Electronegativity and polarity

Polar and non-polar bonds:

1)  Non-Polar bonds:

• A covalent bond shares an electron pair:

• In a hydrogen molecule, the electrons are attracted by the nucleus from each of the hydrogen atoms.

As the 2 atoms are identical, they will have the same number of protons. This means that the electrons are being 'pulled' equally by both of the hydrogen atoms. We say that the H - H bond is non-polar The same is true for any diatomic molecule where the atoms are identical:

2)  Polar bonds:

• If a covalent bond is between 2 different atoms then the attraction from each is more likely to be unequal.

• One atom will have more protons in the nucleus / less shielding.

• This means that that atom will attract the bonding pair of electrons more than the other.

• The power of an atom to attract bonding pair electrons to itself is called electronegativity/

Hydrogen chloride:

• The attraction from the nucleus for the bonding electrons will be different.
• Chlorine is the more electronegative atom.
• This means the bonding electrons will not be closer to the chlorine atom.
• This covalent bond is Polarised.
• Because chlorine atom has the bonding electrons nearer to it, the chlorine atom will have a small negative charge, d-.

                        d  is used to mean ' a little bit of '

• Because the hydrogen doesn't have its fair share of the bonding electrons it will have a small positive charge, d+.
• The molecule has a small negative charge at one end and a small positive charge at the other end.
• We say the molecule has a permanent dipole:

                    Di   -   ' 2 '

                    Pole    -    poles (positive and negative)

Polar molecules:

• Molecules like HCl are called polar molecules.
• This is because over the whole molecule there are ' 2 poles '.

• Some molecules can have polar covalent bonds without being polar.
• It comes down to symmetry:

• Basically there is not a positive end and a negative end.
• So CCl₄ in classed as non - polar even though all of the bonds are polarised.

How is electronegativity measured?

• Linus Pauling came up with the Pauling scale in 1932.

• Basically as you go towards the top right hand side of the Periodic table, the elements become more electronegative:

Pauling definition of electronegativity:-

                 The electronegativity of an atom represents the power of an atom in a molecule to attract electrons to itself.

• Cl, N, O and F are the most electronegative elements.
• Reactive metals, Na - K are the least electronegative elements
• The greater the difference in electronegativities, the greater the permanent dipole and the bigger the d-.

Electronegativity and bonding type   

1)    Covalent

• Elements of very similar electronegativities have their bonding electrons shared equally between the 2 atoms. 

• This is clearly a covalent bond.

2)    Polar covalent

• Elements with a slight difference in electronegativity will still share their bonding electrons.
• The electrons are not evenly shared:

• The covalent bond will however be polar.

3)    Ionic

• Elements with very different electronegativities however have a different type of bonding
• The more electronegative element will attract the bonding electrons to itself so much that that element takes both of the bonding electrons.

• The more electronegative element has now gained an extra electron to become a 1- ion.
• The lesser electronegative element has now lost an electron to become a 1+ ion.

Covalent to ionic:

• As a bond becomes more polar, there is a movement from covalent to ionic bonding.

Questions 1-2 P61

Intermolecular forces

Strengths of bonds and forces:

Ionic - When ionic compounds melt/boil, the forces of attraction are overcome and the ions separate.

NaCl(s)

Na+(g)

+

Cl-(g)

Molecular - When molecular compounds melt/boil, the covalent bonds remain intact.   

H2O(l)

H2O(g)

• Since the molecules do not break up there must be some forces of attraction between the molecules, which are broken.

• These are called intermolecular forces.

• Since molecular substances can exist as solids, liquids and gases by varying the temperature and pressure, intermolecular forces must always exist although they may be very weak.

• There are 3 types of intermolecular forces of attraction:

    1)    Van der Waals' forces

    2)    Permanent dipole - dipole forces

    3)    Hydrogen bonding

Permanent dipole - dipole interactions:

• When we have 2 atoms in a covalent bond with different electronegativities, the bond is polarised.

• If the molecule has a d+ end and a d- end, the molecule is said to have a permanent dipole:

• The d+ end of one molecule will be attracted to the d- end of a neighboring molecule.

• This attraction is called a permanent dipole - dipole force of attraction:

Van der Waals' forces (induced dipole - dipole interactions)

• Helium does not form ionic or covalent bonds but it is possible to condense it to a liquid then to a solid.

• Energy is released when a change of state occurs.  0.105KjMol^-1

• This very weak force of attraction is known as Van der waals forces.

• It is due to the continually changing electric charge interactions between atoms, called induced dipole - dipole forces of attraction.

