Bonding Metallic Ionic Covalent The attraction between two oppositely charged entities
Jan 17, 2016
Bonding
Metallic
Ionic
Covalent
The attraction between two oppositely charged entities
Metallic•Formed between
Metals•Caused by
The metals’ outer shell electrons being delocalised within the lattice of the resultant metal ions•Actual Bond
The attraction between the +ve ions and the –ve electrons•Resultant structure:
A lattice of +ve ions (usually a cube) in a ‘sea’ of mobile delocalised electrons
Properties of Metals•High mp
Strong attraction between the +ve ion and –ve electron.
•Conductivity
Delocalised mobile electrons are able to move towards a +ve plate.
•Malleable
The layers of ions are able to move over each other.
Structure of metals• Giant Structures – HIGH MP / BP• Strong electrostatic forces of attraction ( metallic bond) between every
particle in the structure.
• How does the metallic bond get stronger?Increase the charge on the ion (stronger attraction between ion and
electron, so stronger metallic bond)
How Does MP/BP vary from Na Mg ?
• How do we increase the conductivity of metals?Increase the number of free electrons
How Does conductivity vary from Na Mg ?
Increases as Charge on ion increases from +1 +2
Increases as number of delocalised electrons per mole of lattice increases from 1 2
Ionic•Formed betweenMetals and Non-metals•Caused byThe metals’ outer shell electrons being transferred to the outer shell of the non-metal so that both achieve a full outer shell. The metal atom becomes a +ve ion and the non-metal atom a –ve ion•Actual BondThe attraction between the +ve ions and the –ve ions•Resultant structure: A lattice of +ve ions (usually a cube) surrounded by –ve ions
FLi
LiF
1
2
4 35
6
Represents the electrostatic attraction between opposite charged ions
Properties of ionic compounds•High mp
Strong attraction between the +ve ion and –ve ion.
•Conductivity in solids very poor
Attraction between the +ve and –ve ions hold the ions in a fixed position in the lattice.
•Conductivity in aqueous solutions or liquids very high
The ions become mobile so +ve ions move towards –ve plates and –ve ions move towards +ve plates.
Structure of Ionic compounds• Giant Structures – HIGH MP / BP• Strong electrostatic forces of attraction ( ionic bond)
between every particle in the structure.
• How does the ionic bond get stronger?
Increase the charge on the ions
(so that the attraction between the ions increase,
stronger ionic bond)
How Does MP/BP vary from NaCl Na2O MgO?
Increases as ions charge increase from +1/-1 +1/-2 +2/-2
Polarised Ionic•Small Highly charged cation
•(Nucleus of cation is not shielded greatly)
•Large anion
•Outer shell of anion is shielded from its’ nucleus
RESULT:
Outer shell of anion is attracted towards the nucleus of the cation
Distorts (polarises) the shape of the anion
Polarised ionic
9+3+
+ -
3+
+
-
53+53+
•Formed between
Non-Metals.•Caused by
The un-paired non-metals’ outer shell electrons are shared to form a electron pair between the two atoms.•Actual Bond
The attraction between the shared electron pair and the two nuclei.•Resultant structure:
Small discrete molecules or large giant structures
Covalent
What is really happening to the electrons during covalent bonding?
• GCSE = Shells overlap• A level Electrons are in (atomic) orbitals • In bonding the (atomic) orbitals overlap
QUESTION: WHAT DO THE ORBITALS LOOK LIKE?
S =
3 p = The 3 p orbitals are all at right angles to each other
Molecular orbitals
• Found in a covalent bond • The region of space in which the shared electron
pair in a covalent bond is located• Formed by the overlapping of the ATOMIC
ORBITALS• Two types
IN ALL SINGLE BONDS THE ELECTRON PAIRS ARE LOCATED IN SIGMA MO
IN ALL DOUBLE BONDS 1 ELECTRON PAIR IS LOCTAED IN A SIGMA MO AND THE OTHER ELECTRON PAIR IS LOCATED IN A PI MO
s and sSigma (σ) p and p end onSigma (σ) p and p side onPi (π)
Sigma MO Pi MOSigma MO
Properties of covalent compoundsSmall discrete Molecules
•Low mpWeak attraction between the small molecules
•Poor Conductivity No charged particles free to move
SIMPLE COVALENT (MOLECULAR) STRUCTURES
Properties of covalent compoundsGiant structures•High mp
Strong covalent bond between every atom in structure
•Low Conductivity
No charged particles free to move
EXCEPTION
Graphite – Delocalised electrons within each layer enables conduction.
