Intermolecular Forces Flashcards
Electronegativity
The ability of an atom to attract a BONDING PAIR of electrons in a covalent bond towards ITSELF.
Factor affecting electronegativity 1
SIZE OF THE POSITIVE CHARGE ON THE NUCLEUS:
- moving left to right across a period, the number of PROTONS in the nucleus INCREASES.
- This increased positive charge, increases the ATTRACTION between the NUCLEUS and the pair of electrons in the covalent bond.
- Therefore, elements on the right of the periodic table are more electronegative than those on the left.
( The two electrons in the first shell of an atom CANCEL OUT the effect of TWO PROTONS in each element)
Factor affecting electronegativity - Atomic radius
ATOMIC RADIUS:
- The SMALLER the atomic radius , the CLOSER the bonding electrons will be to the nucleus of an atom.
- As we move from left to right across a period, the atomic radius DECREASES.
- Therefore elements on the right of the periodic table are more electronegative than those on the left.
Factor affecting electronegativity - number of inner shells
THE NUMBER OF INNER SHELLS:
- Electrons in the inner shell SHIELD electrons in the outer shell from the positive charge of the nucleus.
- Therefore , the greater number of inner shells, the LOWER the electronegativity.
Covalent bonds between two IDENTICAL atoms
- The NUCLEUS of each atom attracts the bonding pair of electrons to the EXACT SAME extent.
- The bonding pair will be SHARED EQUALLY between the two atoms and will be found on average HALF WAY between the two nuclei.
- The bond between two identical atoms is NON-POLAR and 100% covalent.
Covalent bonds between two DIFFERENT ATOMS
- The more ELECTRONEGATIVE atom will cause its nucleus to attract the bonding pair of electrons MORE STRONGLY than the other atoms nucleus.
- The bonding pair will NOT BE SHARED EQUALLY between the atoms and will on average found CLOSER to the more electronegative atom.
- The bond between two different atoms is said to be POLAR
Representing Polar covalent bonds
- Write DELTA + and a DELTA - to show the charges.
- Delta means slight and this is because the electron pair has only SHIFTED towards the more electronegative atom.
- The DELTA - always goes on the MORE ELECTRONEGATIVE atom.
- Alternatively , we use a CROSSED ARROW pointing towards the MORE ELECTRONEGATIVE element.
Polar covalent bonds
A separation of OPPOSITE CHARGES.
The separation of opposite charges ACROSS A BOND is called a DIPOLE.
If the difference in charge across the polar covalent bond does not change, it is called a PERMANENT DIPOLE.
Polar molecule
When a PERMANENT DIPOLE exists over an UNSYMMETRICAL MOLECULE.
Non-polar molecule
- The molecule is SYMMETRICAL.
- The PERMANENT DIPOLES in the bonds CANCEL OUT with each other
- This is because they OPPOSE each other EQUALLY due to the SYMMETRY of the molecule.
The polar nature of WATER
- Both O-H bonds have a PERMANENT DIPOLE.
- The molecule is UNSYMMETRICAL with a bond angle of 104.5 degrees and NON- LINEAR shape.
- The dipoles act in different directions and do NOT oppose each other equally and cancel each other out.
- Overall the OXYGEN end has a DELTA - charge and the HYDROGEN end of the molecule has a DELTA + charge.
The non-polar molecule CARBON DIOXIDE
- Both C=O bonds have a PERMANENT DIPOLE.
- The molecule is SYMMETRICAL with a bond angle of 180 degrees and a LINEAR shape.
- The dipoles act in opposite directions and exactly oppose each other.
- The overall dipole over the molecule is ZERO because the dipoles CANCEL OUT.
- Therefore, a carbon dioxide molecule has NO OVERALL POLARITY.
The Pauling scale of Electronegativity
The scale compares the electronegativity of atoms of different elements.
NON-METALS:
- nitrogen
- oxygen
- fluorine
- chlorine
Are the MOST electronegative atoms.
GROUP 1 METALS:
The LEAST electronegative atoms.
Type of bonding and electronegativity
The DIFFERENCE in electronegativity between two atoms indicates the type of BONDING.
PURE COVALENT:
Non-polar bond due to identical atoms bonded together with NO difference in electronegativity.
Electronegativity difference is 0
POLAR COVALENT:
SMALL difference in electronegativity due to NON-IDENTICAL atoms bonded together.
Therefore, PARTIAL charges exist on atoms.
Electronegativity is between 0 and 1.80.
IONIC:
LARGE differences in electronegativity and so ions with FULL charges form.
Electronegativity difference is ABOVE 1.80
INTERmolecular forces
- WEAK interactions BETWEEN the dipoles of different molecules.
- Occur because molecules are attracted to each other.
- Determine the PHYSICAL properties of molecules.
INTRAmolecular forces
- STRONG interactions that occur WITHIN molecules.
- Refer to either COVALENT or IONIC bonds
- Determine the CHEMICAL REACTIVITY of molecules.
Simple molecular substances
- Consist of relatively SMALL molecules and each molecule has a FIXED number of atoms.
LOW boiling point ~ takes very little energy to break intermolecular forces.
- HEATING a simple molecular substance causes the molecules to move faster.
- At a certain temperature, the intermolecular forces break and this allows the molecules to move away from each other.
Induced dipole-dipole interactions
Also referred to as LONDON or DISPERSION forces.
