Vocabulary for Standard 4 Gas Laws

 

 

   

Word

 

Definition

1  

3rd law of thermodynamics

 

 

2  

Absolute zero

(0 K)

The temperature at which the volume of an ideal gas becomes zero; a theoretical coldest temperature that can be approached but never reached. Absolute zero is zero on the Kelvin scale, -273.15°C on the Celsius scale, and -459.67°F on the Fahrenheit scale.

3  

Atmosphere of pressure

(atm)

A unit of pressure, equal to a barometer reading of 760 mm Hg. 1 atmosphere is 101325 pascals and 1.01325 bar.

4  

average kinetic energy

 

 

5  

Avogadro’s law

 

Equal volumes of an ideal gas contain equal numbers of molecules, if both volumes are at the same temperature and pressure. For example, 1 L of ideal gas contains twice as many molecules as 0.5 L of ideal gas at the same temperature and pressure.

6  

Barometer

 

An instrument that measures atmospheric pressure. A mercury barometer is a closed tube filled with mercury inverted in a mercury reservoir. The height of the mercury column indicates atmospheric pressure (with 1 atm = 760 mm of mercury). An aneroid barometer consists of an evacuated container with a flexible wall. When atmospheric pressure changes, the wall flexes and moves a pointer which indicates the changing pressure on a scale.

7  

Boyle’s Law

 

The pressure of a ideal gas is inversely proportional to its volume, if the temperature and amount of gas is held constant. Doubling gas pressure halves gas volume, if temperature and amount of gas don't change. If the initial pressure and volume are P1 and V1 and the final pressure and volume are P2V2, then P1V1 = P2V2 at fixed temperature and gas amount.

8  

C to K conversion

 

 

9  

Charles’ Law

 

The volume of a gas is directly proportional to its temperature in kelvins, if pressure and amount of gas remain constant. Doubling the kelvin temperature of a gas at constant pressure will double its volume. If V1 and T1 are the initial volume and temperature, the final volume and temperature ratio V2/T2 = V1/T1 if pressure and moles of gas are unchanged.

10  

Combined gas law

 

 

11        
12  

Dalton’s law of partial pressure

Dalton's law

The total pressure exerted by a mixture of gases is the sum of the pressures that each gas would exert if it were alone. For example, if dry oxygen gas at 713 torr is saturated with water vapor at 25 torr, the pressure of the wet gas is 738 torr.

13   diffusion   When particles move from areas of high concentration to areas of low concentration.  For example, if you open a bottle of ammonia on one end of the room, the concentration of ammonia molecules in the air is very high on that side of the room.  As a result, they tend to migrate across the room, which explains why you can smell it after a little while.  Be careful not to mix this up with effusion
14  

Directly proportional

 

 

15   effusion   When a gas moves through an opening into a chamber that contains no pressure.  Effusion is much faster than diffusion because there are no other gas molecules to get  in the way.
16   gas laws    
17  

Gay-Lussac’s law

 

 

18  

Gay-Lussac’s law of combining volumes of gases

 

 

19  

Graham’s law of effusion

 

 

20   ideal gas    
21  

Ideal gas constant

(R) ideal gas constant; universal gas constant.

A constant R equal to PV/(nT) for ideal gases, where the pressure, volume, moles, and temperature of the gas are P, V, n, and T, respectively. The value and units of R depend on the units of P, V, and T. Commonly used values and units of R include: 82.055 cm3 atm K-1 mol-1; 0.082055 L atm mol-1 K-1; 8.31434 J mol-1 K-1; 1.9872 cal K-1 mol-1; 8314.34 L Pa mol-1 K-1; 8.31434 Pa m3 mol-1 K-1.

22  

Ideal gas law

ideal gases; perfect gas;

A gas whose pressure P, volume V, and temperature T are related by PV = nRT, where n is the number of moles of gas and R is the ideal gas law constant. Ideal gases have molecules with negligible size, and the average molar kinetic energy of an ideal gas depends only on its temperature. Most gases behave ideally at sufficiently low pressures.

234  

Inversely proportional

 

 

24  

K to C conversion

 

 

25  

Millimeters of mercury

 

 

26   monoatomic gas    
27  

Newton

 

 

28  

Partial pressure

 

The pressure of one gas in a mixture.  For example, if you had a 50:50 mix of helium and hydrogen gases and the total pressure was 2 atm, the partial pressure of hydrogen would be 1 atm.

29  

Pascal

(Pa)

The SI unit of pressure, equal to a force of one newton per square meter. 101325 pascals = 1 atmosphere; 105 pascals = 1 bar.

30  

Pressure

(P)

Force per unit area. The SI unit of pressure is the pascal, defined as one newton per square meter. Other common pressure units are the atmosphere, the bar, and the Torr.

31   Standard Atmospheric Pressure    
32  

Standard molar volume of gas

 

The volume of 1 mole of an ideal gas at STP, equal to 22.414 liters.

33  

STP

standard temperature and pressure.

Used to describe a substance at standard pressure and a temperature of 0°C (273.15 K).

34   Temperature    
35  

Universal gas constant

Ideal gas constant

 

36   volitile   A substance with a high vapor pressure.