3.1 the periodic table Flashcards
metallic bond
strong electrostatic attraction between cation (+ ions) and delocalized electrons
giant metallic lattice
billions of metal ions held together by metallic bonding (EfA between metal cations and delocalised electrons)
giant covalent lattice
3d structure of atoms held together by strong covalent bonds
describe the giant covalent structure of diamond?
-C has 4 unpaired valence electrons so can form 4 covalent bonds
-no lone pairs
-tetrahedral structure w bond angles of 109.5*
-C atoms are small > short bonds
explain th eproperties of diamond?
-short bonds, very strong covalent bonds, hard, high MP.BP
-all electrons contained in covalent bonds so no free mobile charge carriers, cannot conduct
describe the giant covalent structure of graphite?
-each C form 3 covalent bonds and releases 1 delocalised electron
-forms hexagonal layers of C atoms w delcolaized e-s between
explain graphites properties?
-each C delocalised one electron, delocalised electrons, mobile charge carriers, can conduct
-soft/slippery (lubricant/pencils) London forces between layers, can slide
-strong covalent bonds, High MP/BP
describe the structure of graphene?
-giant covalent structure
-each C atom covalently bonded to 3 other C atoms
-1 atom thick layer of hexagonally bonded C atoms
explain the properties of graphene?
-conducts , 1 delocalised electron/ C atom
-high MP/BP, strong covalent bonds
describe the structure of silicon?
-each Si has 4 valence electrons and no lone pairs
-forms tetrahedral structure, 109.5*
-atoms larger than C so bonds are longer and weaker, lower MP/BP
structure and bonding in silicon dioxide/silica/silicon oxide?
-2 properties?
SiO2
-giant covalent lattice sturcture
-Si atoms form 4 covalent bonds w Oxygen
-tetrahedral
-high MP
-doesnt conduct
Boron Nitride
-boron, grp 3, 3 valence electrons
-nitrogen grp 5, 3 to pair and 1 lone pair
-boron is electrodeficient, even when forms 3 bonds outer shell only has 6
-so accepts lone pair from nitrogen and forms dative covalent bond
2 structures:
-4 covalent bonds around each atom- tetrahedral structure (like diamond- “isoelectric”)
-boron form 2 double bonds w N , 1 dative bond w N
-same structure as graphite (no delocalised, cant conduct)
define isoelectric
same electronic structure as the Carbons in diamond
giant ionic lattice properties explained?
high MP,BP
-large amts energy to overcome strong EfAs between oppositely charged ions in lattice (higher for lattices w greater ionic charges/larger ions)
Dissolve in polar substances
-polar (water) molecules break down the lattice and surround each ion in oslution
Conductivity
-only conduct when (l) or (aq), ionic lattice breaks down, ions free to move and carry a charge
-solid, dont bc ions held in fixed position
giant metallic lattice properties?
MP/BP
depends on strength of metallic bond
-MOST high bc high temp to overcome EfA between cations and delocalised electrons
Solubility
-do NOT dissolve, any interaction between solvent and charges in lattice results in reaction NOT dissolving
conductivity
-conduct in (s) or (l), voltage applied across metal, delocalised electrons can move throughout sy=ubstance and carry charge
giant covalent lattice properties?
MP/BP
-High, strong cov. bonds between shared e- pair and nucluei of bonded atoms req high amts energy to overcome
Solubility
-Insoluble, cov, bonds too strong to be overcome by interaction w any solvent
Conductivity
-diamond/silicon, NO, all atoms held in rigid structure, no charged particles
-graphene/graphite, delocalised electrons, mobile, carry a charge
explain the variation in MP across period 2 and 3 in terms of structure and bonding?
Increase from group 1-3
-ions becoming smaller (decreased atomic radius) and more charged [ions have a higher charge-density moving across the period]
-metallic lattice contains more electrons
-attractions are stronger, MP Increases
Increase from grp3-4
(in period 3) move from metallic structure to giant covalent structure
why does atomic radius decrease across a period?
same shielding, increasing nuclear charge pulling outer shell inwards
define first ionisation enrgy
energy needed to remove 1mol of electrons from 1mol of gaseous atoms, to form 1mol of gaseous 1+ ions
general formula for 1st ionisation energy
X(g) > X^+ (g) + e-
define the term second ionisation energy
enrgy required to remove one electron from each ion in 1mol of gaseous 1+ ions of an element, to form 1mol of gaseous 2+ ions
what factors effect ionisation energy?
-distance (nucleus to outermost electrons[atomic radius])
-nuclear charge (on the nucleus)
-electron shielding
how does distance/atomic radius affect ionisation energy?
increase atomic radius>decreases Ie
-increase atomic radius> decrease force of attraction from nucleus to outermost electrons (nuclear attraction)>decreases Ie
how does nuclear charge affect ionisation enegy?
increase nuclear charge> increase Ie
-greater no. protons in nucleus> increases nucluear charge> increases force of attraction nucleus:outermost electrons (nuclear attraction)>increases Ie
how does electron shielding effect ionisation energy?
Increase shielding>decrease Ie
-electrons in outer shell repelled by electrons in inner shells, shielding effect decreases nuclear attraction> decreases Ie
how do ionisation energies provide evidence for existence of shells?
-very large drops in ionisation enrgy between elements 2-3 (He-Li), 10-11 (Ne-Na) and 18-19 (Ar-K)
-the second element in each of the pairs has its outermost electron in a new shell
-new shell is further from the nucleus and experiences more shielding, both decrease nuclear attraction to outermost electron
-lowering of attraction outweighs increased attraction from extra electron
how does ionisation energy provide evidence for subshells?
-all periods general trend of increasing ionisation energies
-increasing nuclear charge (more protons added)
- no extra shielding (same no. sjhells)
- decreasing atomic radius (same shell is pulled inwards
by extra protons)
-however drop between 2nd and 3rd elements 1st ionisation energies
-electron being removed higher energy orbital, easier to remove
trend in ionisation energy down a group?
Ie, depends on attraction nucleus to outermost electrons, affected by nuclear charge/shielding/atomic radius
1st ionisation energy decreases down the group
-down a group ATOMIC RADIUS INCREASES, outer shell electron further from nucleus, decreases nuclear attraction
-down a grp, no. internal energy levels increases, SHIELDING EFFECT INCREASES, decreases nuclear attraction
-down grp, nuclear charge increases but effect offset by increased shielding and aromic radius
-overall down a group attraction from nucleus to outermost electrons decreases, less energy to remove electron, 1st Ie decreases
general trend in 1st Ionisation energy across a period?
across a period:
-nuclear charge increases as no. protons increases> increases nuclear attraction to outermost electrosn
-increased nuclear charge>decreases atomic radius (outershell pulled closer)> increases nuclear attraction
-same period so shielding effects constant
- increased nuclear charge and decreased atomic radius increases attraction nuclueas to outermost electrons, more energy to overcome, increases 1st Ie.
exceptions to trend in 1st Ie across a period?
Be>B
slight decrease
-intro of 2p^1 subshell means that outermost electron in B experiences more shielding which decreases nuclear attraction
N>O, P>S
slight decrease
-moving fromp^3 to p^4 means electron being removed is now paired, electrons repel, decreases nuclear attraction, decreases 1st Ie if
explain why successive ionisation energies always increase?
when each electron is removed the outer hell is drawn closer to the nucleus. this decreases atomic radius which increases nuclear attraction, meaning more energy is required to remove the next electron
