bonding topic 4. terms covalent bonding bonds –breaking them takes energy –making them gives off...
TRANSCRIPT
BONDING
TOPIC 4
Terms
Covalent Bonding• Bonds– Breaking them takes energy– Making them gives off energy
• Exothermic– More energy is given off than put in
• Endothermic – More energy is absorbed than given off
• Intramolecular Forces– Forces within molecules (ionic, covalent and metallic)
• Intermolecular Forces (IMF)– Forces between particles
Metal: K Non-Metal: Cl
Ionic Bonding
• If the electronegative difference between the atoms involved is =>1.8– There are always exceptions to this rule!
• Will conduct electricity in its molten or aqueous state (This test proves ionic)
Less e- = Less e- repulsion More e- = e- more repulsion.
+ -
Intramolecular Forces
Ionic Bonding• Just use the valence shell
• Be sure to include square brackets and charge after electron exchange.
Drawing Ionic Bonding
Na Cl
XElectrons are
in pairs
+
Special Note: The ionic bond is the electrostatic attraction between oppositely charged ions!
-Lewis Dot Diagram
Lewis Diagrams
Lewis Dot diagrams us the atoms valance shell electrons
Combine
C
Cl
Be
Br
Al
F
Fe
Cl
Mg
O
Intramolecular Forces
NaCl
• When in molten or aqueous state, ionic substances WILL conduct electricity, by the movement of (+) and (-) ions.
• This is different from how METALS conduct electricity!
Decomposition
2Na+(aq) + 2Cl-
(aq) 2Na(s) + Cl2(g)CA
THO
DE
(-)
ANO
DE
(+)
-
++
+
++
+
-
--
-+ +
+ +
+ +
Conductivity is FINITE
Metal: K Non-Metal: Cl
Ionic Compounds
+ - + - +-
-+
+ - ++
--
-+-+- +
+ +-+-
• No bonds are made!!!
• Static attractions holds them together. (opposites attract)
• When a force is applied, ionic compounds will make a clean break.
Force
Like charges repel
Giant Ionic Lattices
• Physical characteristics• Hard and brittle• Solid doesn’t conduct Electricity• More soluble in water than other solvents• High MP and BP
CationAnion
Table SaltNaCl
Cubic or IsometricGiant Ionic Lattices
CassiteriteSnO2
Tetragonal
Giant Ionic Lattices
AgagoniteCaCO3
Orthorhombic
• Also found in mollusk shells and coral
Giant Ionic Lattices
BerylBe3Al2(SiO 3)6
Hexagonal
Giant Ionic Lattices
QuartzSiO2
Trigonal
Giant Ionic Lattices
BerylBe3Al2(SiO 3)6
Ionic BondingGiant Ionic Lattices
Copper(II) SulfateCuSO4
Triclinic
Giant Ionic Lattices
Intramolecular Forces
Multiple Ions• Transition metals can have
multiple ions.• Ones you should know.
Transition Metals
Cu2+Cu+Fe2+ Fe3+
Copper(I) Oxide Iron (III) OxideIron(II) Oxide Copper(II) Oxide
Ions
Polyatomic Ions• Be sure to review your polyatomic
ions!!!
Reminder
NO3- SO4
-2 PO4-3
OH- CO3-2
NH4+
HCO3-
Covalent Bonding
Topic 4
Intramolecular Forces
COVALENT BONDING
• If the electronegative difference between the atoms involved is <1.8
• Will NOT conduct electricity• Electrons are shared
Covalent Bonding
X
H Cl
3.02.1
Differences |3-2.1|
=0.9
Special Note: The covalent bond is the electrostatic attraction between pairs of e- and positively charged nuclei!
Questions
• For ionic compounds to form the valance shells of both metal and non-metal must be full!!
Na Ca
Li Na
+
+
+
+
Cl
O
CO3
SO3
• What is the chemical formula? • What is the names for each?
K + NO3
Review
Intramolecular Forces
COVALENT BONDING
• Structural formula
• Lewis structure
Covalent Bonding
H C
H
H
H Cl
X
XX
X
H
H
H
H
Intramolecular Forces
COVALENT
BONDING
• 1) Sum all valence e-
• 2) Subtract 2e- for every bond• 3) Place e- around periphery atoms to
form octets. The remaining around central atom
• 4) All atoms MUST be paired!!!!!!
