We started with learning about the Kinetic-Molecular Theory (KMT) which is a model that aids our understanding of what happened to gas particles as environmental conditions change. A measure of the average kinetic energy of atoms and molecules in a system is temperature. The Kelvin scare is proportional to this, so when kinetic energy doubles the Kelvin temperature doubles. The main parts of the KMT are - gases consist of large numbers of molecules that are in random motion, the combined volume of all the molecules of the gas is negligible relative to the total volume in which the gas is contained (even though they have different masses), masses are considered volume less points of mass (they have mass but no volume), and there are no attractive or repulsive forces between gas molecules. Collisions between gas particles are elastic and kinetic energy is conserved. Average kinetic energy does not change with time, as long as temperature stays constant. Molecules with lower mass go faster than those with a higher mass.
Average kinetic energy of molecules is also proportional to the absolute temperature. Effusion is the escape of gas molecules through a tiny hole into an evacuated space. A smaller molar mass means faster effusion. Diffusion is the spread of one substance throughout a space or throughout a second substance. The average distance between collisions is the mean free path. As pressure increases, mean free path decreases.
In the real world, the behavior of gases only conforms to the ideal gas equation at relatively high temperature and low pressure. Real gases are not ideal because they have attractive forces and gas particles have real volume. In the equation, "a" means the interaction part of the molecules and "b" is the volume component of molecules. Attractions between gas molecules reduces pressure and pulls particles away from the walls where pressure is measured and they're not hitting as hard. IMF forces increase as particles get closer which leads to condensation at low temperature or high pressure. Condensation causes temperature to decrease and kinetic energy to decrease and can cause particles to stick together. Condesation also causes pressure to increase and causes higher rate of collisions. Gases that are closest to ideal gases are non polar, have no hydrogen bonding, smaller molecules, high temperature, smallest LDFs, spaced far apart (low pressure), and have less volume. Gasses furthest from ideal gases have low temperatures, low kinetic energy and more time to intreat, high pressures, and the particle count is high.
This week, I'd give my understanding of our new subject about a 9. I think that I understand the topic well enough that I was able to help my classmates with questions that they had. This unit I've really paid attention during class and watched the lectures really closely, which is really different from what I did last unit. I think that doing the Gases II worksheet really helped, as well as doing the Concept tests at the beginning of class. I hope that I do better on this test than thermodynamics unit, which I know I will because I've really been trying to understand everything in the lectures and in class.
