CHEMISTRY UNIT 1

• The kinetic theory of gases explains the observed behaviour of gases based on the movement of particles of which the gas is composed.
• Matter on Earth is composed of substances in solid, liquid or gas phase, depending on the conditions in which it is found.
• A substance can move from one phase to another without any change in its nature, making these physical changes.
• Plasma is the fourth state of matter, in which a gas is partially ionized and light is emitted (e.g. plasma TV).
• The particle model can explain some of the properties of the phases of matter: all matter is composed of particles (ions, atoms or molecules) with more or less space between them depending on the phase.
• Particles attract or repel each other with varying force depending on the distance between them.
• Particles are always moving.
• Vibration is the only type of motion and is extremely weak in solids.
• Strong forces exist between particles, preventing them from moving in relation to one another, resulting in a definite shape and volume in solids.
• Particles are ordered in solids.
• Motion in liquids is vibrational, rotational and weakly translational.
• Greater motion than in solids because the strong bonds between particles have been broken in liquids.
• In gases, motion is vibrational, rotational and translational.
• In translational motion, particles follow random linear trajectories until they collide with other particles or objects in gases.
• Movement is fast in gases.
• Gases flow in all directions despite gravitation force, to fill their containers.
• Particles are far apart, allowing gases to be compressed.
• Particles are very disordered in gases.
• The kinetic energy of a substance corresponds to the amount of particle movement.
• In gases, increasing temperature increases the rate of vibration, rotation and translation.
• Kinetic energy: energy of a particle in motion.
• The greater that mass of a particle and the faster it moves, the greater its quantity of kinetic energy.
• In a sample of gas, not all particles have the same kinetic energy and therefore velocity.
• It is best to consider mean velocity and kinetic energy for a gas at a given temperature.
• The mean velocity increases with increasing temperature.
• As mean velocity increases, so does mean kinetic energy.
• Avogadro's law states that under the same temperature and pressure conditions, the volume of a gas is directly proportional the its quantity expressed in moles.
• The more frequent collisions cause the pressure inside the container to increase.
• The relationship between pressure and temperature of a gas is another form of Charles' law.
• Charles' law states that at constant pressure, the volume occupied by a given quantity of gas is directly proportional to the absolute temperature of the gas.
• The Kelvin scale, also called the absolute temperature scale, has absolute zero (0 K) as a starting point and thus does not have any negative values.
• Kinetic theory is used to explain the observed behaviour of an ideal gas in a rigid container.
• Kinetic theory states that increasing the temperature of a gas results in more rapid movements.
• Avogadro furthered Gay-Lussac's work and hypothesized that under the same temperature and pressure conditions equal volumes of gas contain the same number of particles (moles).
• Gay-Lussac demonstrated that a chemical reaction between gases always occurs in a simple whole number ratio of volumes provided pressure and temperature are held constant.
• In a container whose volume is not constant (ex balloon or syringe), this increased internal pressure will result in an increase in volume until the internal and external pressures are equal.
• For the same gas sample at constant volume: 𝑃 1 �� 1 = �� 2 �� 2.
• As the temperature of a gas increases, its particles move more quickly and collide with the walls of the container more frequently.
• Using the Kelvin scale makes it possible to observe that the relationship between volume and absolute temperature is directly proportional.
• Gay-Lussac's law states that at constant volume, the pressure of a given quantity of gas is directly proportional to the absolute temperature of the gas.