Lewis StructuresH2O
1
1
6
8
H
H
O
- 4 = 4
Hydrogen can only hold 2e- remaining must be paired on
Oxygen
Intramolecular Forces
Covalent Bonding
• Draw the following Lewis structures
• H2 Cl2• O2 N2
• HCN C2H6
• C2H4 C2H2
Lewis Structures
HL: PCl5, PCl4+, PCl6
-
and XeF4
Intramolecular Forces
Covalent Bonding
• Coordinate or dative covalent bonds• When both e- are shared from the same
atom. (Not one from each as before)• Occurs when a non bonding e- pair
donates an e- to an e- deficient atom.
Special Lewis Structures
H
H
NH
+
Electrophile
H
Lone pair of e-
+
Intramolecular Forces
Covalent Bonding
• Draw the following Lewis structures• CO
• H3O+
Special Lewis Structures
Intramolecular Forces
CO32-
• More bonds = more strength & shorter bonds
• Resonance structures – Bond length is longer than a double bond but
shorter than a single bond
Length, Strength & Hybrid Resonance
CO O
O2-
CO O
O2-
CO O
O2-
Don’t forget to show the e- pairs!!
Intramolecular Forces
CO32-
• Compare the two molecules
• Ethyne has stronger and shorter bonds• • C=O bond is stronger and shorter due to
Oxygen being more electronegative
Length & Strength
Ethene
R = Functional Group
CR OH
OC
H
H
C
H
H
Carboxylic Acid
C HCH
Ethyne
Intramolecular Forces
Covalent
Bonding
• Non-Metals are fighting for e-
• Atom with larger electronegativity will hold the e- closer to itself.
• Atoms become slightly charged.
Bond Polarity
ClXH
δ-δ+
Dipole Moment
Intramolecular Forces
Covalent
Bonding
• BF3
• Actual structure: Boron is e- deficient• This is known because of its reactivity towards
electron rich molecules such as NH3
• CNOF all obey the octet rule.
Exceptions to the Octet Rule
FF
B
F
• SO42-
• Single bonds (8 e- around S)• Double bonds (12 e- around S)
• Formal Charge = (# valence e- on free atom) – (# valence e- assigned to the atom in the molecule)
• (Valence e-)assigned = (# lone pair e-) + ½ (# of shared e-)
• 1) Molecules attempt to achieve Formal Charge as close to 0 as possible.
• 2) Any negative Formal charge will reside on most electronegative atom.
Formal Charge
Covalent Bonding
Intramolecular Forces
Intramolecular Forces
Covalent
Bonding
• VSEPR (Valence Shell Electron Pair Repulsion)
• Paired e- attempt to get as far away from each other as possible.
• Multiple bonds still count only as 1 pair!!
VSEPR (shape)
OO C
FF
B
F
3 Pairs of e-
120o
2 Pairs of e-
180o
OO
C
O2-
Intramolecular Forces
Covalent
Bonding
• Tetrahedral
• Lone pair e- have increased charge density and require more room
• More repulsion from lone pair will decrease bond angle.
VSEPR
H
H
C
H
H
4 Pairs of e-
109.5o
H
H
N
H
Lone pair107o
H
O
H
Lone pair104.5o
Intramolecular Forces
Covalent Bonding
• Predict the shape AND bond angles
• H2S PbCl4 H2CO SO2
• NO3- PH3 NO2
-
• NH2- POCl3 CO2
Home Work
HL VSEPRMolecule Shape Total valance
electrons Bond Pairs Non Bonding
Electron pairAngle
BeF2 Linear 180
BeF3 Triangular Planar
120
SO2 Bent 117
CH4 Tetrahedral 109.5
NH3 Trigonal pyramidal
107
H2O Bent 104
HL VSEPRMolecule Shape Total
Valance electrons
Bond Pairs Non Bonding Electron pair
Angle
PCl5 Triangular Bipyramidal
90 & 120
SF4 Seesaw 90 & ≈117
T-Shape 90
CF6 Octahedral 180
IF5 Square Pyramidal
90
XeF4 Square Planar
≈88
Expanded Valance Shell (14.1)
• Molecules with more than 8 electrons• Electron promotion:
Dipole Moment
Covalent Bonding
• Polarity effects state change (physical change)
• Unequal sharing causes a dipole moment to form
• Q: Why is BF3 non-polar whereas PF3 is polar?
