My books says that at high presure and low temperatures, real gases do not obey Boyle's and Charles's Law, what is the reasoning behind it?[not given in the book]
Whenever your book says that they do not obey Boyle's and Charle's law, it is not stating that 'real gases' act opposite of how ideal gases do. What your book is trying to state is that although real gases will act relatively similar to ideal gases, there are some things that come in to play as to why they will act differently.
When your book says 'Ideal Gases', it basically wants you to get a basic idea on gases before moving on to more complicated things such as intermolecular forces and the van der waals equation. An ideal gas is simply a gas that will occupy 22.4L at STP, however whenever the temperature and pressure are increased significantly, there are other things that have to be accounted for because the ideal gas equation will not suffice in correct calculations. Intermolecular forces cause the deviation from the ideal gas equation, if it's not required in your coursework I suggest you study a little about it anyways.
Okay, what does pV/RT represent or mean? What's the significance of plotting it against pressure?
What they are trying to represent with this (I believe) is that they are showing the significance of pressure and volume versus temperature. As temperature increases, gases volume and pressure increase - however whenever extreme temperatures are reached, the gas will act differently than according to ideal gas law. Think about the laws that were learned before the ideal gas law equation.
What does it mean when the books says "the volume of the gas molecules is negligible in comparison with the total volume of the gas", does it mean that it takes one molecule of the gas and compares with the total volume of the whole gas?
What your book is trying to say (quite poorly I believe) is that the volume of the gas molecules in comparison with the total volume of the gas is going to differ from what you might calculate with the ideal gas law equation. With ideal gases, it is assumed that whenever particles are in their 'gas-flight' that they simply bounce off each other and then assume flight in a different direction (kinetic-molecular theory.) However, with real gases there are intermolecular forces such as dipole-dipole, ion-dipole, london dispersion forces, and hydrogen bonding (look 'em up!) Due to these intermolecular forces, the way that the molecules act is different from ideal gases. As stated above, the ideal gas will simply bounce off the molecules and then continue its path, but with real gases the molecules will be attracted to each other for a brief moment, lowering the amount of volume it occupies because molecules are moving around less freely. Think of it as a large container filled with perpetual motion bouncy balls, and another container with the same conditions only there is a layer of tape on the outside. The container without the tape is going to be much more random, and the molecules will each go their own way, whereas with the taped balls - they will be stuck to other bouncy balls until the 'bonds' are broken apart.
What are postive and negative deviation?
Imagine a basic line of Y=X (a slope of 1, starting from the origin.) Now think of this line as the ideal gas law with temperature on the x-axis, and pressure/volume on the y-axis. A positive deviation will be placed above the ideal-gas line (increase in slope) and a negative deviation will be below the line (decrease in slope.)
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Topic: Some questions on Gas Laws (Read 3514 times)