Module 3 Periodic Table and Energy

• Periodic table


  • Mendeleev arranged in order of atomic mass & lined up those with similar properties. Gaps left where he thought elements would be found.
  • Now: arranged in increasing atomic number, in vertical columns (groups) with same number of outer electrons + similar properties & horizontal rows (periods) giving number of highest energy electron shell
• Periodicity


  • Repeating, periodic pattern across a period
• Electron configuration across a period

  Period 2: 2s then 2p, period 3: 3s then 3p, period 4: 3d filled but highest energy is 4: 4s then 4p
• Periodic table blocks


  Highest energy sub-shell s,p,d,f
• Groups in the periodic table

  • Group 1= alkali metals, 2=alkaline earth metals, 3-12=transition elements, 15=pnictogens, 16=chalcogens, 17=halogens, 18=noble gases
• First ionisation energy


  Energy required to remove one electron from one mole of gaseous atoms, forming one mole gaseous 1+ ions
• Factors affecting ionisation energies

  • Atomic radius (greater distance nucleus to outer electrons=less attraction, big effect), nuclear charge (more protons, more attraction), electron shielding (shielding effect- inner shell electrons repel outer shell electrons→ reduced attraction nucleus to outer electrons)
• Successive ionisation energies are greater: after 1st electron lost, remaining electrons pulled closer to nucleus, nuclear attraction increases so more energy needed
• Large increase successive ionisation energies→ electron has been removed from shell closer to nucleus with less shielding
• Successive ionisation energies allow predictions about: no. electrons outer shell, group in periodic table → element can be identified
• Trends in ionisation energy

  • General increase across period, sharp decrease between end of a period to the start of the next
  • Down a group: atomic radius increases, more inner shells so shielding increases, nuclear attraction on outer electrons decreases, 1st IE decreases
  • Across a period: nuclear charge increases, same shell so similar shielding, nuclear attraction increases, atomic radius decreases, 1st IE increases. Exceptions period 2+3: group 2-3 fall (2p subshell higher energy than 2s, so 2p electron easier to remove) & group 5-6 fall (highest energy in 2p, but paired in 6- electrons repel making it easier for them to be removed)
• Metallic bonding


  Strong electrostatic attraction between cations (+ve) and delocalised electrons. Cations fixed in position (maintains shape) & delocalised electrons mobile
• Metals w/ 2+ cations have 2x electrons
• Properties of metals


  • Electrical conductivity (electrons can move when voltage is applied), high mpt/bpt (high temp needed to overcome strong electrostatic attraction between cations/electrons) & insoluble (any interactions lead to reaction not dissolving)
• Giant covalent lattice


  Many billions of atoms held together by network strong covalent bonds (boron, carbon, silicon)
• Carbon (diamond) + silicon


  • Use 4 outer electrons forming covalent bonds with other atoms→ tetrahedral, 109.5°, can be shown w/ dot+cross diagram
• Properties of giant covalent structures

  • High mpt/bpt (covalent bonds strong so high energy to break), insoluble in almost all solvents (bonds too strong to be broken by interactions w/ solvents), electrical conductivity (diamond/silicone no- no electrons not involved in bonding + graphene/graphite- yes)
• Graphene


  Single layer graphite, hexagonally arranged (planar 120°) carbons, conducts electricity & thinnest + strongest material in existence
• Graphite


  Parallel layers hexagonally arranged carbon atoms (planar 120°). Layers bonded by weak london forces, spare electron delocalised between layers→ conducts electricity
• Periodic trend melting points period 2+3: increases group 1-14, sharp decrease 14-15, comparatively low 15-18
• Group 2 reactions


  • Most common reactions of group 2= redox. They act as reducing agents
• Reaction with oxygen


  2M (s) + O2 (g) →2MO (s) (Mg burns w/ white light, MgO=white) (M=generic group 2 metal)
• Reaction with water


  M (s) + 2H2O (l) →M(OH)2 + H2 (g). Reaction more vigorous as reactivity increases down group
• Reaction with dilute acids


  Reactivity increases down group, metal + acid→salt + hydrogen e.g. M (s) + 2HCl (aq) →MCl2 (aq) + H2 (g)
• Reasons for increasing reactivity down group 2


  • Lose 2 electrons, requiring energy for 1st+2nd ionisation energies, these decrease down group b/c attraction decreases b/c atomic radius + shielding increase
• Reaction of oxides with water


  MO(s) + H2O (l)→ M2+ (aq) + 2OH- (aq), only slightly soluble so once solution saturated, any further ions: M2+ (aq) + 2OH- (aq) -> M(OH)2 (s)
• Solubility of group 2 hydroxides


  • Increases down group→ more OH- ions→ more alkaline (higher pH)
• Compounds in agriculture


  Ca(OH)2 (s) + 2H+ (aq) → Ca2+ (aq) + 2H2O (l)
• Compounds in medicine


  1. Mg(OH)2 (s) + 2HCl (aq) → MgCl2 (aq) + 2H2O (l)

  2. CaCO3 (s) + 2HCl (aq) -> CaCl2 (aq) + H2O (l) + CO2 (g)
• Halogens


  On earth occur dissolved in seawater OR combined with sodium/potassium as solid deposits
• Appearance and state of halogens at room temperature and pressure

  • F2 : pale yellow gas (reacts w/ almost any substance)
  • Cl2 : pale green gas
  • Br2 : red-brown liquid
  • I2 : shiny grey-black solid
  • At2 never been seen (radioactive + decays rapidly)
• Trend in boiling points of halogens

  • More electrons, stronger London forces, more energy to break intermolecular forces, bpt increases
• Most common halogen reactions

  • Redox, oxidising agents
• Halogen-halide displacement reactions


  1. Cl2 reacts with Br- (Cl2 (aq) + 2Br- (aq) →2Cl- (aq) + Br2 (aq) orange)

  2. Cl2 reacts with I- (Cl2 (aq) + 2I- (aq) → 2Cl- (aq) + I2 (aq) violet)

  3. Br2 reacts with I- only (Br2 (aq) + 2I- (aq) → 2Br- (aq)+ I2 (aq) violet)

  4. I2 doesn't react at all
• Appearance of halogen solutions

  • Cl=pale green, Br=orange, I=brown (in water)
  • Cl=pale green, Br=orange, I=violet (in cyclohexane /nonpolar solvent)
• Trend in reactivity of halogens

  • Atomic radius increases, more inner shells so shielding increases, less nuclear attraction to capture another electron, reactivity decreases
  • Become weaker oxidising agents down group
• Disproportionation of chlorine


  1. Cl2 (aq) + H2O (l) → HClO (aq) + HCl (aq) (bacteria killed by chloric (I) acid/ions, chloric (I) acid acts as weak bleach- indicator paper will turn red then white)

  2. Cl2 (aq) + 2NaOH (aq)→ NaCl (aq) + H2O (l) + NaClO (household bleach)
• Benefits and risks of chlorine in water

  • Benefits: kills bacteria preventing disease e.g.cholera
  • Risks: can react with organic hydrocarbons (formed by decaying vegetation) to make chlorinated hydrocarbons- carcinogens, could use other methods purification
• Carbonate test


  1. Add dilute HNO3 , test bubbles produced by bubbling through limewater→ white ppt

  2. Eq: Na2CO3 + 2HNO3 → 2NaNO3 + CO2 + H2O

  3. CO2 + Ca(OH)2 → CaCO3 (s) + H2O
• Sulfate test


  1. Add barium chloride(not if testing for halides after)/nitrate→ white ppt

  2. Eq: Ba2+ + SO4 2- → BaSO4 (s)