Molecule Polarity (4.2.6)
δ+
2δ-
δ+H
O
H
ClH δ-
δ+
Non Polar
H
H
C
H
δ-
δ+ H
H
C
Cl
Cl
Hybridization (14.2.2)• Sigma bond: σ (single bond)
– Axial overlap of orbital’s
Cl
1s1
H
2px2 py
2 pz2
Hybridization (14.2)• Sigma bond: σ (single bond)
– Axial overlap of orbital’s
ClCl
Hybridization (14.2)• Pi bond: π(Double bond, one σ bond)
– Parallel overlap of orbital’s
NN OO
Hybridization (14.2.3)• Hybridization electron promotion
– New Orbital sp3
2px2 py
2 pz2
2s2Ground
StateC
Excited State
4 Equal orbital`s capable of holding a maximum of 2 electrons each
Hybridization (14.2)
• How to determine Hybridized orbital`s– Look at the shape
Shape High Electron dense regions
Hybridized Orbital
sp 2 Tetrahedral
Sp2 3 Trigonal planar
sp3 4 Linear
sp3d 5 Trigonal bi-pyramidal
sp3d2 6 Octahedral (Square bi-pyramidal)
Allotropes
Giant Covalent
• 1) Diamond (Tetrahedron, localized e-)– Very hard and does not conduct electricity
• 2) Fullerenes (C60) Hexagonal and pentagonal rings– Nanotubes
Carbon
C
CC
C
C
Allotropes
Giant Covalent
• 3) Graphite (Planar, delocalized e-)– Weak pi bonding between sheets cause it to
conduct electricity and be slippery.– Bonds are shorter than a tetrahedral due to the
pi bonding
Carbon
C
C
CC
C
C
HL: sp hybridDelocalized electrons
able to move
Weak Pi Bonds
C6H6• Planar, delocalized e-
– Regular bonding would predict an alternating double bond (Resonance structure)
– Hybrid theory shows sp2 configuration
Benzene (14.3)
C
C
CC
C
C
Pi bonds overlap allowing for electrons to be delocalized over the entire molecule.
Intramolecular Forces
SiliconTetrahedron Configuration
Similar to diamond
Silicon
Si
Si
Si
Si
Si
Si
Si
SiSi
Si
Intramolecular Forces
Quartz
• Single bonds formed between Oxygen to satisfy the octet.
• HL: Less overlap in the P-sub orbital due to atomic size difference therefore Pi bonds do not form.
Silicon & Silicon dioxide
SiOO
O
O
SiO2 but based on a network
of SiO4
Metallic Bonding
Topic 4
Intramolecular Forces
Metallic Bonding
• In solid state
• Outer e- are delocalized and free to move about
• Bond is a result of electrostatic attraction between Fixed positive metal ions and delocalized e-
Metallic Bonding+ +
+
+
+
+ +
+
+
++
++
+
+ ++-
-
-
- --
Conductivity is INFINITE
-
-
Sea of electrons
Metallic Bonding
• The ability for a material to be pounded into thin sheets.
• Aluminum Foil
• Swords and Folding
+ +
+
+
+
+ +
+
+
++
++
+
+ ++
-
-
-
- --
-
-
MalleabilityPhysical Properties
Ductility
Metallic Bonding
• The ability for a material to be pulled into wire
• Or in this case extruded into a wire
Physical Properties
+
+
+ +
+
+
+
+
++
+
+
Electrons have been excluded
Metallic Bonding
• Because e- can move easily it can conduct energy. (Heat or electricity)
• MP related to attractive force (between atoms) • 1) Size of Cation(+)
• 2) # of valence e-
• 3) Atom packing
• Size increases MP decreases: • Giant Covalent substances have very high
mp
Physical Properties
Metallic Bonding
• Same element but different structure• Carbon
• Diamond• Graphite• Fulluron
Allotrops
INTERMOLECULAR FORCES
Topic 4
Intermolecular Forces
IMF• Van der Waals Forces
van der Waals’ ForcesIntermolecular Forces (4.3.1)
Charge Induction Charge Induction
d+
d- d+
d- d-d+
Intermolecular Forces
IMF• Polar molecules (polar covalent) have
slightly charged ends• Opposites attract.• Large electronegative difference =
stronger attraction.
