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The Kinetic-Molecular Theory Of Matter.  The Kinetic-Molecular Theory was developed to explain the observed properties of matter.  Since matter can.

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Presentation on theme: "The Kinetic-Molecular Theory Of Matter.  The Kinetic-Molecular Theory was developed to explain the observed properties of matter.  Since matter can."— Presentation transcript:

1 The Kinetic-Molecular Theory Of Matter

2  The Kinetic-Molecular Theory was developed to explain the observed properties of matter.  Since matter can be found in three common states (?), the theory has been tailored to fit these states.

3 3  The basic assumptions of kinetic- molecular theory are:  Postulate 1 ◦ Gases consist of discrete molecules that are relatively far apart. ◦ Gases have few intermolecular attractions. ◦ The volume of individual molecules is very small compared to the gas’s volume.  Proof - Gases are easily compressible.

4 4  Postulate 2 ◦ Gas molecules are in constant, random, straight line motion with varying velocities.  Proof - Brownian motion displays molecular motion.

5 5  Postulate 3 ◦ Gas molecules have elastic collisions with themselves and the container. ◦ Total energy is conserved during a collision.  Proof - A sealed, confined gas exhibits no pressure drop over time.

6 6  Postulate 4 ◦ The kinetic energy of the molecules is proportional to the absolute temperature. ◦ The average kinetic energies of molecules of different gases are equal at a given temperature.  Proof - Brownian motion increases as temperature increases.

7  The basis of the theory is that as the temperature increases the velocity of the gas molecules increase. This is because there kinetic energy is increasing.  Some of the properties of gases explained by the theory include:  Expansion: Gases will fill their container since they are moving rapidly and they do not attract each other. 

8  Fluidity: Since there is little attraction between the molecules they easily flow past one another. Liquids can also flow and both are referred to as fluids.  Low Density: Since there is a lot of space between gas molecules, gases have a low density.

9 9  The density of gases is much less than that of solids or liquids. Densities (g/mL) SolidLiquidGas H2OH2O0.9170.9980.000588 CCl 4 1.701.590.00503  Gas molecules must be very far apart compared to liquids and solids.

10  Compressibility: Since the molecules are far apart they can be forced closer together.  Diffusion and effusion:

11 11  Diffusion is the intermingling of gases.  Effusion is the escape of gases through tiny holes.

12  Gases that obey all the postulates of the Kinetic Molecular Theory are called Ideal gases. Most of the gases we are familiar with behave ideally at normal temperatures and presures.

13 13 Gas% by Volume N2N2 78.09 O2O2 20.94 Ar0.93 CO 2 0.03 He, Ne, Kr, Xe0.002 CH 4 0.00015 H2H2 0.00005 Composition of Dry Air

14  The main difference between a liquid and a gas is that the molecules in a liquid do attract one another. This restricts their motion so they stay in the liquid state. The molecules in a liquid have enough motion for them to flow past one another (fluid).  Some properties of liquids include:  Higher density than gases. There are more molecules (mass) in a given volume so the density goes up.

15  Relatively incompressible: The molecules in the liquid are already close together so increasing pressure will not compress the liquid.  Liquids can mix with one another, they diffuse into one another. See p334.  Liquids have the property of surface tension. The molecules at the surface of a liquid are attracted less that the molecules within the liquid. This creates a tension at the surface.  See p 335.

16  Capillary action refers to the attraction of a liquid to a solid surface. (fig. 7, p335)  Boiling and evaporation: If the molecules in a liquid possess sufficient energy it is possible for them to overcome the attraction between them and move into the gaseous state. If we supply the energy in the form of heat we can speed up the process and we say the liquid is boiling. If the attractive forces in the liquid are weak the liquid will evaporate rapidly.

17  The molecules in a solid are closely packed together giving solids a definite shape and volume. The attractive forces between molecules are slightly stronger in the solid state than in the liquid state.

18 18  Amorphous solids do not have a well ordered molecular structure. ◦ Examples of amorphous solids include waxes, glasses, asphalt. unit cells  Crystalline solids have well defined structures that consist of extended array of repeating units called unit cells. ◦ Crystalline solids display X-ray diffraction patterns which reflect the molecular structure. ◦ The Bragg equation, detailed in the textbook, describes how an X-ray diffraction pattern can be used to determine the interatomic distances in crystals.

19 19  Unit cells are the smallest repeating unit of a crystal. ◦ As an analogy, bricks are repeating units for buildings.  There are seven basic crystal systems.

20  We can think of solids as falling into two groups: ◦ Crystalline— particles are in highly ordered arrangement.

21 ◦ Amorphous—no particular order in the arrangement of particles.

22 In ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.

23 Because of the order in a crystal, we can focus on the repeating pattern of arrangement called the unit cell.


25  Diamonds are an example of a covalent- network solid in which atoms are covalently bonded to each other. ◦ They tend to be hard and have high melting points.

26  Graphite is an example of a molecular solid in which atoms are held together with van der Waals forces. ◦ They tend to be softer and have lower melting points.

27  Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces.  In metals, valence electrons are delocalized throughout the solid.

28  As we have mentioned all phase changes also involve changes in energy.

29 29  Solids

30 30  Solids  Liquids

31 31  Solids  Liquids  Gases

32 32  Change States ◦ heating ◦ cooling

33 33  Illustration of changes in state ◦ requires energy

34 34 Vapor Pressure  Vapor pressure is the pressure exerted by a liquid’s vapor on its surface at equilibrium.  Vapor Pressure (torr) and boiling point for three liquids at different temperatures. 0 o C 20 o C 30 o Cnormal boiling point diethyl ether185 442 647 36 o C ethanol12 44 7478 o C Water 5 18 32100 o C

35 35 Vapor Pressure as a function of temperature.

36 36 Boiling Points and Distillation boiling point  The boiling point is the temperature at which the liquid’s vapor pressure is equal to the applied pressure. normalboiling point  The normal boiling point is the boiling point when the pressure is exactly 1 atm.  Distillation is a method we use to separate mixtures of liquids based on their differences in boiling points.

37 37  Phase diagrams are a convenient way to display all of the different phase transitions of a substance.  This is the phase diagram for water.

38 38  Compare water’s phase diagram to carbon dioxide’s phase diagram.

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