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The Behavior of Gases. Properties of Gases Section 1.

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Presentation on theme: "The Behavior of Gases. Properties of Gases Section 1."— Presentation transcript:

1 The Behavior of Gases

2 Properties of Gases Section 1

3 Compressibility  Compressibility – a measure of how much the volume of matter decreases under pressure.  Gases are easily compressed because of the space between the particles.

4 Factors Affecting Gas Pressure  The amount of gas, volume, and temperature are factors that affect gas pressure.  Pressure (P) in kilopascals (kPa)  Volume (V) in liters (L)  Temperature (T) in Kelvin (K)  Number of moles (n) in mole (mol)

5 Amount of Gas  You can use kinetic theory to predict and explain how gases will respond to a change of conditions.  As you add more gas particles the pressure increases.

6 Amount of Gas  Once the pressure exceeds the strength of the container the container will burst.

7  Aerosol cans depend on the movement of gas from a region of high pressure to a region of low pressure.  Pushing the spray button creates an opening between the inside of the can and the outside.

8  The gas flows through the opening to the lower pressure region outside.  The movement of the gas propels the paint out of the can until the gas can no longer propel paint out.

9 Volume  You can raise the pressure exerted by a contained gas by reducing its volume.  The more a gas is compressed the greater the pressure.

10 Temperature  As a gas is heated, the temperature increases and the average kinetic energy also increases.

11  When the volume of a container is held constant and the temperature increases and the pressure increases.

12 The Gas Laws Section 2

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14 Boyle’s Law: Pressure and Volume  If the temperature is constant, as the pressure of a gas increases, the volume decreases.

15  Boyle’s law – states that for a given mass of gas at a constant temperature, the volume of the gas varies inversely with pressure. Boyle’s law

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17 Charles’s Law: Temperature and Volume  As the temperature of an enclosed gas increases, the volume increases, if the pressure is constant.

18  Charles’s law – states that the volume of a fixed mass of gas is directly proportional to its Kelvin temperature if the pressure is kept constant. Charles’s law

19 How can you tell from the picture that there is a fixed amount of gas in the cylinder? Describe what is happening in the cylinder as it’s being heated.

20 Gay-Lussac’s Law: Pressure and Temperature  As the temperature of an enclosed gas increases, the pressure increases, if the volume is constant.

21  Gay-Lussac’s law – states that the pressure of a gas is directly proportional to the Kelvin temperature if the volume remains constant. Gay-Lussac’s law

22 PV T Boyle’s Law Gay-Lussac’s Law Charles’s Law

23 The Combined Gas Law  Combined gas law – describes the relationships among the pressure, temperature, and volume of an enclosed gas.

24  The combined gas law allows you to do calculations for situations in which only the amount of gas is constant.

25 PV T Boyle’s Law Gay-Lussac’s Law Charles’s Law

26 Ideal Gases Section 3

27 Ideal Gas Law  The combined gas law is good when the amount of gas does not change – this does not always stay constant though.

28  To calculate the number of moles of a contained gas requires an expression that contains the variable n.  The number of moles is directly proportional to the number of particles and can be introduced into the combined gas law by dividing each side by n.

29  Ideal gas constant – (R) has the value of 8.31 (LkPa)/(Kmol).  Ideal gas law – includes the variables of P, V, T, and n.  P is the pressure (units of kPa)  V is the volume (units of L)  T is the temperature (units of K)  n is the number of moles (units of mol)

30 PV=nRT Song

31 Ideal Gases and Real Gases  An ideal gas is one that follows the gas laws under all conditions of temperature and pressure.  Real gases differ most from an ideal gas at low temperatures and high pressures.

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33 Gases: Mixtures and Movements Section 4

34 Dalton’s Law  Partial pressure – the contribution of each gas in a mixture makes to the total pressure.  In a mixture of gases, the total pressure is the sum of the partial pressures of the gases.

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37  Dalton’s law of partial pressures – states that at constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases.

38 Example:  Determine the total pressure of a gas mixture that contains oxygen, nitrogen, and helium. The partial pressures are: P O2 = 20kpa, P N2 = 46.7kPa; and P He = 26.7kPa.

39 Graham’s Law  Diffusion – tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout. Diffusion

40 Bromine gas is put in a cylinder and after several hours you can see how the gas has diffused.

41  Effusion – a gas escapes through a tiny hole in its container.  Gases of lower molar mass diffuse and effuse faster than gases of higher molar mass.

42 Thomas Graham’s Contribution  Scottish chemist Thomas Graham studied rates of effusion in the 1840’s  Relates to KE = ½ mv 2.  Kinetic energy of the particles (KE) is related to the mass (m) and their velocity (v).

43  Graham’s law of effusion – states that the rate of effusion of a gas is inversely proportional to the square root of the gas’s molar mass.

44 Example:  Determine the rate of effusion for helium compared to nitrogen.  This result tells me that the helium effuses/diffuses faster than the nitrogen at the same temperature.


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