96
BONDING
BONDING
CHEMICAL BONDING
Chemical Bonds
Ionic(Metal &
Non-Metal)
Covalent(2 or > non-
metals)
Intermolecular forces
Hydrogen bond
Van Der Waals forces
Metallic
CHEMICAL BONDINGSubstances
Giant Structure- High melting &
boiling points
Ionic compounds – forms giant ionic
structures
Giant covalent structure
- Atoms are held by strong covalent
bonds-No van der waals
forces
Metals
Simple molecular structure
- Low melting & boiling points
Covalent- Molecules are held
by weak van der waals forces
Diamond Silicon dioxide
Graphite
Ionic bonding typically occurs between metal and non-metal. E.g. Barium fluoride, BaF2
The reactivity of metals and non-metals can be assessed using electro-negativity
Electro-negativity ability of an atom in a covalent bond to attract ___________________________ to
itself.
PREDICTING THE TYPE OF BONDING FROM ELECTRONEGATIVITY VALUES
shared paired of electrons
Metals generally have low electronegativity values, while non-metals have relatively high electronegativity values.
Fluorine, which has the greatest attraction for electrons in bond-forming situations (highest E value).
Type of Bond Electronegativity Difference
Ionic 2.0 – 4.0
Polar Covalent 0.5 – 1.9
Nonpolar covalent 0 – 0.4
Polar covalent bonds are covalent bonds with ionic character.
Electrons are not shared.E.g. Na+ Cl- , electron is transferred.
Electrons are equally shared.E.g.Cl-Cl
Ionic bond
Non polar (pure covalent) bond
Polar covalent bond
Electrons are not equally shared.E.g.
Atoms have differentelectronegativity values
-+
ClH
What type of bond is the following?(a) N (3.0) and H (2.1)(b) H (2.1) and H(2.1)(c) Ca(1.0) and Cl(3.0)(d) Al (1.5) and Cl(3.0)(e) H (2.1) and F(4.0)
EXAMPLEType of Bond Electronegativity Difference
Ionic 2.0 – 4.0
Polar Covalent 0.5 – 1.9
Nonpolar covalent 0 – 0.4
DIPOLE
No bond is purely ionic or covalent..they have a little bit of both characters.
When there is unequal sharing of electrons a dipole exists.
Dipole- is a molecule that has 2 poles or regions
with opposite charges.- is represented by a dipole arrow towards the
more negative end.
ATTRACTIONS BETWEEN MOLECULES
Intermolecular attractions are weaker than ionic, covalent, and metallic bonds
Besides ionic, metallic, and covalent bonds, there are also attractions between molecules
There are 2 main types of attractions between molecules: Van der Waals and Hydrogen
VAN DER WAALS FORCES Van der Waals forces consists of the two
weak attractions between molecules
1. dipole interactions – polar molecules attracted to one another
2. dispersion forces – caused by the motion of electrons (weakest of all forces)
HYDROGEN BOND Hydrogen Bonds are forces where a hydrogen
atom is weakly attracted to an unshared electron pair of another atom
HYDROGEN BOND
This other atom may be in the same molecule or in a nearby molecule, but always has to include hydrogen.
INTERMOLECULAR HYDROGEN BONDING
Hydrogen Bonds have about 5% of the strength of an average covalent bond
Hydrogen Bond is the strongest of all intermolecular forces
INTRAMOLECULAR HYDROGEN BONDING
INTERMOLECULAR ATTRACTIONS
A Network Solid contains atoms that are all covalently bonded to each other
A few solids that consist of molecules do not melt until the temperature reaches 1000ºC or higher called network solids (Example: diamond, silicon carbide)
• Melting a network solid would require breaking bonds throughout the solid (which is difficult to do)
http://library.thinkquest.org/C006669/data/Chem/bonding/inter.html
The bonding pair of electrons spends most of its time between the two atomic nuclei.
screening the positive charges from one another and enabling the nuclei to come closer together.
Negative charge on the electron pair attracts both nuclei and holds them together in a covalent bond.
When two atoms are chemically bonded, the two atoms close together have less energy and therefore are more stable than when separated.
Energy is given off by the atoms to form a bond, and energy must be supplied (absorbed) to break the bond.
