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Shapes of molecules and intermolecular forces (Intermolecular forces…
Shapes of molecules and intermolecular forces
Shapes of molecules and ions
Electron pair repulsion theory
• The electron pairs surrounding a central atom determine the shape of the molecule or ion
• The electron pairs repel one another so that they are arranged as far apart as possible
• The arrangement of the electron pairs minimises the repulsion and thus holds the bonded atoms in a definite shape
Shapes of molecules
•
Tetrahedral (109.5 bond angle)
- 4 bonded pairs eg. CH4
• Lone-lone pair > bonded-lone pair > bonded-bonded pair (in terms of strength)
•
Pyramidal (107)
- 3 bonded pairs one lone pair eg. ammonia
•
Bent or non-linear (104.5)
- 2 bonded pairs 2 lone pairs eg. water
• Lone pairs repel more strongly than bonded pairs, so they repel the bonded pairs slightly closer together, which decreases the bond angle (by about 2.5 degrees for each lone pair)
•
Trigonal planar (120)
- 2 bonded pairs eg. boron trifluoride
•
Octahedral (90)
- 6 bonded pairs eg. sulfur hexafluoride
Electronegativity and polarity
Electronegativity
• When bonded atoms are different elements the nuclear charges are different, the atoms may be different sizes and the shared pair of electrons may be closer to one atom than another
• The shared pair of electrons in the covalent bond may experience a stronger attraction from one of the nuclei
• The attraction of a bonded atom for the pair of electrons in a covalent bond is
electronegativity
• Electronegativity is measured using the Pauling scale – electronegativity increases as you go up and to the right of the group (fluorine has the highest electronegativity)
• Covalent if electronegativity is 0
• Polar is electronegativity is 0 – 1.8
• Ionic if electronegativity > 1.8
Bond polarity
• In a non-polar bond, electron pair is shared equally – this is a
pure covalent bond
• An example is C-H, as they have very similar electronegativities so form non-polar bonds
• In a polar bond, the bonded electron is shared unequally between the bonded atoms – results in a
polar covalent bond
• Example H-Cl. The H is slightly positive (δ+), and the Cl is slightly negative (δ-)
• The separation of opposite charges is called a
dipole
• A dipole in a covalent bond doesn’t change and is called a
permanent dipole
• A molecule with polar bonds can be non-polar if the dipoles act in opposite directions and so exactly oppose one another
• Polar molecules can be by polar solvents – the water molecules attract the anions and cations which causes the ionic lattice to break down. The resulting solution is surrounded by water ions, where the cations are attracted to the oxygen and the anions to the hydrogen ions
Intermolecular forces
Forces between molecules
•
Intermolecular forces are weak interactions between dipoles of different molecules
• 3 different categories -
induced dipole-dipole interactions (London forces), permanent dipole-dipole interactions, hydrogen bonding
• London forces are the weakest, hydrogen bonds are the strongest, but single covalent bonds are much stronger than all of them
Induced dipole-dipole interactions
• London forces are weak intermolecular forces that exist between
all
molecules, polar or non-polar
• Act between induced diploes in different molecules
• Movement of electrons produces a changing dipole in a molecule
• At any instant, an instantaneous dipole will exist, but its position is constantly shifting
• The instantaneous dipole induces a dipole on a neighbouring molecule
• The induced dipole induces further dipoles on neighbouring molecules, which then attract one another
• Induced dipoles are only temporary
• The more electrons in each molecule the larger the instantaneous and induced dipoles so the greater the induced dipole-dipole interactions so the stronger the London forces
Permanent dipole-dipole interactions are bonds between polar molecules. They are stronger than London forces as extra energy needed to break the additional permanent dipole-dipole interactions.
Simple molecular substances
•
A simple molecular substance is made up of simple molecules such as H2
• In solid state, simple molecules form a regular structure – a
simple molecular lattice
• Molecules held in place by weak molecular forces while atoms within each molecule bonded together strongly by covalent bonds
• All simple molecular substances have very low boiling points as the weak intermolecular forces are easily broken by low temps (they don’t need much energy to be overcome)
• However only the weak intermolecular forces are broken – covalent bonds stay
• Don’t conduct as no mobile charged particles
Solubility
• Non-polar simple molecular substances are soluble as when added to a non-polar solvent intermolecular forces form between the substance and the solvent
• The interactions weaken the intermolecular forces in the simple lattice which causes the intermolecular forces to break and the compound dissolves
• When simple molecular substance added to a polar solvent, little happens as the intermolecular bonding within the polar solvent is too strong to be broken
• Simple molecular substances tend to be insoluble in polar solvents
Hydrogen bonding
Hydrogen bonds
• A hydrogen bond is a special type of permanent dipole-dipole interaction
• Happens between molecules containing an electronegative atom with a lone pair of electrons and a hydrogen atom attached to an electronegative atom – eg. H-F, H-O, H-N
• Strongest type of intermolecular attractions
• Shown by a dashed line (see textbook for diagram) – don’t forget to label dipoles
Anomalous properties of water
• Ice (solid) is less dense than the liquid (water)
• The hydrogen bonds hold water molecular apart in an open lattice structure
• The water molecules are further apart in ice than in water, as the open lattice structure has holes in, which are due to water molecules being able to form 4 hydrogen bonds
• When ice melts, the lattice collapses and the molecules move closer together, causing water to be denser than ice
• Water has a relatively high melting and boiling point
• Hydrogen bonds are extra forces which need large amounts of energy to be broken
• When water boils, hydrogen bonds break completely
• Water also has high surface tension and viscosity
• Hydrogen bonds play key part in structure of DNA – between two bases