What causes Van der Waals' forces:

• These are present in all molecules but are the only forces of attraction present in non - polar molecules

• The electrons in shells are continually moving.

• In the turmoil we get an uneven distribution of electrons / charge.

• At any moment or snap shot in time there would be an instantaneous dipole across the whole atom / molecule.

• The negative end of the dipole induces a dipole of opposite charge in neighbouring atoms

• A force of attraction results.

• These induced dipole – dipole interactions produces a cohesive force.

Imagine an atom to be like a large spherical jelly with a golf ball at the center. The golf ball is the nucleus, the jelly is the cloud of electrons whizzing about this. The net average field will be zero because the (+)ve nucleus field will be exactly balanced by the electron cloud. Atoms vibrate, at any instant the cloud is likely to be slightly off center.  This creates an instantaneous dipole. If we have another atom next to it, this atom will be affected by the instantaneous dipole. This will induce a dipole in the neighbouring atom. The 2 dipoles attract one another – producing an attractive interaction.

• The forces of attraction are so weak that we use dd to represent extra small charges:

• Molecules with very similar electronegativities will also only have Van der Waals' forces of attraction, eg:

            Hydrocarbons        Diatomic elements

• The greater the number of electrons (and protons) in an atom / molecule, the greater these fluctuations are and the greater the fluctuations

• This will give greater VDW attraction.

• If you consider the alkanes – The boiling point increases as the Mr increases.

• This is because of the increased number of electrons, which increase the VDW attraction.

Questions 1-2 P63

Hydrogen Bonding:

Water is peculiar

• Hydrogen bonding explains these observations.

• It is the strongest of the intermolecular forces.

Why are hydrogen bonds so strong?

• The atoms O,N,F are so strongly electronegative that the bonding pair of electrons are so far from the H that they are almost able to be donated.

• This along with the small size of the H atom means that the H in the molecule is very positive.

• The pair of bonding electrons are very near the O,N,F.  With their small sizes they are very negative. 

• With its own lone pair(s) of electrons a strong force of attraction is able to occur between the H and the lone pair of electrons on neighbouring molecules.

• This force of attraction is known as a Hydrogen bond and is represented by a dotted line:

Maximum bond strength is when the bond angle O-H-O is 180o. The strength of a hydrogen bond is typically ~30Kj. Compare this with the strength of a covalent bond ~300Kj.  A Hydrogen bond is ~1/10th a covalent bond. Similarly with VDW attraction ~3Kj.

1)    Within a molecule, the hydrogen must be highly polarised (very positive)

2)    Within a molecule, the atom joined to the hydrogen must be very electronegative, O,N,F.

3)    Within a molecule, the atom joined to the hydrogen must also have a lone pair of electrons.

• H

• O,N,F

• O,N,F must have a lone pair

• Hydrogen bonding is strong enough to change physical properties but not chemical properties.

• Water would be a gas at room temperature and pressure if it was unable to hydrogen bond.

Ice is less dense than water:

• The 2 hydrogen bonds per water molecule of water sets  up a 3D structure.

• The Hydrogen bonds are longer than the O - H bonds. 

• This means they are held further apart than in water making ice less dense.

• Hydrogen bonds give water a skin effect and this contributes to the high surface tension.

• Other liquids with hydrogen bonding also have a surface tension.

Hydrogen bonding in biological molecules:

• Hydrogen bonds exist in biological molecules due to NH₂ and OH groups.

• Hydrogen bonding is responsible for the shapes of proteins and the double helix in DNA:

Questions P65  1-2 / P73  6,8P74  2,5

Metallic bonding and structure

• Positive metal ions are in a fixed position while the outer shell electrons are delocalised between all the atoms in the metallic structure:

The model consists of metal ions surrounded by ‘Mobile sea of electrons’. Attraction occurs between the ions and the delocalised electrons The sea of electrons explains its electrical conductivity and its thermal conductivity. The sea of electrons bonds the metal ions tightly into the lattice. This explains its high melting point. Since strong forces of attraction exists even in the liquid phase, metals tend to have a wide temperature range over which they remain liquid. The giant metallic lattice is often referred to as

Properties of giant metallic lattices:

1)  High melting and boiling points -

• Attraction occurs between the fixed ions and the delocalised electrons.

• The attraction between the positive ions and the 'sea of electrons' is strong

• The sea of electrons bonds the metal ions tightly into the lattice.

• This explains its high melting point and boiling points.

2)  Good electrical conductors -

Mobile electrons will be attracted to a positive terminal. Electrons are replaced from the negative terminal.