A covalent bond is the result of electrostatic attraction between the nuclei of the 2 atoms and the pair of shared electrons.
ELECTRON DOT (LEWIS) STRUCTURE
VALENCE ELECTRONS FOR ELEMENTS
VALENCE ELECTRONS AND NUMBER OF BONDSNumber of bonds elements prefers depending on the number of valence
electrons. In general -
X
Family # Covalent Bonds*
Halogens
F, Br, Cl, I
Calcogens
O, S
Nitrogen
N, P
Carbon
C, Si
O
N
C
1 bond often
2 bond often
3 bond often
4 bond always
LEWIS STRUCTURE, OCTET RULE GUIDELINES
When compounds are formed they tend to follow the Octet Rule.Octet Rule: Atoms will share electrons (e-) until it is surrounded by eight valence electrons.
Rules of the (VSEPR) game-i) O.R. works mostly for second period elements.
Many exceptions especially with 3rd period elements (d-orbitals)
ii) H prefers 2 e- (electron deficient)
iii) :C: N: :O: :F:4 unpaired 3unpaired 2unpaired 1unpaired up = unpaired e-
4 bonds 3 bonds 2 bonds 1 bond
O=C=O NN O = O F - F
iv) H & F are terminal in the structural formula (Never central)
.. . .. ..
.
ATOMIC CONNECTIVITYThe atomic arrangement for a molecule is usually given.
CH2ClF HNO3 CH3COOH H2Se H2SO4 O3
H C F
Cl
HH
N OO
OH O S O H
O
OO O O
H C C
O H H
H O H Se H
In general when there is a single central atom in the molecule, CH2ClF, SeCl2, O3 (CO2, NH3, PO4
3-), the central atom is the first atom in the chemical formula.
Except when the first atom in the chemical formula is Hydrogen (H) or fluorine (F). In which case the central atom is the second atom in the chemical formula.
Find the central atom for the following:
1) H2O a) H b) O 2) PCl3 a) P b) Cl
3) SO3 a) S b) O 4) CO32- a) C b) O
5) BeH2 a) Be b) H 6) IO3- a) I b) O
Count the no. of valence electrons.(i) If the species has a –n charge, add n to the electrons(ii) If the species has a +n charge, subtract n from the electrons
Draw a skeletal structure.(i) If C is present, place C at the centre.(ii) If C is not present, place the LEAST electronegative atom at the centre.Note: H is never the in the center.
Complete the octets of the outer atoms (except for H) by adding lone pairs of electrons (including the 2 electrons shared with the central atom)
If there are any electrons left over, place them on the central atom as lone pairs.
If the central atom does not have a complete octet, rearrange lone pairs on the outer atoms to form double bonds between the central and outer atoms. Continue doing until the central atom’s octet is satisfied.
If the species is charged, place it inside brackets and write the charge outside the brackets.
RULES FOR DRAWING LEWIS STRUCTURES
Draw Lewis structures of the following molecules:
(a) H2O , NH3 , CO2 , OCl2 , PCl4+
(b) SO2 , NO+ , OCN- , COF2 , CO32- , NO2
- , O3
(c) BeCl2 , BH3 , PCl5In which of the above obey the octet rule?
LEWIS STRUCTURESBond pair _____Lone pair . .
