Topic 2 - Bonding and Structure Flashcards
Ionic bond
the strong electrostatic attraction between oppositely charged ions
Effect of ionic radius on the strength of ionic bonding
larger ions with larger ionic radii will have a weaker attraction to the oppositely charged ion because attractive forces have to act over a greater distance
Effect of ionic charge on the strength of ionic bonding
Ions with a greater charge will have a greater attraction to the other ions, resulting in stronger forces of attraction and therefore stronger ionic bonding
Trend in ionic radii down a group
As you go down each group, the ions have more electron shells. Therefore, ionic radius increases.
Trend in ionic radii across a period
Ionic radius decreases as number of protons increases. Positive charge of the nucleus increases, so electrons are pulled closer to the nucleus.
Evidence for the existence of ions
- conduct electricity
- during electrolysis, positive ions in solution are attracted to cathode and
negative ions in solution are attracted to the anode - physical properties: high MP, soluble in water but not in non-polar solvents
Relationship between bond length and bond strength in covalent bonds
inversely proportional
As bond length decreases, the strength of the covalent bond increases, as electrons involved are more tightly held when the distance between the nuclei of the bonded atoms is smaller
Covalent bond
the strong electrostatic attraction between two nuclei and the shared pair of electrons between them
Order of repulsion in covalent bonds
lone pair – lone pair > lone pair – bond pair > bond pair – bond pair
Linear structure bond angle
180˚
Trigonal planar structure bond angle
120˚
Bent structure bond angle
104.5˚
Pyramidal structure bond angle
107˚
Tetrahedral structure bond angle
109.5˚
Trigonal bipyramidal structure bond angles
120˚
90˚
Octahedral structure bond angle
90˚
Electronegativity
the ability of an atom to attract the bonding electrons in a covalent bond
Electronegativity difference needed for ionic bonds
𝜟x > 1.7
Electronegativity difference needed for polar covalent bonds
1.7 ≥ 𝜟x ≥ 0.5
Electronegativity needed for pure covalent bonds
𝜟x < 0.5
London force (instantaneous dipole – induced dipole)
a temporary attractive force due to the formation of temporary dipoles in a non-polar molecule
the constant “sloshing around” of the electrons in the molecule causes rapidly fluctuating dipoles
Permanent dipoles
weak intermolecular forces of attraction that arise between permanently polar molecules
Occur when two atoms in a molecule have substantially different electronegativity: one atom attracts electrons more than the other, becoming more negative, while the other atom becomes more positive
Hydrogen bond
a special type of permanent dipole-dipole force that forms when hydrogen forms a covalent bond with a very electronegative element: either nitrogen, oxygen or fluorine
Why water has a high melting/boiling temperature
Hydrogen bonds are relatively strong. These extra forces in addition to London forces require more energy to be overcome
Why ice is less dense than water
Open lattice structure in ice means rigid hydrogen bonds hold the water molecules apart (hydrogen bonds in ice are longer)
When ice melts, hydrogen bonds collapse, allowing water molecules to come closer together
Trend in boiling temperature of alkanes with increasing chain length
Boiling point increases as there are more points of contact with each adjacent molecule, causing stronger London forces between adjacent molecules
Branching on boiling temperatures in alkanes
Branching makes molecules more compact, reducing surface area. Branched alkanes will have lower boiling points than straight chain alkanes.
Why alcohols have a low volatility and higher BP compared to alkanes with a similar number of electrons
They possess hydrogen bonds in addition to London forces and dipole-dipole interactions, whereas alkanes only have London forces
The trends in boiling temperatures of the hydrogen halides, HF to HI
HF > HI > HBr > HCl
HF can form hydrogen bonds while the others can’t
Increase from HCl to HI is caused by increasing London forces due to increasing number of electrons
Water dissolving ionic compounds
water can break down or disrupt the ionic lattice and surround each ion in solution
the greater the ionic charge, the less soluble an ionic compound is
Water dissolving simple alcohols
Alcohols have a hydroxyl group that can form hydrogen bonds with water
alcohols with longer carbon chains show decreased solubility due to the hydrocarbon chain’s non-polar character overruling the polar hydroxyl group
Water as a poor solvent for halogenoalkanes
The dipole moment of halogenoalkanes is too weak to form hydrogen bonds with water molecules. The hydrogen bonds between water molecules are stronger than the dipole interactions that can form between water molecules and the halogenoalkane so the compound does not dissolve
Like dissolves like
Compounds which have similar intermolecular forces to those in the solvent will generally dissolve
e.g., non-polar solvents will dissolve non-polar solutes as London forces can form between them
Metallic bonding
the strong electrostatic attraction between metal ions and the delocalised electrons
In what are giant lattices present?
- ionic solids (giant ionic lattices)
- covalently bonded solids, such as diamond, graphite and silicon(IV) oxide (giant covalent lattices)
- solid metals (giant metallic lattices)
Diamond structure
- a giant lattice of carbon atoms with strong bonds in all directions
- each carbon is covalently bonded to four others in a tetrahedral arrangement with a bond angle of 109.5˚
- the hardest substance known
for this reason it is used in drills and glass-cutting tools
Graphite structure
- 3 bonds
- trigonal planar
- sheets of hexagons -> slippery, single electrons are delocalised and conduct electricity
- strong/lightweight
- insoluble
Graphene structure
- a single layer of carbon atoms that are bonded together in a repeating pattern of hexagons
- 3 covalent bonds per atom and the 4th outer electron per atom is delocalised
Dative covalent bond
both electrons of the covalent bond come from one atom (arrow instead of dash) e.g., NH4+