ATOM 1 ~
- electrons are randomly moving from place to place and are not evenly spread out.
- For a fraction of a second , more electrons are on the RIGHT hand side of the atom.
- Forms an INSTANTANEOUS DIPOLE ~ right hand side has a delta negative charge and the left hand side has a delta positive charge.
ATOM 2 ~
- Because the right hand side of atom 1 has a delta negative charge, it will REPEL the electrons in atom 2.
- This causes the electrons in atom 2 to move towards the right hand side.
- Forms an INDUCED DIPOLE as it was caused by the dipole in atom 1.
Induced dipoles form in further neighbouring molecules.
ALL of these dipoles now experience a force of attraction known as LONDON forces.
- Typically 0.001 the strength of a covalent bond.
Properties of LONDON forces
- Weakest IMF and easily broken
- Caused by RANDOM electron movement.
- Every single atom or molecule will experience London forces even if they experience other IMF as well.
- The STRENGTH depends on the number of ELECTRONS.
This means atoms with a GREATER number of electrons have a HIGHER BP as they experience strong London forces.
PERMANENT dipole-dipole interactions
When two molecules that have a PERMANENT DIPOLE get near enough then their permanent dipoles can lead to an ATTRACTION.
H-F —— H-F ——- H-F—— H-F—–H-F ——–
- POLAR molecules like H-F have permanent dipoles.
- The delta negative charge on the fluorine atom is attracted to the delta positive charge on the hydrogen atom on the neighbouring molecule.
- This results in a WEAK permanent dipole-dipole force.
- This force is typically 0.01 the strength of a covalent bond.
The TWO types of intermolecular forces
VAN DER WAALS FORCES:
Dipole-dipole interactions & induced dipole-dipole interactions.
HYDROGEN BONDING
PERMANENT dipole-INDUCED forces
When a PERMANENT DIPOLE in a molecule induces a dipole in another molecule.
EXAMPLE:
H-F —– F-F
- The delta negative charge on the F atom of H-F repels the electrons in F2 , moving the electrons in F2 to the right side of the atom.
- This makes the right side delta negative and the left side delta positive.
- The permanent dipole in HF has induced a dipole in F2 and results in a weak permanent dipole-induced force.
- Typically 0.01 the strength of a covalent bond.
Hydrogen bonding
- The strongest type of intermolecular force
- Around one tenth the strength of a covalent bond
- They are a specialised type of permanent dipole-dipole force.
TWO conditions for hydrogen bonds to form
- A hydrogen atom must be bonded to a STRONGLY ELECTRONEGATIVE element.
- The electronegative atom must have at least one LONE PAIR of electrons.
This means the only elements that can from hydrogen bonds are:
- NITROGEN
- OXYGEN
- FLUORINE
The formation of a hydrogen bond in HYDROGEN FLUORIDE
- The hydrogen atom is highly electron deficient d+, due to being covalently bonded to the highly electronegative fluorine atom d-
- The Hd+ is attracted to the LONE PAIR of electrons on the fluorine atom in a neighbouring molecule.
- This attraction is known as a hydrogen bond and is indicated using a HASHED LINE.
- The hydrogen atom is found along the straight line drawn between the two fluorine atoms - LINEAR SHAPE.
Properties of water : DENSITY
- In its LIQUID form, the water molecules are moving RANDOMLY , sometimes close together, other times further apart.
- This means as a liquid , hydrogen bonds are constantly being FORMED and BROKEN.
- In its solid form, ICE, the water molecules arrange themselves into an ordered structure , stabilised by hydrogen bonds.
- This means the water molecules are further apart.
- Therefore, the solid form of water, is LESS DENSE than the liquid form.
- This means that ice can float on the surface of water.
Properties of water : BOILING POINT
- HIGHER THAN EXPECTED temperatures are required .
- This is because EXTRA ENERGY is required to overcome the strength of the hydrogen bonds , IN ADDITION to the permanent dipole-dipole and induced dipole-dipole forces.
- This means water has a RELATIVELY high melting and boiling point.
Properties of water : SURFACE TENSION
- The presence of hydrogen bonding gives water a HIGHER THAN EXPECTED surface tension.
- This means water ha a RELATIVELY high surface tension.
Properties of water : VISCOSITY
- Liquid water flows LESS easy than expected.
- This means water has a RELATIVELY HIGH viscosity.
The structure of ice
- Forms an OPEN TETRAHEDRAL LATTICE structure of water molecules.
- Each water molecule forms FOUR HYDRGEN BONDS due to having TWO lone pairs and TWO hydrogen atoms.
How to figure out if a molecule is POLAR or NON-POLAR
Identify the CHARGES on each element:
- If the elements are different , there will be a difference in electronegativity.
- The more electronegative element (further right) will have a delta negative charge, whilst the other element has a delta positive charge.
- The bonds have a PERMANENT DIPOLE
Is the molecule SYMMETRICAL or NOT :
- Are the charges evenly distributed
- Draw the molecule and identify the shape and bond angle.
POLAR:
- Unsymmetrical
- The dipoles act in different directions and DO NOT oppose each other equally and cancel out.
- The molecule has an OVERALL DIPOLE.
NON-POLAR:
-Symmetrical
- The permanent dipoles in the bonds oppose each other equally and CANCEL OUT.