Strange Compounds
• GCSE idea: After sharing electrons, atoms posses a full outer shell
• Dot cross diagrams for NaCl, CaO, CO2, HCl, SO2, SO3, BH3
• Note the problems with last three!
NaCl
CaO
CO2,
HCl
SO2
SO3
NOTE S has 10 electrons in outer shell
NOTE S has 12 electrons in outer shell
BH3
NOTE B has 6 electrons in outer shell
What causes Bonding to happen?Electronegativity
The ability of an atom to attract electrons towards itself in a covalent bond
AFFECTED BY?
a. No of shells (and hence the shielding of the nuclei)
b. No of protons in the nucleus
Electronegativity and Periodic Table• Across a PeriodNo of shells the sameShielding the sameNo of Protons increasesNuclear pull on electrons increasesElectronegativity increases
• Down a GroupNo of shells increasesShielding increasesNuclear pull on electrons decreasesElectronegativity decreases
MOST ELECTRONEGATIVE ATOM = FLEAST ELECTRONEGATIVE ATOM = CsELECTRONEGATIVITY OF C < ELECTRONEGATIVITY OF HALOGENSELECTRONEGATIVITY OF C = ELECTRONEGATIVITY OF HYDROGEN
Consider a Covalent bond X-Y
Elec X = Elec Y
Elec X > Elec Y
Elec X >>> Elec Y
ElecX >>>>>>>>>>Elec Y
Hence
Bond between a metal (low elec) and a non-metal (high elec) usually IONIC
Bond between 2 non-metals (similar elec) usually Covalent
Polarised covalent or ionic bonds ensue if somewhere in between
IONIC BOND= No overlap of shells
COVALENT Bond = Overlap of shells
Boron Compounds
• Boron is in group 3
• Expect ionic compounds (lose 3 electrons to achieve a full outer shell)
• Electronic configuration of B+3 = 1s2
• Polarisation so great: Covalent bonds ensue (even with BF3)
NB: B-F bond electron pair nearer to F atom once covalent
Aluminium compounds
• Al+3
• Electronic configuration = 1s22s22p6
• Most compounds form a degree of polarised ionic (except for AlF3)
• All aluminium tri-halides (except for AlF3) are covalent molecules
Covalent Bonding
1.Shared electron pair(s) of electrons between 2 nuclei
2.Pair usually results from 1 unpaired electron from one atom, pairing up with an unpaired electron from another atom. Doesn’t have to be the case
3.Two types of electron pairs are seen in molecules.
Square = Bonding Electron Pair Circle= Non-Bonding Electron Pair or LONE pair
Types of electron pairs in a molecule
Dative Covalent Bonding•2 H2 + O2 2 H2O•H+ + OH- H2O
One of the bonds in the water formed in the lower example, must involve a shared electron pair in which BOTH electrons originated from the same atom.Dative Covalent (or co-ordinate) bonding is where the shared electron pair originates from the same atom. Dative Covalent bonds behave like normal covalent bonds once formed
Requirements for Dative Covalent Bonding
1.An atom or an atom in a molecule or ion which has a lone pair of electrons available for donation. (often –ve charged)CALLED A NUCLEOPHILE
2.An atom or an atom in a molecule or ion which has an empty orbital which can accept an electron pair. (often +ve charged)
CALLED AN ELECTROPHILE
H+ and OH- example
•H+ ion (electronic configuration = 1s0)Has an empty s orbital to accept a lone pair
•OH- ion Has a lone pair which it can donate into the empty orbital
During the reaction, this happens to form a O-H dative covalent bond
O H H
Dative covalent bond
Boron compounds• All covalent compounds
• All have 6 electrons in outer shell of B
• All have an empty orbital
• All can accept an electron pairNN B
B
Shapes of Molecules•Valence Shell Electron Pair Repulsion theory (VSEPR theory)
•The shape of a molecule is dependent on the angle between the bonds in the molecule
