3.1 The periodic table Flashcards
What is periodicity?
Periodicity is the trend in properties that is repeated across each period.
What do elements in the same group have in common?
They have:
- the same number of outer shell e-
- the same type of orbitals
- similar properties
How is abbreviating the electron configuration of some elements useful?
It allows us to clearly see which sub-shell the outer e- are in and where the element can be found in the periodic table.
What is the electron configuration of Na?
1s2. 2s2. 2p6. 3s1.
or [Ne]3s1
What is first ionisation energy?
The first ionisation energy of an element is the energy required to remove 1 electron from each atom in 1 mole of the gaseous element to form 1 mole of gaseous 1+ ions.
What is the equation for the first I.E. of Na?
Na(g) -> Na+(g) + e-
What is I.E. measured in?
kJmol-1
What are the factors that affect I.E.?
- atomic radius - larger atomic radius = smaller nuclear attraction = smaller I.E.
- nuclear charge - higher nuclear charge = greater nuclear attraction = greater I.E.
- electron shielding - inner shells of e- repel outer e- (all -) - more inner shells = greater shielding effect = smaller nuclear attraction = smaller I.E.
What is successive ionisation energy?
Successive ionisation energy is the energy required to remove each e- in turn from an ion in 1 mole of gaseous ions (ion x+) to form 1 mole of gaseous ions (ion (x+1)+)
Explain why each successive I.E. is greater than the one before
- As each e- is removed, there is less repulsion between remaining e- and each shell will be drawn in closer to the nucleus
- due to the positive charge of the nucleus beginning to outweigh the negative charge of the e-.
- As the distance between the nucleus and each e- decreases, the nuclear attraction increases, so more energy is needed to remove each successive e-.
How does successive I.E. support the Bohr model?
Successive I.E.’s provide evidence for shells as once all the e- from the outer shell have been removed, e- from the next shell begin to be removed. This requires a larger amount of energy as there is a smaller distance and less electron shielding, so there is a stronger attraction between the nucleus and outer e-, so we will see a bigger jump in I.E. indicating a different shell closer to the nucleus.
Why is the I.E. of each noble gas the highest in the period?
These atoms have a full outer shell of electrons and a high positive attraction from the nucleus so more energy is needed to remove an electron.
Explain the trend in I.E. going across a period
- nuclear charge (no. of protons in nucleus) increases, so there is a higher nuclear attraction on the outer e-
- atomic radii decreases as increased nuclear charge pulls e- inwards, decreasing the distance between the nucleus and outer e-, causing there to be a higher nuclear attraction on the outer e-
- same no. of inner shells, so electron shielding stays the same
- as the attraction between the nucleus and outer e- increases, more energy needed to remove an e-
- I.E. increases across a period
Why is there a decrease in I.E. between Mg (gr 2) and Al (gr 13)?
- gr 13 elements have their outer e- in a p-orbital whereas gr 2 elements have theirs in an s-orbital
- p-orbitals have slightly higher energy than s-orbitals so the outer shell is further from the nucleus
- therefore the nuclear attraction on the outer e- is slightly weaker so the I.E. is lower
Why is there a decrease in I.E. between N (gr 15) and O (gr 16)
- both gr 15 and gr 16 elements have their outer e- in a p-orbital
- however, for gr 15 elements, each p-orbital only contains a single e-
- in gr 16, the outermost e- is spin paired with another e- in the orbital
- so the e- experience some repulsion, making the first outer e- slightly easier to remove so the I.E. is lower
Explain the large decrease in 1st I.E. between the end of one period and the start of the next.
- at the start of another period, another shell has been added which is further from the nucleus
- this leads to an increase in the distance between the nucleus and the outer shell and an increase in electron shielding
- this causes the nuclear attraction on the outer e- to decrease
- so less energy is needed to remove an e-
- therefore the I.E. is much lower
Explain the trend in I.E. going down a group
- no. of shells increases, atomic radii increases/distance between nucleus and outer shell increases, so weaker nuclear attraction on outer e-
- more inner shells, shielding effect increases, so weaker nuclear attraction
- no. of protons/nuclear charge increases however the resulting increased attraction is outweighed by the increases in distance and shielding
- as the attraction between the nucleus and outer e- decreases, less energy needed to remove an e-
- so I.E. decreases going down a group
Describe the structure of a giant metallic lattice
- positively charged metal ions are held in fixed positions in regular arrangement
- the negatively charged outer e- are delocalised and are spread throughout the structure and are shared between all the atoms in the structure
- the delocalised e- are free to move
- over the whole structure, the charges must balance
- the metal is held together by the strong attractions between the positive metal ions and the negative delocalised e-
Explain why metals have high melting and boiling points
- e- are free to move through the structure but the positive metal ions stay in fixed positions
- the electrostatic attraction between the negative delocalised e- and the positive metal ions is very strong
- a lot of energy is required to overcome the strong metallic bonds and dislodge the ions from their rigid positions in the lattice
Explain why metals are good electrical conductors
- delocalised e- can move freely and carry charge through the structure
Explain why metals are both ductile and malleable
- the delocalised e- are able to move so the layers of atoms are able to slide over each other easily
How do elements change moving across periods 2 and 3?
- metals to non-metals
- solids to gases
Why is Si harder to classify as a metal/non-metal?
It is shiny like metals, but is brittle.
It conducts electricity, but very poorly.
It is an ‘in-between’ element, usually classified as a semi-metal/metalloid.
What happens to the melting points of elements between gr 1 and gr 14?
Between gr 1 and 14, melting points increase steadily because the elements have giant structures.
For each successive group, if an element has a giant metallic lattice, the nuclear charge increases across the period as well as the no. of outer e-, resulting in a stronger attraction.
If the element has a giant covalent lattice, each successive group has more e- which form covalent bonds.
What happens to the melting points of elements between gr 14 and gr 15?
Between gr 14 and gr 15, there is a sharp decrease in melting point because the elements have simple molecular structures. Molecules are held together by weak intermolecular forces.
What happens to the melting points of elements between gr 15 and gr 18?
Between gr 15 and gr 18, the melting points remain relatively low as the elements have simple molecular structures.
Briefly describe the trend in boiling points across periods 2 and 3
Boiling point…
- increases steadily up to gr 14
- decreases sharply between gr 14 and 15
- slays low from gr 15-18
Describe the structure and bonding of diamond
Diamond forms a giant covalent lattice where each carbon atom bonds to 4 other carbon atoms making it extremely strong,
Describe the structure and bonding of graphene
Graphene forms a 2D giant covalent lattice, one carbon atom thick, of interlocking hexagonal carbon rings. It is extremely strong, light and can conduct electricity.
Explain the difference in melting point for the elements Na and Mg
- Mg has more delocalised electrons
- so there will be a stronger electrostatic attraction between the positive metal ions and negative delocalised electrons
- so metallic bonds are stronger
Explain the decrease in atomic radii across the period from Na to Cl
Going across the period…
- the number of shells and amount of electron shielding stays the same
- the nuclear charge increases as the number of protons increases
- this causes the atomic radii to decrease as the increase in positive charge pulls the negative outer electrons closer to the nucleus
Which group are known as the alkali earth metals?
Group 2
Which group has the highest first I.E’s?
Group 8
Which metal is the only metal not to be a solid at room temperature?
Mercury
What shape is the structure of diamond based on?
Tetrahedral