1.Molecules with an odd number of electrons
2.Molecules in which an atom has less than an
octet
3.Molecules in which an atom has more than an
octet
EXCEPTIONS TO OCTET RULE
1. Odd Number of Electrons
NO Number of valence electrons = 11
N O N O Resonace Arrows
O O O OOxygen is a ground state"diradical"
NO2Number of valence electrons = 17
O2
Resonance occurs when more than one valid Lewis structure can be written for a particular molecule (i.e. rearrange electrons)
Molecules and atoms which are neutral (contain no formal charge) and
with an unpaired electron are called Radicals
N OO N OO N OO
2. Less than an Octet
Includes Lewis acids such as halides of B, Al and compounds of Be
BCl3
Group 3A atom only has six electrons around it
However, Lewis acids “accept” a pair of electrons readily from Lewis bases to establish a stable octet
Cl
AlCl
Cl
N
H
H
H
Cl
AlCl
Cl
N
H
H
H
+
Lewis acid Lewis base salt
+_
B
Cl
Cl Cl
AlX3
Aluminium chloride is an ionic solid in which Al3+ is surrounded by six Cl-. However, it sublimes at 192°C to vapour Al2Cl6 molecules
AlCl
Cl Cl
ClAl
Cl
ClB2H6
A Lewis structure cannot be written for diborane. This is explained by a three-centre bond – single electron is delocalized over a B-H-B
BH
H H
HB
H
H
Octet Rule Always Applies to the Second Period = n2 ; number of orbitals
2s, 2px, 2py, 2pz ---orbitals cannot hold more than two
electrons
Ne [He]; 2s2, 2px2, 2py
2, 2pz2
n = 2
n = 3
THIRD PERIOD ; N2 = 32 = 9 ORBITALS
Ar [Ne]; 3s2, 3px2, 3py
2, 3pz2 3d0 3d0 3d0 3d0 3d0
n = 3
3. More than an Octet
PCl5
Elements from the third Period and beyond, have ns, np and unfilled nd orbitals which can be used in bonding
P : (Ne) 3s2 3p3 3d0
Number of valence electrons = 5 + (5 x 7) = 40
P
Cl
ClCl
ClCl
10 electrons around the phosphorus
SF4
S : (Ne) 3s2 3p4 3d0
Number of valence electrons = 6 + (4 x 7) = 34
SF
F
F
FThe Larger the central atom, the more atoms you can bond to it – usually small atoms such as F, Cl and O allow central atoms such as P and S to expand their valency.
MULTIPLE BONDS Bond StrengthTriple bonds > Double bonds > Single bonds
Page 93
(1)The attraction between the 2 nuclei for 3 electron pairs in a triple bond is > that for 2 electron pairs in a double bond which is > than that for 1 electron pair in a single bond.
(2) Triple bonds are shorter due to greater attraction between the bonding electrons and the nuclei with more electrons in the bond.
StrengthTriple bonds > Double bonds > Single
bonds LengthSingle bonds > Double bonds > Triple
bonds
BOND STRENGTH AND LENGTH OF COVALENT BONDS
Bond Type Length (nm)
Strength (kJmol-1 )
C-C 0.154 348
C=C 0.134 612
CΞC 0.120 837
In some molecules and polyatomic ions, both electrons to be shared come from the same atom forming the coordinate or dative bond.
Carbon monoxide (CO) can be viewed as containing one coordinate bond and two "normal" covalent bonds between the C atom and the O atom.
COORDINATE (DATIVE) BONDING
Page 94How do you draw the Lewis structure?
DISSOLVING HYDROGEN CHLORIDE GAS IN WATER Something similar happens. A hydrogen
ion (H+) is transferred from the chlorine to one of the lone pairs on the oxygen atom.
The H3O+ ion is variously called the hydroxonium ion.
Other examples: The reaction between ammonia and
boron trifluoride, BF3
In BF3, there are only 6 electrons in the outer shell of boron. There is space for the B to accept a pair of electrons.
BOND POLARITY Due to difference in electronegativity value between
the 2 atoms in the bond.
Unequal distribution of electron density results in small charges on the atoms
( δ+ and δ- )
Example
A dipole is established when two electrical charge of opposite sign are separated by a small distance.Dipole
moment
Element F O N Cl C H
Electronegativity
4.0 3.5 3.5 3.0 2.5 2.1
NON-POLAR MOLECULE A molecule can possess polar bonds and
still non-polar. Check the Geometry of the molecule: The polar bonds are arranged
symmetrically so as to give zero net direction of charge.
i.e. Overall dipoles cancel so that there is no overall dipole.
NON-POLAR & POLAR COVALENT BONDSNon-polar Covalent
bond No difference in
electronegativity value – bond consists of
2 ____________ atoms. _______ net charge.
Examples :
Polar Covalent bond Due to the difference in
electronegativity value – bond consists of
2 ____________ atoms. _______ net charge.
Examples:
In the water molecule, (i) O-H bonds are significantly polar(ii) The bent structure makes the
distribution of those polar bonds asymmetrical.
