Gas Laws & the Mole.pdf

Gas Laws & the Mole Chemistry  Particles in Solids, Liquids and Gases Solids, particles are… Closely packed together, tightly bound to each other. Solids have a fixed volume and shape. Constantly vibrate. Vibrate on heating, more and more, until at melting poi...

Gas Laws & the Mole Chemistry  Particles in Solids, Liquids and Gases Solids, particles are… Closely packed together, tightly bound to each other. Solids have a fixed volume and shape. Constantly vibrate. Vibrate on heating, more and more, until at melting point they break free from each other and liquid is formed. Liquids, particles are… Still close together. Can slip by one another easily. On heating, particles move with greater speed and, at boiling point they escape completely from the other particles and a gas is formed. Gas, particles are… Relatively free of each other, therefore gas has no fixed shape or volume at a particular temperature. Move rapidly and randomly, colliding with each other and with walls of container. Gases are easily compressed, unlike solids and liquids.  Diffusion Diffusion is the spontaneous spreading out of a substance, and is due to the natural movement of its particles. When ammonia and hydrogen chloride come into contact with each other, a white cloud of ammonium chloride is formed. Diffusion of liquids is much slower than diffusion in gases.  Gas Laws - Boyle’s Law & Charles’ Law Boyle’s Law - At a constant temperature, the volume of a given mass of any gas is inversely proportional to the pressure of the gas. Boyle found, that a constant temperature P x V = constant. Charles’ Law - At a constant pressure, the volume of a given mass of any gas is directly proportional to the Kelvin temperature. Charles discovered that equal volumes of different gases at constant pressure all expanded by the same amount. V / T = constant  Gas Laws - Gay-Lussac’s Law & Avogadro’s Law Gay-Lussac’s Law of combining volumes - When gases react, the volumes consumed in the reaction bear a simple whole number ratio to each other, and to the volumes of any gaseous product of the reaction, all volumes being measured under the same conditions of temperature and pressure. Avogadro’s Law - Equal volumes of gases, under the same conditions of temperature and pressure, contain equal numbers of molecules.  The Mole A mole of any substance is defined as the amount of substance that contains as many particles (atoms or molecules or ions) as there are atoms of C-12 in 12g of C-12. The number of atoms of the C-12 isotope in 12g of C-12 is found to be approximately 6 x 10²³. This is the number of particles per mole for all substances, and it is called Avogadro’s Constant (L). L = 6 x 10²³mol-¹  Molar Volume of Gases Standard temperature and pressure (s.t.p.) Pressure = 101,325 Nm−² = 101,325 Pa Temperature = 273 K Molar volume at s.t.p. Room Temperature & Pressure 1 mole = 22.4 L 1 mole = 24 L 1 mole = 22,400 cm³ 1 mole = 24,000 cm³ 1 mole =2.24 x 10−² m³  Relative Molecular Mass The relative molecular mass of a substance is the average mass of a molecule of the substance relative to one-twelfth of the mass of an atom of C-12. Relative molecular mass can be calculated by adding the relative atomic masses of all the atoms in the molecule. Examples of RMM Copper = 63.5 Zinc = 65 Iron = 56 Etc. *Mass spectrometer is used to measure R.M.M. of a substance.  Relative Molecular Mass & Molar Mass The molar mass of a substance is the mass in grams of one mole of the substance. Molar mass is measured in g mol-¹. Substance R.M.M. Molar Mass x R.M.M. H₂SO₄ 98 98g mol-¹ C12H22O11 342 342g mol-¹ Moles Mass ÷ R.M.M. Combined Diagram  The Combined Gas Law Boyle’s Law, Charles’ Law and Avogadro’s Law can be combined to give the combined gas law: Where P1, V1 and T1 are the initial pressure, volume and Kelvin temperature respectively, and P2, V2 and T2 are the final pressure, volume and Kelvin temperature respectively.  The Kinetic Theory of Gases Clerk Maxwell and Ludwig Boltzmann assumed that: Gases are made up of particles whose diameters are negligible compared to the distance between them. There are no attractive or repulsive forces between these particles. The particles are in constant rapid random motion, colliding with each other and with the walls of the container. The average kinetic energy of the particles is proportional to the Kelvin temperature. All collisions are perfectly elastic (for example, if a particle travelling at 450 m/sec collides with a wall of its container, it rebounds with the same speed. An ideal gas is a gas that perfectly obeys all of the gas laws under all conditions of temperature and at high pressures. A real gas obeys kinetic theory under high temperature and low pressure. Real gases behave least like an ideal gas under low temperature and high pressure. CONDITIONS; Temperature and Pressure.  The Kinetic Theory of Gases Reasons why a real gas deviates from ideal behaviour ? 1. Diameters are no longer negligible compared to the distance between molecules. 2. Attractive and repulsive forces now exist between molecules as they are close together. 3. Collisions are no longer perfectly elastic. The Equation of State for an Ideal Gas Pressure PV = nRT KPa → Pa x1000 P = pressure (in Pa) Volume cm³ → m³ x10−⁶ V = volume (in m³) L → m³ x10−³ n = no. of moles Temperature R = universal gas constant (8.31) Celsius → K +273 T = temperature (in K)  Expt: Estimation of the R.M.M. of a volatile liquid Text Book Pg 154 To find the relative molecular mass of a volatile liquid; 1. Find the number of moles using PV = nRT, n = PV ➗ RT 2. Calculate the relative molecular mass, Mr = M➗n, M = mass in g (normally given in the question, n = no. of moles (from part 1.) Volatile liquid; has a low boiling point and is easily vaporised.