2nd IB Chem HL

• Boyle's Law


  P1V1=P2V2 (inversely proportional)
• Charle's Law


  V1/T1=V2/T2
• Gay-Lussac's Law


  P1/T1=P2/T2
• Combined Gas Law


  P1V1/T1=P2V2/T2
• Ideal Gas Law


  PV=nRT
• conditions of PV=nRT


  Pascals, m3, mol, 8.314 JK-1mol-1

  -real gases deviate at high pressure (closer particles, volume significant)

  -deviate at low temp (slow movement, IMF)
• ideal gas


  low molar mass and weak to no IMF, negligible dimension, no change in kinetic energy (perfectly elastic), STP conditions
• effective nuclear charge formula

  # of protons - # of protons of previous noble gas
• effective nuclear charge


  extent to which the nucleus is attracting the valence electrons (more difference, stronger)
• effective nuclear charge increase across periods

  linear
• atomic radius


  half the difference between the nuclei of an element (depende on what it is bonded with)
• atomic radius tendencies


  increases down a group because more shells, decreases across periods because more protons means more attraction, noble gases have none
• more effective nuclear charge means

  smaller atomic radius
• atomic radius graph

  
• ionic radius


  half the distance between the nucleus of the cation and the anion (Average)
• parent ion has _______ ionic radius than cation

  higher (because loses electron and less shells)
• ionic radius trend


  decreases across a period and increases down a group
• ionization energy


  minimum amount of energy required to remove one mol of electrons from a neutral gaseous atom in its ground state
• ionization energy trend


  decreases from top to bottom in a group; increases from left to right in a period
• electron affinity


  energy released when 1 mol of electrons is attached to 1 mol of neutral gaseous atoms or molecules
• electron affinity trend


  -greater nucleus, more effective nuclear charge, more attraction, more energy released

  -down group energy decreases because there are more shells and it is easier to it an electron in a bigger atom
• electronegativity


  tendency of an atom to gain electron
• electronegativity trend


  increases across a period, decreases down a group
• melting point trend


  1. decreases down group 1

  2. increase down 17

  3. increase across a period and reach a maximum at group 14
• alkali metal properties (group 1)

  -very reactive

  -form ionic compounds with non-metals

  -react with H2O to produce hydroxide and gas

  -intensity increases down group (valence further away)
• Halogen properties (group 7)

  -some gas/liquid/solid

  -colored

  -very reactive

  -diatomic

  -Fluorine, chlorine, bromine, iodine (order reactivity)
• Oxides from Period 3 elements


  Giant ionic: Na2O, MgO, Al2O3

  Giant Covalent: SiO2

  Molecular Covalent: P, S, Cl
• oxides from period 3 conductivity/pH

  -Na, Mg, Al (basic and conductors bc mobile ions)

  -Al amphoteric

  -Si, P, S, Cl (acidic and non conductive)
• Na oxide reaction


  Na2O + water = 2NaOH
• Mg oxide reaction


  MgO + water = Mg(OH)2
• Aluminum oxide


  Al2O3
• Phosphoric acid reaction


  P4O10+6water=4H3PO4
• phosphorous acid reaction


  P4O8 + water = H3PO3
• sulfuric acid reaction


  SO3 + H2O = H2SO4
• sulfurous acid reaction


  SO2 + H2O = H2SO4
• ionic bond electronegativity difference

  more than 1.8
• coordination number


  how many anions around cation in a lattice
• lattice energy


  how much energy needed to break ionic bond
• Coulomb's Law


  electrostatic force inversely proportional to distance between charged particles, but directly proportional to the charge of the particles
• ionic compound physical properties

  brittle, soluble, non-volatile, conductive