POLAR MOLECULE
Some molecules have very low polarity - so low as to be regarded as non-polar,
Name of molecule
Formula Polarity of molecule
Hydrogen chloride HCl Polar
Water H2O Polar
Ammonia NH3 Polar
Benzene C6H6 Non-polar
Boron trichloride BCl3 Non-polar
Methane CH4 Non-polar
Bromobenzene C6H5Br Polar
Carbon dioxide CO2 Non-polar
Sulfur dioxide SO2 Polar
Tetrachloromethane
CCl4 Non-polar
For CO2 each C-O bond is polar since O is moreelectronegative than C. Why is the molecule non-polar?
VSEPR THEORY The shapes of simple molecules and
ions can be determined by using the Valence Shell Electron Repulsion (VSEPR) theory.
Electron pairs around the central atom repel each other
Bonding pairs and lone pairs arrange themselves to be as far apart as possible
PRINCIPLES OF VSEPR THEORY
Find the number of electron pairs / charge centres in the valence shell of the central atom.
Electron pairs / charge centres repel each other to the positions of minimum energy in order to gain maximum stability.
Pairs forming a double or triple bond act as a single bond
Non-bonding pairs repel more than bonding pairs.
A FEW VSEPR SHAPES
PRACTICE
Methane (CH4) – tetrahedral Ammonia (NH3) – pyramidal Water (H2O) – bent Carbon Dioxide (CO2) - linear
METHANE, AMMONIA AND WATER
Valence Shell Electron-Pair Repulsion Theory (VSEPR)
Procedure
1. Sum the total Number of Valence Electrons
Drawing the Lewis Structure
2. The atom usually written first in the chemical formula is the Central atom in the Lewis structure
3. Complete the octet bonded to the Central atom. However, elements in the third row have empty d-orbitals which can be used for bonding.
4. If there are not enough electrons to give the central atom an octet try multiple bonds.
Predicting the Shape of the Molecule
5. Sum the Number of Electron Domains around the Central Atom in the Lewis Structure; Single = Double = Triple Bonds = Non-Bonding Lone Pair of Electrons = One Electron Domain
6. From the Total Number of Electron Domains, Predict the Geometry and Bond Angle(s); 2 (Linear = 180º); 3 (Trigonal Planar = 120º); 4 (Tetrahedral = 109.5º); 5 (Trigonal Bipyramidal = 120º and 90º); 6 (Octahedral = 90º)
7. Lone Pair Electron Domains exert a greater repulsive force than Bonding Domains. Electron Domains of Multiple Bonds exert a greater repulsive force than Single Bonds. Thus they tend to compress the bond angle.
LONE PAIRS AND BOND ANGLES
Lone pairs are held closer to the nucleus than the bonding pairs.
The distance between the lone pair electrons and the bonding pairs of electrons is shorter than the distance between the bonding pairs to each other.
Repulsion due to lone pairs causes the bond angles to become smaller
Order of repulsion : lone pair-lone pair > lone pair- bonding pair > bonding pair – bonding pair
Ammonia, NH3
Greater repulsion by lone pair of electrons. Bond angle is smaller than 109.50(1050)
Water, H2OEven greater repulsion bytwo lone pair of electrons. Bond angle is even smaller (1050)
Methane, CH4
Bond angle is 109.50
PREDICTING THE SHAPE OF IONS
ConsiderNH4
+ , H3O+ , NO2
-
NH4+
As the 4 negative charge centres repel each other and take up positions in space to be as far apart as possible, the electron pairs are distributed in tetrahedral arrangement.
H3O+
As the 4 negative charge centres repel each other and take up positions in space to be as far apart as possible, the electron pairs are distributed in tetrahedral arrangement. With one lone pair of electrons, the actual structure is trigonal pyramid with a bond angle of 1070 for H-O-H bond.
Read page 104
NO2-
As the 3 negative charge centres repel each other and take up positions in space to be as far apart as possible, the electron pairs are distributed in trigonal planar arrangement. With one lone pair of electrons, the actual structure of the ion is bent with a bond angle of about 1170 for O-N-O bond.