3)  Malleability and ductile -

• Malleable - can be hammered or pressed into shape.

• Ductile - can be drawn / stretched into wire.

• Due to the delocalised electrons, the metallic structure has a degree of 'give' which allow layers to slide past each other.

Alloys

• These are mixtures of metals.  This do not however form compounds.

• Compounds have a definite ratio of the elements.

• Metals can be mixed in different proportions.

• In the structure one metal ion is replaced with another.

• Often one ion is bigger than the original metal ion. 

• This acts as a barrier preventing the layers from sliding past each other.

• This makes the alloy harder than the original metal.

Q1-3  P 67

Structure of ionic compounds

Giant ionic lattices

Each sodium ion is surrounded by 6 chloride ions. Each chloride ion is surrounded by 6 sodium ions. This continues in all directions and is describes as a Giant Ionic Lattiice

Properties of ionic compounds

1)  High melting and boiling points:

• There are strong electrostatic forces of attraction between the ions.

• This means that they are not easily broken which is why they have high melting and boiling points.

• The higher the charges between the ions, the stronger the electrostatic forces of attraction.

• The stronger the forces of attraction, the more heat energy is required to overcome those forces and hence melt / boil.

2)  Electrical conductivity:

• In a giant ionic lattice, the ions are held in a fixed position.  This means that the ions cannot move.

• This is why they do not conduct electricity as a solid.

• When the ions are molten or dissolved - the ions are now free to move.

• This means they will conduct electricity:

Solubility:

• When an ionic solid dissolves, the ions in the lattice separate.

• The energy required to separate ions within a lattice is large. ie high melting points.

• The energy required on dissolving must be equal and opposite to the energy required to separate ions (as dissolving separates the ions).

Where does the energy come from?

• Where does this large amount of energy come from if all we are doing is dissolving the solid in water?

• There must be some process during dissolving that can release enough energy to separate the ions from the lattice.

• If energy is being released there must be some type of attractive interaction to release energy:-

• The force of attraction comes between the d+ /  d- end of the polar water molecule and the opposite charges on the ion.

• This attraction releases energy (as all ‘bond forming’ reaction do).

• Many water molecules surround the anion as shown.

• This releases energy which is used to break up the lattice structure:

Qu 1-3  P69

Structures of covalent compounds

• Covalent compounds fall into 1 of 2 categories:

A)    Simple molecular lattice

B)    Giant covalent lattice

A)    Simple molecular structures:

• These are made up from simple (small) molecules such as:  CO₂, N₂, O₂, I₂ and H₂O.

• In its solid forms, the molecules are held together by weak intermolecular forces (VDW / Dipole / H Bonding).

• The atoms within the molecules are made up from strong covalent bonds.

Properties of simple molecular structures:

1)    Low melting and boiling points:

• Simple molecular molecules have low melting and boiling points due to weak forces of attraction between the molecules.

• When you melt or boil these molecules, you do not break up the molecule, only overcome the forces of attraction between them.

• From the animation - can you explain why water is a liquid at room temperature while carbon dioxide is a gas?

2)    Electrical conductivity:

• Simple molecules have no free moving electrons or charges.

• This means they do not conduct electricity.

3)    Solubility:

• Simple molecules are only soluble in non - polar solvents like hexane.

• This is because both molecule and solvent have weak forces of attraction between their molecules.

B)    Giant covalent structures

• These are structures that have extensive covalent bonded atoms in a giant lattice structure.

• Diamond, graphite and quartz are examples of these:

Diamond	Graphite

Properties of giant covalent structures

1)    High melting and boiling points:

• Due to the extensive covalent lattice structure of these types of compounds, a lot of energy is required to break these covalent bonds.

• This gives them very high melting and boiling points.

• Diamond is the hardest natural substance due to its covalent lattice structure.

2)    Electrical conductivity:

• As there are no free moving electrons or charges, they are non - conductors of electricity.

• Graphite however is the exception to the rule.

• Each carbon in graphite has 3 covalent bonds.

• The 4th electron from each carbon atom is delocalised and free moving between the layers.

• This means it is able to conduct electricity.

3)    Solubility:

• Giant covalent structures are insoluble in both polar and non - polar solvents.

• The covalent bonds are too strong to be over come by any solvent.

4)    Hardness:

• Due to the extensive covalent lattice structure, all giant covalent structures are hard.

• Graphite however is the exception to this.

• Due to its layered structure and that the delocalised electrons are between the layers.

• This allows the layers to slide over and past each other.

Questions 1-2 P71