Answer: A) Forces of attraction and repulsion exist between gas particles at close range.
Explanation:
The <u>Ideal Gas equation</u> is:
Where:
is the pressure of the gas
is the volume of the gas
the number of moles of gas
is the gas constant
is the absolute temperature of the gas in Kelvin
According to this law, molecules in gaseous state do not exert any force among them (attraction or repulsion) and the volume of these molecules is small, therefore negligible in comparison with the volume of the container that contains them. In this sense, real gases can behave approximately to an ideal gas, under conditions of high temperature and low pressures.
However, at low temperatures or high pressures, real gases deviate significantly from ideal gas behavior. This is because at low temperatures molecules begin to move slower, allowing the repulsive and attractive forces among them to take effect. In fact, <u>the attraction forces are responsible for the condensation of the gas</u>. In addition, at high pressures the volume of molecules cannot be approximated to zero, hence the volume of these molecules is not negligible anymore.
Answer:
A) 20 N, B) 20 N, & C) 8 N
Explanation:
For the object to be in equilibrium, the upward forces must be equal to the downward forces and the forward forces must be equal to the backward forces.
1. Determination of A and B.
Forward forces = Backward forces
A + 10 + B = 25 + 25
A + 10 + B = 50
Collect like terms
A + B = 50 – 10
A + B = 40
Assume A and B to be equal. Thus, A is 20 N and B is 20 N.
2. Determination of C
Upward forces = Downward forces
C + 112 = 20 + 100
C + 112 = 120
Collect like terms
C = 120 – 112
C = 8 N
Thus, for the object to be in equilibrium, A must be 20 N, B must be 20 N and C must be 8N.
Since it is restricted at both ends, λ/2 = length of string
λ/2 = 1.5m
λ = 1.5*2 = 3m
Answer:
D. The ice-to-liquid phase change of water requires less energy than the liquid-to-vapor phase.
Explanation:
In the phase change from liquid to gas, the bonds between atoms are completely broken. The phase change from liquid to gas requires more energy because the bonds must be completely broken for it to take place, rather than just loosened as in the phase change of solid to liquid.
Phase changes can have a strong stabilizing effect on temperatures that are not near the melting and boiling points, since evaporation and condensation occur even at temperatures below the boiling point.
More energy is required to evaporate water below the boiling point than at the boiling point, because the kinetic energy of water molecules at temperatures below 100°C is less than that at 100°C, so less energy is available from random thermal motions.
