Specific Heat Capacity In a room, that has a constant temperature (say 23°C), all the objects have the same temperature (23°C – thermal equilibrium). However, if we humans touched a metallic object in the room it would feel much more cold than a non-metallic object in the same room. This is due to the fact that metals are good conductors of heat. The heat from our bodies is conducted faster to the metals than to the non-metals. And because our body senses the rate at which heat is transferred to or away from our body, the metals feel more cold (remember that the metals still have the same temperature). “Good conductors of heat” refers to substances with a low heat capacity – i.e. they require relatively less amounts of energy to raise its temperature (refer to next slide for a more precise and detailed description of specific heat capacity).
9. 9. 4.1 - Specific Heat Capacity Cont.  Specific Heat Capacity is the measure of how much energy is required to raise the temperature of 1 kg of a substance by 1 K or 1°C (note – a change of 1 K is exactly the same as a change of 1°C).  Different substances have different specific heat capacities. On the left side is a table containing the specific heat capacities of some commonly seen substances  On the bottom left is the formula for specific heat capacity. In this formula, Q is the heat energy required (J), m is the mass (kg), c is the specific heat capacity, ΔT is the change in temperature (measured in either °C or K – refer to the first dot point in this slide).
10. 10. 5.0 - Calorimetry  When two substances are placed together in a closed system, thermal equilibrium occurs.  In practice, there is always some heat lost to the surroundings. There are two main ways in which such heat loss could be minimised;  Carrying out the experiment quickly.  Use calorimeters, which have good insulation to limit the loss of heat to the surroundings. This process is called Calorimetry.Return to Slide 2.0 – Thermal Equilibrium
11. 11. 6.0 - Change of State  The amount of energy required to melt 1 kg of an object is called the specific latent heat of fusion.  The amount of energy required to vaporise 1 kg of an object is called the specific latent heat of vaporisation.  On the left is the formula for the energy required to change the state of a substance. In this formula; Q is the heat energy required (J), m is the mass (kg), and L (specific latent heat) of the object becomes;  Lf for the specific latent heat of fusion (or)  Lv for the specific latent heat of vaporisation.
12. 12. 7.0 - Changing the melting and boilingPoints Most substances have fixed melting and boiling points as long as they are in pure form. To change the melting and boiling points of various substances, there are two main methods which could be used;  Adding Impurities to the substance (and/or)  Changing the pressure of the substance and/or its environment.
13. 13. 8.0 - Evaporation  Liquids turn into gas without boiling. <Image of clothes This process is called evaporation and occurs all the time. drying>  For a substance to change state, energy is required. But note that not all the individual particles of a substance have exactly the same energy (also note – temperature is the measure of the average random kinetic energy of the particles of a substance).  This is the reason for evaporation – individual particles with relatively higher energy are able to reach the surface of the substance and escape (e.g. when you leave a bowl of water at room temperature, it will eventually evaporate to nothing).
14. 14. 9.0 - Laws of Thermodynamics There are two laws of thermodynamics;  The first law of thermodynamics states that the total increase in the thermal energy of an isolated system is equal to the sum of the heat added to it and the work done on it. Note that this is just an extension of the law of conservation of energy.  The second law of thermodynamics relates heat transfer to differences in temperature.
15. 15. 9.1 - Laws of Thermodynamics Cont.  The laws of thermodynamics also <Image of Sand Castle helped develop a new term in – Entropy> physics, called entropy.  Entropy is the measure of the disorder of a system – the more disorder, the more entropy. It states that in nature, all ordered systems head towards becoming disordered.  An example of this (from thermodynamics) would be when two objects, say water and ice, are placed in contact with each other and allowed to reach thermal equilibrium. After equilibrium is reached, the ordered molecules of the ice become less ordered; and therefore it now has a lower entropy.
16. 16. FINALLY! WE’RE DONE!
17. 17. REMEMBER! This presentation is only designed to help you learn easier not thorough. So, refer to you textbook for detailed information on this chapter! And practice the questions in your textbook if any!