Dipole-Dipole (4.3.1)
H
H
C
Cl
Cl
d-
d+
H
H
C
Cl
Cl
d-
d+
Intermolecular Forces
IMF• Hydrogen Bonding (F, O or N bonded to H)
• Due to small size and high electronegativity of non metals
• Creates a large charge difference • Basically a super strong dipole-dipole bond
van der Waals’ ForcesHydrogen Bonding (4.3.1)
d+d+ H
O
H
d-
d+
d+
H
OH
d-
Get a picture of group 4,5,6,7 boiling points for hydrides
Key question is why does water have an abnormally high BP?
H bonding with O, F and N
IMF• Phase change when IMF are overcome• Be sure to explain using the words IMF and
how they affect the bonds BETWEEN particles.• Van der Waals’ Forces are ALWAYS present!!!
Boiling Point Trends (4.3.2)Intermolecular
Forces
• Van der Waal’s: Lowest MP, Non polar• Butane (C4H10)
• Dipole-dipole: Slightly miscible• Propanone C3H6O
• Hydrogen Bonding: Miscible with polar substances• H2O
• Ionic Bonding: Only conducts electricity when liquid or aqueous. (Decomposition when it does)• NaCl
• Metallic Bonding: Conducts electricity, not water soluble, MP regulated by, valance, size and packing.• Fe
• Giant Covalent: Highest MP, Insoluble in both non-polar and polar solvents. Does not conduct electricity except for graphite.• Diamond and Graphite (Allotropes)
Physical Properties
Increasing Melting Point
Bonding Questions
• Compare the following for B.P
• HF and HCl• H2O and H2S
• NH3 and PH3
• CH3OCH3 and CH3CH2OH
• CH3CH2CH3, CH3CHO and CH3CH2OH
HL Material
Hybridization (14.2)
• Sigma bond: σ (single bond)– Axial overlap of orbital’s
Hybridization (14.2)
• Sigma bond: σ (single bond)– Axial overlap of orbital’s
Hybridization (14.2)
• Sigma bond: σ (single bond)– Axial overlap of orbital’s
Hybridization (14.2)
• Sigma bond: σ (single bond)– Axial overlap of orbital’s
Lattice Formation
• Where the heat comes from
• Route 1: A + B + C + E• Route 2: F
• Hess’s law: A + B +C + E = F+107 + 122 + 496 + (-349) + E = -411E = -787 kJ mol-1
Intramolecular Forces
NaCl
• 1) Production of Gaseous atoms• 2) Formation of Gaseous ions• 3) Production of solid ionic lattice
Lattice Enthalpy
Na(s) + ½ Cl2(g) Na+Cl- or NaCl
1) Na(s) Na(g) ½ Cl2(g) Cl(g)
3) Na+(g) + Cl-
(g) NaCl(s)
2) Na(g) Na+(g) + e- Cl(g) + e- Cl- (g)
Born-Haber CycleNa+
(g) + Cl-(g)
Na(s) + Cl2(g)
Na(g) Cl(g)
NaCl(s)
ΔHθI.E.
1st Ionization of Na+496 kJ mol-1
ΔHθat
Atomization of Na-107 kJ mol-1
ΔHθat
Atomization of Cl+122kJ mol-1
ΔHθE.A.
1st electron affinity of Cl-349 kJ mol-1
ΔHθf
Formation of NaCl-411 kJ mol-1
Lattice Enthalpy
BA
DC
E
FExothermicEndothermic
OO C
Spare Parts
O C
ClXH δ-
δ+
Cl
HH
C
H
N
H
H
NH
+H
-
H
O
H
H
H
C
H
H
C
C
CC
C
C