Chem 102 Yerkes University of Illinois Lecture
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Transcript of Chem 102 Yerkes University of Illinois Lecture
Chemical Bonding
Why do bonds form? to lower the potential energybetween positive and negative charges
positive charges protonscations
negative charges electronsanions
metals
Periodic Tablelose e-
non-metalsgain e-
Lewis electron-dot symbols
element symbol = nucleus + core e-
one “dot” = valence e-
metals dot = e- it loses to form cationnon-metal unpaired dot = e- paired through
e- gain or sharing
metal + non-metal
Ionic bonding
Na (s)2 + Cl2 (g) NaCl (s)2
[Ne] [Ne] [Ne] [Ar]
Ca O: :..
..
[Ar]4s2[He]2s2 2p4
Ca2+ O2-
[Ar] [Ne]
Ca(s)2 2+ O2(g) CaO (s)
Na. + Cl-::
::+ Cl:
:: . Na+
3s1 3s23p5
metal + non-metal
Ionic bonding
low Ionization Energy
lose 1 or 2 valence e-
high Electron Affinity
gain e-
electron transfer takes place
electrostatic attraction between cation and anion
formula = ratio of anions to cations
e-
+-
Ionic sizes
isoelectronic series
46 e-
+49 +50 +51 ions get smaller
same # electrons
+-
metal + non-metal
Ionic bonding
Na (s)2 + Cl2 (g) NaCl (s)2
metal + non-metal
Ionic bonding
Na (s)2 + Cl2 (g) NaCl (s)2
exothermic heat given off
Ionization Energy Na + 496 kJ/mol
Electron Affinity Cl -349 kJ/mol
Lattice Energy
Na+
Cl-
E =d
-787 kJ/mol
Na++ Cl- NaCl
negative
Coulomb’s law
k Q1Q2
-640 kJ/mol
Ionic solids
cation + anion+ -
metal non-metal
lithium oxygenLi O+ 2-
+Li O lithium oxide
magnesium nitrogen+
Mg N2+ 3- Mg N magnesium nitride3 2
strong interactions (ion-ion)
sodium chlorineNa Cl+ -
+NaCl sodium chloride
2
high melting points
801o C
> 1700oC
Transition metals
more than 1 form except Ag+
Zn2+
Cd2+
Al3+
aluminum+sulfurAl3+ S2-
Al S2 3
aluminum sulfide
manganese oxygen+Mn1+ Mn2+ Mn3+ Mn4+
Mn4+ O2- MnO2
manganese(IV) oxide
Mn3+ Mn2O3
manganese(III) oxide
O2-
+ non-metalnon-metal
Covalent bonding
electrons shared between atoms
high Ionization Energies
high Electron Affinities
electron density between the atoms
distance between atoms = bond length
formula = actual # atoms
F.: ::
:F.:
:
+ .
: :F .:
:
F:
:
[He]2s22p5 [Ne]
F:
:: F:
::
e- not used in bonding lone pairs
Lewis structure
shared equally between F
H O.
::. .
:O:
. .H. H.
oxygen 2 lone pairsbonding pair
1s1 [He]2s22p4
[He]
shared e- bonding pair
not shared equally
[Ne]
+ non-metalnon-metal
Covalent bonding
H. + H . H ..H1s1 1s1 [He]
Covalent compounds
share valence electrons = chemical bondscarbon + chlorine
C Cl -4+
CCl4
carbon chloridetetra
1 mono2 di3 tri4 tetra5 penta6 hexa7 hepta8 octa
nitrogen + oxygenN O 2-?
NO nitrogen monoxide 2+
NO2 nitrogen dioxide 4+
N2O4 dinitrogen 4+
non-metal + non-metal
tetroxide
Covalent compounds
H
N H1-3+ NH3 nitrogen trihydride ammonia
O2-1+ H2O dihydrogen monoxide water
weak forces low m.p. 0.0oC
Polyatomic ions
NH4+ ClO3
-
Table 2.5 p. 62
ammoniumOH- hydroxideNO3
- nitrateSO4
2- sulfatePO4
3- phosphate
chlorateMnO4
- permanganateCrO4
2- chromateCO3
2-carbonate
metal + metalMetallic Bonding
metals valence e- well shielded
low Ionization Energy
low Electron Affinities
share valence e- not localized between atoms
delocalized
move freely throughout metal
Nae- “sea”
(nucleus and core e-)(valence e-)
Electronegativity
ability of an atom in a moleculeto attract e- to itself
related to Ionization EnergyElectron Affinity
Pauling scale
non-polar covalent 0.4polar covalent 0.5-1.8
NaCl 2.1 ionic 801oC
AlCl3 1.5 polar covalent 178oCPCl3 0.9 polar covalent 76oCCl2 0.0 covalent
ionic > 1.8
C HC O-+
Li2O
BeCl2 1.5 polar covalent 405oC Boiling/Melting points