Module 3 : Periodic table and energy

• Mendeleev's periodic table


  arranged in order of atomic mass. he lined up those with similar properties into groups. if they didn't fit he left gaps, confident that the missing elements would be found; he even predicted their properties
• modern periodic table


  arranged by increasing atomic number

  group number = number of electrons in outer shell

  period number = number of electron shells
• periodicity


  the repeating pattern of chemical and physical properties of the elements across each period
• sub-shell blocks


  s-block, p-block, d-block, f-block

  tells us electron configuration

  e.g. s-block elements have an outer shell config of s1 or s2

  Li = 1S2 2S1 Mg = 1s2 2s2 2p6 3s2

  
• what is ionisation?


  involves the loss of an electron to form a positive ion. energy is needed (endothermic process) - the amount of energy needed depends on the size of the nuclear charge and the energy of the electron being removed
• first ionisation enthalpy


  the energy required to remove 1 electron from each atom in a mole of gaseous atoms to form a mole of 1+ ions

  
• second ionisation enthalpy


  the energy required to remove 1 electron from each ion in a mole of gaseous 1+ ions to form a mole of 2+ ions

  
• ionisation energy of oxygen


  1st: 0(g) → O+(g) + e-

  2nd: O+(g) → 02+ + e-

  6th: O5+ → O6+ + e-

  
• factors affecting ionisation energy

  nuclear charge, atomic radius, electron shielding
• how does nuclear charge affect IE?

  the more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons (so the IE is higher)
• how does atomic radius affect IE?

  the greater the distance between the nucleus and the outer electrons, the weaker the nuclear attraction. (less energy required to remove electron further from the nucleus = lower IE)
• how does electron shielding affect IE?

  as the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge. (lower IE)
• trends in ionisation energy

  decreases down a group - more electron shells, atomic radius is larger, outer electrons further from nucleus, inner shells shield outer electrons from attraction of the nucleus

  
  

  increases across a period - nuclear charge increases, increases attraction of electrons, pulls them closer to the nucleus, atomic radius slightly decreases, more energy to overcome attraction
• trend in ionisation energy across a period

  general increasein IE due toincreasing nuclear charge. occasional drops occur e.g. due to filling a p sub-shell (only in period 2/3 elements)

  
• why is there a drop in ionisation energy between group 2 and 3?

  theouter electroningroup 3 elementsis in ap-orbital rather than an s-orbital.ap-orbitalisslightly higher in energythan an s-orbital in the same shell, so the electron is, on average, to befound further from the nucleus.the p-orbital also hasadditional shieldingprovided by the s electrons.these factorsoverride the effect of the increased nuclear charge, resulting in theIE dropping slightly.

  
• why is there a drop in ionisation energy between group 5 and 6?

  ingroup 5 elements, the electron is being removed fromsingly-occupied orbitals.ingroup 6 elements, the electron is being removed from anorbital containing 2 electrons.therepulsion between 2 electronsin an orbital means that electrons areeasier to removefrom shared orbitals

  
• what are simple molecular substances?

  they contain only a few atoms. e.g. 02, Cl2, P4, S8
• melting/boiling points of simple molecular substances


  thecovalent bonds between the atomsof SMS arevery strong, but themp & bpof SMSdepend upon the strength of the induced dipole-dipole forcesbetween the molecules. these intermolecular forces areweak and easily overcome, so these substances havelow mp & bp.more atomsin a molecule =strongerinduced D-D forces =highermp & bp. e.g.S8 has more atomsand ahigher mp/bpthanCl2 and P4.noble gaseshavevery lowmp & bp as they exist asindividual atoms(monatomic) -very weak ID-D forces.
• define metallic bonding


  theelectrostatic attractionbetweenpositive metal ionsanddelocalised electrons
• What is a giant metallic lattice?

  repeating layers of metal atoms joined together by metallic bonding
• how do metallic structures form?

  in a solid metal structure,each atom has donated its negative outer-shell electronsto ashared pool of electrons, which aredelocalisedthroughout the whole structure. thepositive ions (cations)left behind consist of thenucleus & the inner electron shellsof the metal atoms.thecations are fixed in position,maintaining the structure and shapeof the metal. thedelocalised electronsaremobileand are able tomove throughout the structure.only the electrons move.