•Electron Pairs will repel each other to the maximum extent
•Lone pairs repel to a greater extent than bonding pairs
•The bond angle (and hence shape) is therefore dependent on the number of bonding and lone pairs surrounds the central atom in the molecule
Working out shapes• Draw dot cross diagrams (or displayed formulae)• Work out the number of bonding and lone pairs around the central atom• Treat multiple bonds as 1 bonding pair
• H2
• BeI2 (covalent)• AlCl3 (covalent)• CF4 • H2S• PH3
• SiH4
• SO3
Molecule Number of electron Pairs
Max Distance repelled
Number of lone pairs
ActualBond angle
Shape
H2 1 180 0 180 Linear
BeI2 2 180 0 180 Linear
AlCl3 3 120 0 120 Trigonal Planar
CF4 4 109.5 0 109.5 Tetrahedral
H2S 4 109.5 2 104.5 Bent
PH3 4 109.5 1 107 Pyramidal
SiH4 4 109.5 0 109.5 Tetrahedral
SO3 3 120 0 120 Trigonal Planar
SF6 6 90 0 90 Octahedral
BeI2
CO2
AlCl3BH3
CF4
CH4
SO3
SF6
NF3
PH3
H2S
H2O
Complex molecules
• C2Cl4• CH3CH2OH
• Work out the bond angle (and hence shape around each central atom)
Intermolecular Forces
•What makes a substance a solid at RT?•Why is F2 a gas, while I2 is a solid?•How can you get liquid F2?•Why does H2O boil at 100oC while H2S boils at -60.3oC?
Intermolecular forces: THE FORCE OF ATTRACTION BETWEEN MOLECULES
Still a force of attraction between 2 oppositely charged entities
VAN DER WAALS FORCES
• Consider two F2 molecules
• What force of attraction could exist between them to enable F2 to liquefy?
VAN DER WAALS FORCES
• Electrons in a bond are not static
• Closer to one atom Temporary Dipole
• Causes dipole to form in neighbouring molecules Induced Dipoles
• Attraction between these two dipoles = VDW
Permanent Dipole- Dipole
• Molecules containing atoms of different electronegativity eg HCl, CH3Cl,
• Permanent dipoles are present
• Interactions occur between the δ+ve atom in one molecule and the δ-ve atom in the other molecule
Molecules without permanent dipoles when expected
• Symmetrical molecules
O=C=O
Dipoles cancel out
Hydrogen Bonding
• Especially large dipoles in molecules
• Especially large dipole-dipole attraction
• Called H-Bond
• Formed between molecules containing OH, NH or HF bonds
Drawing Hydrogen bonds
• Need at least 2 molecules
• Need to draw on dipoles on O and H atoms
• Need to draw lone pairs on O
• Draw straight line (often dotted) between lone pair on O to H on DIFFERENT Molecule
Water and DNA
WATER
• High MP than expected
• High surface tension
• Low density of ice
DNA
• Helical structure
• Strands held together
H-Bonds need to be broken
H-Bonds on surface pull molecules on surface in
H-Bonds hold Water molecules apart in Solid Water
Boiling Points•Non-Polar compoundsLOW as VDW weakBP increases as no of electrons in molecules increasesTemporary dipoles / induced dipoles increaseForce of attraction between molecules increase
•Polar compoundsHIGHER AS permanent dipole dipole attractions stronger than VDWNeed to look for a bond in the molecule containing two atoms of different electronegativities.
•Compounds containing OH, NH or HF bondsMuch higher as Hydrogen bonds much stronger than permanent dipole dipole
BUT• Intermolecular forces only found in simple molecular
structures and are MUCH WEAKER than the covalent, ionic or metallic bonds found in GIANT STRUCTURES
BP 0CNaCl 1250MgO 2390Cu 2600Carbon 4500
CH4 -79
CH3Cl -25
H2O 100
Giant Structures: BREAK STRONG CHEMICAL BONDS TO MELT
Simple Molecular Structures: BREAK Weak intermolecular forces to melt
In solid Methane Particle = CH4
In Gaseous Methane Particle = CH4
No Chemical bond is broken on melting / boiling