1. حرارة وديناميكا حرارية المحاضرة الثالثة د/عبدالرحمن لاشين قسم الفيزياء - كلية العلوم التطبيقية – جامعة أم القرى - المملكة العربية السعودية قسم الفيزياء - كلية العلوم – جامعة المنصورة – جمهورة مصر العربية

2. Three basic assumptions of the kinetic theory as it applies to gases: 1. Gas is composed of particles- usually molecules or atoms – Small, hard spheres – Insignificant volume; relatively far apart from each other – No attraction or repulsion between particles Kinetic Theory

3. 2. Particles in a gas move rapidly in constant random motion – Move in straight paths, changing direction only when colliding with one another or other objects – Average speed of O 2 in air at 20 o C is an amazing 1660 km/h! (1.6km=1mile) Kinetic Theory

4. 3. Collisions are perfectly elastic- meaning kinetic energy is transferred without loss from one particle to another- the total kinetic energy remains constant Kinetic Theory

5. Kinetic Theory of the Ideal Gas 1.The gas consists of a very large number of molecules, each with mass but with negligible size and no internal structure. 2.The molecules do not exert any force on each other except during a collision. This means that there is no potential energy. 3.The molecules are moving in random directions with a distribution of speeds that is independent of direction. 4.Collision with each other and the walls are elastic.

6. Pressure and Temperature Pressure: Results from collisions of molecules on the surface Pressure: Force Area Force: Rate of momentum given to the surface

7. Kinetic Theory Kinetic Theory of the Ideal Gas So we need the change in momentum

8. Kinetic Theory Kinetic Theory of the Ideal Gas Now using this in the pressure we find Thus, we find for the force on the wall

9. Kinetic Theory Kinetic Theory of the Ideal Gas But now

10. Kinetic Theory Kinetic Theory of the Ideal Gas

11. Van der Waals Equation The ideal gas equation is based upon the model that: –Gases are composed of particles so small compared to the volume of the gas that they can be considered to be zero-volume points in space. –There are no interactions, attractive or repulsive, between the individual gas particles. In 1873, Johannes H. Van der Waals (1837-1923), a Dutch physicist, developed an approximate : –A semiempirical equation, based upon experimental evidence, as well as thermodynamic arguments. –Awarded 1910 Nobel Prize in Physics for his work.

12. Van der Waals Equation Assuming that the particles behave as small, impenetrable hard spheres. If b is the volume of one mole of gas molecules n is the number of moles nb is the volume of the molecules The actual volume is given by:

13. Van der Waals Equation Correcting for the volume of the particles into the ideal gas equation: Since Taking into account the actual volume of the gas particle results in an increase in pressure relative to that predicted by the ideal gas law.

14. Van der Waals Equation the weak long-range attraction: the long- range attractive forces between the molecules tend to keep them closer together; these forces have the same effect as an additional compression of the gas. The attractive forces are proportional to the number density of the gas molecules in the container:

15. Van der Waals Equation The pressure produced by the attractive forces is proportion to the square of the number density. Taking this into account, the corrected pressure is:. The van der Waals Equation Correction for molecular volume Correction for molecular attraction

16. Van der Waals Equation The van der Waals constants: –a Pressure correction Represents the magnitude of attractive forces between gas particles Does not specify any physical origin to these forces –b Volume correction Related to the size of the particles These constants: –Are unique to each type of gas. –Are not related to any specific molecular properties.

17. Substancea’ (J. m 3 /mol 2 ) b’ (x10 -5 m 3 /mol) P c (MPa) T c (K) Air.13583.643.77133 K Carbon Dioxide (CO 2 ).36434.277.39304.2 K Nitrogen (N 2 ).13613.853.39126.2 K Hydrogen (H 2 ).02472.651.3033.2 K Water (H 2 O).55073.0422.09647.3 K Ammonia (NH 3 ).42333.7311.28406 K Helium (He).003412.340.235.2 K Freon (CCl 2 F 2 )1.0789.984.12385 K

18. van der Waals (1837-1923), Dutch

19. The critical constants of real gas critical temperature Tc Critical pressure p c Critical volume V c

20. Critical point 1.4 Molecular interactions

21. Critical Constants For the van der Waals Equation The critical constants are related to the van der Waals coefficients. The critical constants are found by solving for the critical point, which occurs with both the slope and curvature of the van der Waals isotherms are zero. The slope is found from taking the first derivative of the pressure with respect to the molar volume (i.e. n=1): The curvature is found from taking the second derivative of the pressure with respect to the molar volume:

22. Critical Constants For the van der Waals Equation Solving these two equations in two unknowns (temperature and molar volume) gives the critical temperature and critical molar volume: أسهل طريقة لحل المعادلتان بالقسمة تحصل على قيمة V c ثم بالتعويض فى المعادلة الأولى للحصول على قيمة T C The critical pressure may be calculated by substituting the expressions for the critical temperature and critical molar volume into the van der Waals equation:

23. Example The critical temperature, T c, and critical pressure, p c, for methane are 191K and 46.4×10 5 Pa respectively. Calculate the van der Waals constants and estimate the radius of a methane molecule. Solution

24. Boyle’s temperature The Boyle temperature T B is defined as:

25. At very low pressure b may ignored with respect to V

26. Compressibility Factor Z Compressibility Z is a measure of the deviation from ideal gas Z = 1 ideal gas Z can be either positive or negative For real gas:

27. We have critical compression factor Z c : For most gases, the values of Z c are approximately constant, at 0.26 ～ 0.31. Example 15, 16