ConsiderN2H4 , C2H2
MOLECULES WITH MORE THAN ONE CENTRAL ATOM
INTERMOLECULAR FORCES Forces between molecules. Does not exist in giant structure (ionic
cpds, metals & giant covalent structure)
Intermolecular forces
Van der Waals forces
Hydrogen bond
http://chemtools.chem.soton.ac.uk/projects/emalaria/index.php?page=13
Intermolecular Forces: are generally much weaker than covalent or ionic bonds. Less energy is thus required to vaporize a liquid or melt a solid. Boiling points can be used to reflect the strengths of intermolecular forces (the higher the Bpt, the stronger the forces)
Hydrogen Bonding : the attractive force between hydrogen in a polar bond (particularly H-F, H-O, H-N bond) and an unshared electron pair on a nearby small electronegative atom or ion
Very polar bond in H-F.The other hydrogen halides don’t form hydrogen bonds, since H-X bond is less polar. As well as that, their lone pairs are at higher energy levels. That makes the lone pairs bigger, and so they don't carry such an intensely concentrated negative charge for the hydrogens to be attracted to.
Hydrogen Bonding & Water
VAN DER WAALS FORCES Electrons can at any moment
be unevenly spread producing a temporary instantaneous (fluctuating) dipole.
An instantaneous dipole can induce another dipole in a neighbouring particle resulting in a weak attraction between the two particles.
The forces of attraction between temporary or induced dipoles are known as Van der Waals’ forces (London Dispersion Forces).
Van der Waals’ forces increases with increasing mass.
London Dispersion Forces – significant only when molecules are close to each other
Prof. Fritz London
Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom
The ease with which an external electric field can induce a
dipole (alter the electron distribution) with a molecule is
referred to as the "polarizability" of that molecule
The greater the polarizability of a molecule the easier it is
to induce a momentary dipole and the stronger the
dispersion forces
Larger molecules tend to have greater polarizability
Their electrons are further away from the nucleus (any
asymmetric distribution produces a larger dipole due
to larger charge separation)
The number of electrons is greater (higher probability
of asymmetric distribution)
thus, dispersion forces tend to increase with increasing
molecular mass
Dispersion forces are also present between polar/non-
polar and polar/polar molecules (i.e. between all
molecules)
The strength of the intermolecular forces determines how easily the molecules will separate and hence the melting and boiling points.
Why is there an increasing boiling points of the noble gases as you go down the group?
The boiling points of the noble gases are
Helium -269°C Neon -246°C Argon -186°C Krypton -152°C Xenon -108°C Radon -62°C
Because electrons are always moving around very quickly, the charges switch around all the time.
the more electrons in a molecule / atom, the stronger these Van der Waals or London forces are.
This is seen in the increasing boiling points of the noble gases as you go down the group.
The boiling points of the noble gases are
Helium-269°C Neon -246°C Argon -186°C Krypton -152°C Xenon -108°C Radon -62°C
GROUP 7
Cl2 : gas
I2 : solid
Iodine molecule is made up of larger atoms with more electrons compared to chlorine.
With more electrons moving around, the temporary dipole will be larger.
The larger atoms in the molecule means that the valence electrons are less strongly held,
Hence the induced dipoles will be larger.
PERMANENT DIPOLES
van der Waals forces are present between covalent molecules with no H atom attached to N, O or F.
E.g. van der Waals forces are present between HCl molecules.
There’s also other intermolecular forces beween the molecules :
permanent dipole - permanent dipole
These intermolecular forces between polar molecules are stronger than between non-polar molecules. (all things being equal)
For polar substances with similar RMM, the higher the dipole moment, the stronger the dipole-dipole attractions and the higher the boiling points.
COMPARE MOLECULES WITH SAME RMM Propane (C3H8) and ethanal (CH3CHO) both with
RMM = 44 Ethanal has a higher bp.Ethanal - is a polar molecule - has stronger intermolecular forces (van der waals &
dipole-dipole interactions) between the molecules of ethanal than between the propane molecules.
Read further for a few exceptions
It is not true that polar molecules have stronger intermolecular forces and hence higher bp than non-polar molecules.
Non-polar molecules with higher RMM might have higher bp.
