Bonding and Structure Flashcards
What is meant by an ionic bond?
A strong electrostatic attraction between the oppositely charged ions
Explain why ionic compounds conduct electricity only when molten/dissolved in water?
There are free moving ions when dissolved in water/molten
Explain why ionic compounds have high melting points:
Because in order to melt an ionic compound, you have to break many strong ionic bonds (between oppositely charged ions), this takes a lot of energy
Explain why magnesium oxide has a higher melting point than lithium chloride:
The charge of an ion is related to the strength of the ionic bond. MgO has greater charge, greater attraction, stronger forces of attraction, .˙. stronger ionic bonding .˙. more energy required to break the bonds
Why do larger ions with greater ionic radius have weaker attractions?
Larger ions that have a greater ionic radius will have a greater attraction to the oppositely charged ion because the attractive forces have to act over a greater distance
Why do larger ions with greater ionic radius have weaker attractions?
Larger ions that have a greater ionic radius will have a greater attraction to the oppositely charged ion because the attractive forces have to act over a greater distance
Explain why converting a lattice into gaseous ions is an endothermic process:
We have to put energy in, to break the strong ionic bonds between oppositely charged ions
What is the trend in lattice energy as you go down the group?
Gets less exothermic as you go down the group
- Lattice enthalpy becomes less exothermic and less negative as the size of the negative ions increases. This indicates weaker attraction between ions and hence weaker ionic bonding
What is the trend in lattice energy as you go across the period?
Gets more exothermic across a period
- Charge increases and produces a greater attraction between these positive ions and negative ions. The ionic radius decreases resulting in the ions in the lattice being closer together producing more attraction
What happens to the ionic radius as you go down the group?
Increases as the atomic number increases. This is because extra electron shells are added
What happens to the ionic radius of a set of isoelectronic?
The ionic radius of a set of isoelectronic ions decreases as the atomic number increases
- As you go through this series of ions the number of electrons stays the same, but the number of protons increases. This means that the electrons are attracted to the nucleus more strongly, pulling them in a little, so the ionic radius decreases
Define covalent bond:
When two atoms share one or more pairs of electrons. Strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond
What is a pure (non-polar) covalent bond:
Bonding electrons shared equally between two atoms
No charges on atoms
What is a polar covalent bond?
Bonding electrons shared unequally between two atoms
Partial charges on atoms
Electrons shared unequally, if electron is pulled to one end, that end will be negative. Electronegativity of oxygen: 3.5, hydrogen: 2.1
What is a dative (coordinate) bond:
Shared pair of electrons that came from the same atom
- In AlCl3, Al has 3 electron pairs around it. Hence, Al has an empty orbital
What are the positive nuclei in covalent bonds attracted to?
- The positive nuclei are attracted to the area of electron density between the two nuclei. But there’s also a repulsion.
- The two positively charged nuclei repel each other, as do the electrons.
What is bond length?
The distance between the two nuclei is the distance where the attractive and repulsive forces balance each other.
What is the trend in bond lengths? Why?
- Increases down a group
- Decreases across a period
This is because the smaller the atomic radius the closer the two atoms can get and the shorter the bond will be. This also tends to make the bond stronger
What is the trend in bond lengths? Why?
- Increases down a group
- Decreases across a period
This is because the smaller the atomic radius the closer the two atoms can get and the shorter the bond will be. This also tends to make the bond stronger
Explain the bond lengths in NH3 vs PH3
Bond length will be greater in PH3 than NH3. This is because P is lower down the group, so has more shells of electrons, and therefore a bigger atomic radius. The nuclei of the two atoms are further from the outer shell electrons so therefore the attraction for the shared pair of electrons will be weaker and the bond will be longer.
What happens when electron density increases?
The higher the electron density between the nuclei, the stronger the attraction between the atoms, the higher the bond enthalpy and the shorter the bond length.
What does molecular shape depend on?
Depends on electron pairs around the central atom
- Shape depends on the number of pairs of electrons in the outer shell of the central atom
What is a pi bond?
The two P orbitals overlap sideways (laterally) to produce regions of electro density above and below the axis joining the two nuclear centres -pi bonds
What is a pi bond?
The two P orbitals overlap sideways (laterally) to produce regions of electro density above and below the axis joining the two nuclear centres -pi bonds
What is the order of the VSEPR shapes?
2 bonded pairs - 180° linear
3 bonded pairs - 120° trigonal planar
2 bonded pairs, one lone pair - 119° bent
4 bonded pairs - 109.5° tetrahedral
3 bonded pairs, one lone pair - 107° trigonal pyramidal
2 bonded pairs, two lone pairs - 104.5° bent
5 bonded pairs - 90/120° trigonal bipyramidal
What is the order of strengths of repulsion:
Lone pair/lone pair > bonded pair/lone pair > bonded pair/bonded pair
What is the lone pair repulsion?
Each lone pair reduces the bond angle by 2.5 degrees
Explain the bond angles in NH3 vs NH4+
- In NH3 all electron pairs repel to achieve maximum separation, but lone pairs repel more than bonding pairs so the shape is pyramidal and the bond angle is 107
- NH3 - three bonding pairs and one lone pair
- In NH4+ all electron pairs are bonding pairs, therefore they all repel equally. Therefore the shape is tetrahedral and has a bond angle of 109.5
- NH4+ - has four bonding pairs (one is dative covalent)
What are London Forces?
- Even in molecules with no polar bonds, there are temporary dipoles due to uneven electron distribution due to the constant movement of electrons
- This induces a temporary/instantaneous dipole in a neighbouring molecule, producing a temporary induced dipole-dipole attraction
What happens if the molecules get bigger?
The bigger the molecule, the greater the London forces
What are permanent dipole-dipole attraction?
Some molecules with polar bonds have an overall dipole (eg/HCl)
- But individual dipoles cancel each other out as the molecule is symmetrical in every direction
What are hydrogen bonds?
- Where an H atom is bonded to a very electronegative atom (i.e. F,O,N)
- The polar bond leaves the H nucleus exposed as H only has one electron
- Therefore there is a strong attraction from the lone pair on the N,O or F of one molecule to the exposed H nucleus of another molecule
Eg/ NH3, H2O.HF
What is the order of the strength of intermolecular forces?
Hydrogen bonding > permanent dipole-dipole > London forces
Explain the boiling point trend across the hydrides of group 7:
HF: highest boiling point as it has hydrogen bonding which is the strongest IMF
HCl, HBr and HI all have dipole-dipole interactions, which are less strong, so have a lower bpt
From HCl → HI, Mr increases, so stronger London forces
What is metallic bonding:
Array of cations in a sea of delocalised electrons
Explain why molecular substances have low melting points
They have weak intermolecular forces in between the molecules therefore require less energy to overcome the forces
Explain why molecular substances do not conduct electricity
They have no charge (no delocalised electrons/ions) .˙. cannot conduct
Why do giant covalent substances have high melting points?
Many and strong covalent bonds, which require a lot of energy to break
Why is diamond hard?
Giant covalent structure therefore many covalent bonds
Why does graphite conduct electricity?
Layers with delocalised electron, therefore have free moving electrons and can carry current
Why is graphite soft and slippery?
Weak forces of attraction between layers therefore can slide
Use of fullerenes and carbon nanotubes:
Delivery tube for cancer drugs
Why do metals conduct electricity?
Delocalised electrons can move and when a current is applied (eg/circuit) they all move in the same direction
Why do metals have high melting points?
There are very strong electrostatic forces of attraction between the oppositely charged particles. Requires a lot of energy to break the forces
