Charles Law Lab Report

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  • Topic: Pressure, Vapor pressure, Vapor
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  • Published : August 23, 2010
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125 mL Erlenmeyer flask, one-hole rubber stopper, glass and rubber tubing, pneumatic trough, thermometer, screw clamp.

The quantitative relationship between the volume and the absolute temperature of a gas is summartzed in Charles'law. This law states: at constant pressure, the volume of a particular sample of gas is directly proportional to the absolute temperature. Charles' law may be expressed mathematically: V ". T (constant pressure) V = kT o, : T = k (constant pressure) (1) (2)

where V is volume, T is Kelvin temperature, and k is a proportionality constant. dependent on the number of moles and the pressure of the gas. If the volume of the same sample of gas is measured at two temperatures, V1/T1 = k and V2/T2- k, and we may say that V, V, or V" = (V,f[]) z \ T1 T2 "[Tr/ {.o.rrt.nt pressure)


where V1 and T, represent one set of conditions and V2 and T2 a different set of conditions, with pressure the same at both conditions. Experimental Verification of Charles' Law

This experiment measures the volume of an air sample at two temperatures, a high temperature, Ts, and a low temperature, T1. The volume of the air sample at the high temperature, (Vn),decreases when the sample is cooled to the low temperature and becomesV1. All of these measurements are made directly. The experimental data is then used to verify Charles'law by two methods: 1. The experimental volume (V""o) measured at the low temperature is compared to the V1 predicted by Charles' law where

Yy(t oretic (vH,[ he at)= + )


2. The V/T ratios for the air sample measured at both the high and the low temperatures are compared. Charles'law predicts that these ratios will be equal.


Pressure Considerations The relationship between temperature and volume defined by Charles' law is valid only if the pressure is the same when the volume is measured at each temperature. That is not the case in this experiment. 1. The volume, Vs, of air at the higher temperature, Ts, is measured at atmospheric pressure' P"t* in a dry Erlenmeyer flask. The air is assumed to be dry and the" is obtained from a barometer. 2. The experimental air volume, (V"*p) at the lower temperature, Tp, is measured. over water. This volume is saturated with water vapor that contributes to the total pressure in the flask. Therefore, the experimental volume must be corrected to the volume of dry anrat atmospheric pressure. This is done using Boyle's law as follows: a. The partial pressure of the dry air, Poo, is calculated by subtracting the vapor pressure of water from atmospheric pressure: P.r--PffrO=POA b. The volume that this dry air would occupy at Pur,''is then calculated using the Boyle's law equation:

= (%,.oXp*) (voo)(%,_) (%,.oXp*) . =Sffi (voo)

Wear protective glasses.

NOTE: It is essential that the Erlenmeyer flask and rubber stopper assemblvbe as drv as possiblein order to obtain reproducibleresults. Dry a L25 mL Erlenmeyer flask by gently heating the entire outer surface with a burner flame. Care must be used in heating to avoid breaking the flask. If the flask is wet, first wipe the inner and outer surfaces with a towel to remove nearly all the water. Then, holding the flask with a test tube holder, gently heat the entire flask. Avoid placing the flask directly in the flame. Allow to cool. While the flask is cooling select a l-hole rubber stopper to fit the flask and insert a b cm piece of glass tubing into the stopper so that the end of the tubing is flush with the bottom of


the stopper.Attach a 3 cm piece of rubbertubingto the glass tubing (see Figure 19.1-). Insert (wax pencil) the distance that it is inserted. Clamp the the stopper into the flask and mark flask so that it is submerged as far as possible in water contained in a 400 mL beaker (without the flask touching the bottom of the beaker) (see Figure I9.2). Heat the water to...
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