Boiling points increase for polar molecules of similar mass, but increasing dipole:
SubstanceMolecular Mass
(amu)Dipole
moment, u (D)Boiling Point
(°K)
Propane C3H8 44 0.1 231
Dimethyl ether CH3OCH3
46 1.3 248
Methyl chlorideCH3Cl 50 2.0 249
AcetaldehydeCH3CHO 44 2.7 294
AcetonitrileCH3CN 41 3.9 355
HYDROGEN BONDING A hydrogen bond is a weak type of force
that forms a special type of dipole-dipole attraction which occurs when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of another electronegative atom with a lone pair of electrons. These bonds are generally stronger than ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds.
Hydrogen bonding is present between covalent molecules with H atoms attached to O, N and F
2 FACTORS AFFECTING THE BOND Strength of H bondThe larger the electronegativity of H and
the other atom (N, O or F), the stronger the H bond.
Strength F > O > N Number of them that can be formed
between neighbouring molecules
HYDROGEN BONDINGAlthough the strength is such F > O > N, HF can only form 1 H
bond to 1 neighbour.
H2O can form 2, thus promoting more intermolecular interactions .
The collective strength of the H bonds in water is greater than the strength of the H bonds in HF because each O atom (with 2 lone pairs) in the water molecule can form 2 H bonds with 2 other water molecules, whereas each F atom in HF molecule can only form 1 H bond with another HF molecule.
Ammonia molecule has 3 N-H bonds. N is larger and < electronegative than F, has far weaker H bonds due to lower electron density on the N atom (only 1 lone pair) compared to O and F.
When the RMM is large, we expect the boiling point to be high because larger molecules have more space for electron distribution and more possibilities for instantaneous dipole moment.
However,Compound
RMM Boiling pt (K)
H2O 18 373
HF 20 292.5
NH3 17 239.8
H2S 34 212
HCl 36.4 187
PH3 34 185.2
Greater intermolecularforce because H2O, HF,NH3 all exhibit hydrogen bonding .
tend to have higher viscositythan those that do not have H bond.Substances which have multipleH bonds exhibit even higher Viscosity.
FACTORS AFFECTING H BONDING
ElectronegativityCannot occur without significant
electronegativity difference between H and the atom it is boded to.
E.g. Both PH3 and NH3
have trigonal pyramidal shape but only NH3 has H bonding.
Atomic sizeWhen the radii of the 2 atoms differ greatly,
their nuclei cannot achieve close proximity when they interact resulting in weak interaction.
Is there any hydrogen bonding between the molecules if CH3F?
H H
H C F H C F
H H
H is not joined directly to F in each molecule, hence no hydrogen bonding between the molecules.
Is there any hydrogen bonding between the molecules of ethanol?
Hydrogen bonding affects the boiling points of water, ammonia,
hydorgen fluoride and other molecules the solubility of simple covalent
molecules such as ammonia, methanol and ethanoic acid in water
the density of water and ice. the viscosity of liquids, e.g. the alcohols.
EFFECTS OF H BONDING ON PHYSICAL PROPERTIES
Group 4A hydrides
Groups 4, 5, 6A hydrides
Van der Waals forces are made of dipole-dipole and London dispersion forces
COMPARISONS OF THE PHYSICAL PROPERTIES OF
IONIC COMPOUNDS VS COVALENT SUBSTANCES
Recall:
In ionic compounds, the ions are held together by strong ionic bonds in a giant ionic lattice.
In simple covalent molecules, the attractive forces between the molecules are known as intermolecular forces or van der Waal’s forces, which is weaker than the ionic bonds.
In giant covalent molecules, the atoms are held together by strong covalent bonds in a giant covalent lattice.
Types of Covalent Substances
Covalent substances can be divided into 2 categories as shown in the table below:
Simple Covalent Molecules
Macromolecules
Examples: Hydrogen, Oxygen, Water, Carbon Dioxide and Methane
Examples: Diamond, Graphite and sand (Silicon Dioxide)
The atoms can be either same like silicon and carbon (graphite and diamond) or of 2 different elements such as silicon dioxide.
Allotropes are two (or more) crystalline forms of the same element, in which the atoms ( or molecules) are bonded differently.