  
• properties of metals


  good conductors of heat and electricity, malleable, ductile, high tensile strength, lustrous, hard & durable, high mp/bp, insoluble
• electrical conductivity of metals

  metals are good conductors of electricity because thedelocalised electrons can move and carry charge. metalscan conduct in solid and molten states. theions are fixed- do not move.only the electrons move.
• melting and boiling points of metals

  most metals havehigh mp/bps-strong metallic bondsrequirea lot of energytoovercometheforces of attractionbetweenions and delocalised electrons.themp depends upon the strength of the metallic bondsholding together the atoms in the giant metallic structure.
• solubility of metals


  metalsdon't dissolve- themetal bonds are too strongto allowwater molecules to make new interactionswith thepositive ions. if there is aninteraction, between polar solvents and the charges in a metallic lattice, it would be achemical reaction(rather than dissolving)
• how are metals malleable and ductile?

  there areno bonds holding specific ions togetherso the metal ions canslide past each other, so metals can be hammered into different shapes
• properties of giant covalent structures

  very high melting and boiling pointsdon't conduct electricity- not even when molten (only graphite can)insolublethe properties are dominated by thestrong covalent bonds, which make forvery stable structuresthat arevery difficult to break down
• melting and boiling points of giant covalent structures

  giant covalent lattices havehigh mp & bp, due tostrong covalent bonds.high tempsare required to provide thelarge quantity of energyneeded tobreakthe strong covalent bonds.
• solubility of giant covalent structures

  giant covalent lattices areinsoluble in almost all solvents. thecovalent bondsholding together that atoms in the lattice arefar too strongto bebroken by interaction with solvents.
• electrical conductivity of giant covalent structures

  giant covalent lattices arenon-conductors of electricity. onlyexceptions are graphene and graphite.incarbon (diamond) & silicon,all 4 outer shell electronsare involved incovalent bonding, sononeare available forconducting electricity.carbon also formsgraphene & graphitewhich have1 electron available for conducting, so theycan conduct electricity.
• structure & properties of graphene

  asingle layer of graphite, composed ofhexagonally arranged carbon atomslinked bystrong covalent bonds. it has thesame electrical conductivity as copper, and is thethinnest & strongest material ever made.

  
• structure & properties of graphite


  it's composed ofparallel layersofhexagonally arranged carbon atoms, like astack of graphene layers. the layers areheld together(bonded) byweak London forces. the bonding in the hexagonal layers only uses3 of carbons 4 outer-shell electrons.thespare electron is delocalised between the layers, soelectricity can be conductedlike in metals

  
• periodic trend (across 2/3) in melting point

  mp increasesfromgroup 1 to group 4.sharp decrease in mp between group 4 & 5- marks a change fromgiant to simple molecular structures. onmelting,giantstructures havestrong forces to overcomeso havehigh mp.simplemolecular substances haveweak forcesto overcome, so have amuch lowermp.themp are comparatively low from group 5 to group 0.

  
• group 2 elements


  alkaline earth metals - the elements are reactive metals & don't occur in their elemental form naturally; found in stable compounds, e.g. CaCO3. 2 valence electrons
• redox reactions of group 2 elements


  the elements have2 electrons in their outer shell, found in theouter s sub-shell (s2)in redox reactions,each metal atom is oxidisedas theylose 2 electronsto form a2+ ion(electron config. of noble gas is achieved)another species will gain these 2 electronsand bereduced, so thegroup 2 element is a reducing agent
• reactivity of group 2 elements

  reactivity increases down the groupas you godownthe group, theionisation energies decrease. this is due to theincreasing atomic radius & shielding effect, ∴attractionbetween nucleus and outer electronsdecreases.when these elements react theylose electrons, formingpositive ions(cations). theeasier it is to lose electrons(lower the 1st & 2nd IE), themore reactivethe element.so group 2 elements becomemore reactive and stronger reducing agents down the group
• group 2 reactions with water

  group 2 elements react with water to form analkaline hydroxide, general formula M(OH)₂, andhydrogen gas. reactionbecomes more vigorous with metals further down the group.Ca (s) + 2H₂O (l) → Ca(OH)₂ (aq) + H₂ (g)0 +2 (ox)+1 0 (red)
• group 2 reactions with oxygen

  group 2 elements react with oxygen to form asolid, white metal oxide. general formula MO, made up of M2+ ions and O2- ions.2Ca (s) + O₂ (g) → 2CaO (s)0 +2 (ox)0 -2 (red)
• group 2 reactions with dilute acid


  metal + acid → salt + hydrogen gaswhen group 2 metals react withdilute hydrochloric acid, ametal chlorideand hydrogen is formed:Ca (s) + 2HCl (aq) → CaCl₂ (aq) + H₂ (g)0 +2 (ox)+1 0 (red)
• group 2 oxides


  the oxides & hydroxides of group 2 metals are bases. most of them are soluble in water, so are also alkalis.