Topic 2 + 13A - Structure and bonding Flashcards
Define an ionic bond
Electrostatic attraction between two oppositely charged ions in an ionic compound. Formed by the transfer of electrons from one atom to another.
What are the properties of an ionic compound?
- Giant lattice structure
- Conduct electricity when molten or in solution because ions are free to move and carry charge
- high melting points because of strong electrostatic forces between oppositely charge ions
How does the charge of an ion affect melting point in ionic compound?
The larger the charge, the higher the charge density.
Therefore stronger forces of attraction, so more energy is required to overcome, so increased melting point
How does the size of an ion affect melting point in ionic compounds?
The larger the ion, the lower the charge density.
Therefore weaker forces of attraction, less energy required to overcome forces, lower melting point.
Explain the trend in ionic radius down a group
- increases down a group
- number of shells increases so valance electron further from nucleus
- increased shielding
- higher nuclear charge but this is outweighed by increased distance and shielding so ion gets larger
Define a covalent bond
The linkage of two atoms held together by the electrostatic attraction between positive nuclei and the negative charge on the shared pair of electrons.
What are the properties of diamond?
- giant covalent structure
- very high mpt due to strong COVALENT bonds
- does not conduct electricity as no free ions or delocalised electrons
- rigid structure allows good heat conduction
What are the features of graphite?
- layers slide over each other use to weak IMF between layers
- high mpt due to strong covalent bonds
- conducts electricity due to delocalised electrons
What is a dative covalent bond?
Where both the shared electrons come from the same atom.
Using the example of ammonia, NH3, how would you answer the Q: “state and explain the shape of a molecule”
- State the number of electron pairs.
e.g. N has 3 bond pairs and 1 lone pair so 4 electron pairs. - State the shape this is based on and why
e.g The electron pairs want to maximise separation to minimise repulsion so this shape is based on tetrahedral (4 electron pairs) - Explain lone pair repulsion
e.g Lone pairs repel more than bond pairs which decreases the angle between the bonds. - State name of shape and bond angle
e.g. the bond angle is reduced from 109.5 to 107 and the shape name is pyramidal
Shape name and bond angle: 2 bond pairs, 0 lone pairs
linear, 180 degrees
shape name and bond angle: 3 bond pairs, 0 lone pairs
trigonal planar, 120 degrees
shape name and bond angle: 4 bond pairs, 0 lone pairs
tetrahedral, 109.5 degrees
shape name and bond angle: 5 bond pairs, 0 lone pairs
trigonal bipyramidal, 90 degrees and 120 degrees
shape name and bond angle: 6 bond pairs, 0 lone pairs
octahedral, 90 degrees
shape name and bond angle: 3 bond pairs, 1 lone pairs
(trigonal) Pyramidal, 107 degrees
shape name and bond angle: 2 bond pairs, 2 lone pairs
v-shaped/ bent, 104.5 degrees
Define electronegativity
The ability of an atom to attract the bonding electrons in a covalent bond.
Define a metallic bond
Electrostatic attractive force between the positively charged metallic ions and the delocalised electrons in a metallic lattice.
Describe an experiment which can be done to prove the existence of ionic bonding
Electrolysis: Ions migrate to the oppositely charged electrodes. e.g. when done with copper (II) chromate (VI), the solution next to the negative electrode turns blue, and the solution next to the positive electrode turns yellow
Name the three types of intermolecular forces
London forces
Permanent dipole-dipole forces
Hydrogen bonds
How do the boiling points differ between two isomers: one branched, one unbranched (e.g. hexane vs 2,3-dimethylbutane)
The branched isomer has a lower boiling point because there is less surface contact between the molecules. So there are fewer London forces between the chains which take less energy to overcome compared to the straight chain.
Why does CCl4 have a higher boiling point than CF4?
CCl4 has more electrons than CF4
Stronger London Forces in CCl4
Both compounds have polar bonds so there are permanent dipole-dipole forces between molecules
The dipole-dipole forces are stronger in CF4
However the London forces are stronger so CCl4 has the higher bpt
Why does HF have a higher boiling point than HCl
Both have london forces - stronger in HCl as it has more electrons
Both have permanent dipole-dipole forces
HF has hydrogen bonding between molecules. Strongest IMF so takes a lot of energy to overcome so bpt higher for HF
Explain whether the molecules PCl3 and PCl5 are polar
Both have polar P-Cl bonds
PCl5 is not polar: the molecule is symmetrical so the dipoles cancel out
PCl3 is polar: because of the lone pair, the molecule is not symmetrical (it has a pyramidal shape) so the dipoles do not cancel.
Define enthalpy change of solution
The enthalpy change when one mole of an ionic compound is dissolved in a large excess of water
Define enthalpy change of hydration
The enthalpy change when one mole of aqueous ions are formed from their gaseous ions under standard conditions.
Define lattice enthalpy
The enthalpy change when one mole of an ionic lattice is formed from its isolated gaseous ions.
Define enthalpy change of atomisation
The enthalpy change when one mole of gaseous atoms is formed from an element in its standard state.
Define first ionisation energy
The enthalpy change when one mole of gaseous atoms acquires one mole of electrons to form one mole of gaseous negative ions.
Define first electron affinity
The enthalpy change when one mole of an ionic substance is dissolved in a large excess of water.
Why is the second electron affinity of an element always endothermic?
Energy is required to add an electron to a negatively charged anion to overcome the repulsion between them.
What does a lattice enthalpy value tell us?
The strength of ionic bonds:
more exothermic = stronger bonds
What factors affect the value of lattice enthalpy
- Ion size (smaller ions –> stronger bonds)
- Ion charge (higher charge –> stronger bonds)
What factors affect the value of hydration enthalpy?
- Ion size (smaller ions –> stronger interactions with water )
- Ion charge (higher charge –> stronger interactions with water)
What does the value of enthalpy of solution tell you?
If negative –> compound dissolves
If positive –> might dissolve (but because of entropy not enthalpy - see Y13)
If a compound has a theoretical lattice enthalpy and an experimental lattice enthalpy that are nearly the same, what does this tell you?
The bonding in the compound is purely ionc/ 100% ionic.
(The compound matches the theoretical structure of ionic compounds i.e. the ions are perfect spheres which are just touching each other)
If a compound has a theoretical lattice enthalpy and an experimental lattice enthalpy that quite different, what does this tell you?
When the experimental LE is more exothermic than the theoretical LE this tells you the compound has covalent character.
The cation is small (and probably highly charged) and polarises the anion.
The anion is large and polarisable.
The electron cloud of the anion is distorted and it overlaps the electron cloud of the cation increasing the strength of the bond.