GIANT COVALENT LATTICE
Diamond GraphiteGiant Molecular Structure
Giant covalent structure
Each carbon atom is covalently bonded to 4 other carbon atoms in a tetrahedral arrangement.
Rings of C atoms in layers
Each carbon atom is covalently bonded to 3 other carbon atoms in each layer.
The fourth electron is not used in chemical bonding.
This gives hexagonal rings of six atoms that join together to from flat layers that are held together by van der Waal’s forces.
C60 Fullerene Molecular structure – consists of individual
C60 molecules
with covalent bonds within the molecule. 60 C atoms are arranged in
hexagons and pentagons. Van der Waal’s forces between molecules.
Diamond Graphite C60 fullerene
Melting/Boiling point
1. Hard2. Very high melting point
and boiling point
Strong covalent bond in the giant structure between C atoms.
A lot of energy needed to break them when melts.
1. Soft and slippery2. High melting point and
boiling point
Layers can slip pass each other.
Covalent bond in layers must be broken when it is melted.
Van der Waal’s force is easier to break.
1. Soft and slippery2. Low melting point
and boiling point
Few covalent bonds hold the molecules together but only weak Vander Waals forces between molecules.
Only Van der Waal's interactions have to be broken for melting.
Properties
Diamond Graphite C60 fullerene
Electrical conductivity
1. Non-conductor of electricity
No free electrons to conduct electricity.
1. Good conductor of electricity
Free electron is not used in bonding , able to move within the layers and free to conduct electricity.
1. Non-conductor of electricity
No movement of electrons available from one molecule to the next.
Properties
Diamond Graphite C60 fullerene
Solubility
1. Insoluble in water
There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another.
1. Insoluble in water
There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another.
1. in water but soluble in methyl benzene
There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another in the molecules.
Uses 1. Gemstones2. As tips of grinding,
cutting and polishing tools.
1. In pencils2. As a dry lubricant3. Brushes for electric
motors
Properties
4 Metallic Bonds
Metals consist of positive ions surrounded by a 'sea of moving electrons'. The negative 'sea of electrons' attracts all the positive ions and cements everything together.
Metallic bonds are the results of the strong forces of attraction between the negative electrons and the positive ions.
Hence, metals have high melting points and high boiling points.
Physical Properties of Metals
Physical Properties
Explanation Diagram
High Density The close packing of atoms explains why most metals have a high density
Good Conductor of Electricity
Metals are good conductors of electricity in the solid and molten state. This is due to the sea of delocalized electrons.
Physical Properties
Explanation Diagram
Malleable and Ductile
Metals are malleable (can be bent or flattened) and ductile (can be drawn into wires) because the layers of metal ions can slide over each other when a force is applied.
High melting and boiling point
Metallic bonds are strong bonds. Except for: Mercury has a low melting point and is a liquid at room temperature. Similarly, sodium and potassium have low melting and boiling point.
Explanation
The valence electrons do not belong to any particular atom, hence, if sufficient force is applied to the metal, 1 layer of metals can slide over another without disrupting the metallic bonding.
The metallic bonding in metal is strong and flexible and so metals can be hammered into thin sheets (malleability) or drawn into lonng wires (ductility) without breaking.
If atoms of other elements are added by alloying, the layers of ions will not slide over each other so readily. The alloy is thus less malleable and ductile and consequently harder and stronger.
‘Like tends to dissolve like’. Polar substances tend to dissolve in polar solvents, such as water, whereas non-polar substances tend to dissolve in non-polar solvents, such as heptane or tetrachloromethane.
Organic molecules often contain a polar head and a non-polar carbon chain tail. As the non-polar carbon chain length increases in an homologous series the molecules become less soluble in water.
Ethanol is a good solvent for other substances as it contains both polar and non-polar ends.
SOLUBILITY
Water will mix with polar liquids such as ethanol. The oppositely charged ends of the different molecules attract one another forming hydrogen bonds.
Gases are generally slightly soluble in water.
A small number of gases are highly soluble because they react with water to release ions.
Example,SO2(g) + H2O (g) H+(aq) + HSO3
-(aq)
This solution is known as sulfurous acid , a major component of acid rain
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