Chapter 7 - Periodicity Flashcards
The periodic table, Ionisation energies and Periodic trends in bonding and structure.
Who created the modern periodic table?
Dmitri Mendeleev
How were the elements ordered by Mendeleev?
By increasing atomic mass
What is the name for the vertical columns of the periodic table?
Groups
What is the name for the horizontal rows of the periodic table?
Periods
What is the periodic trend in electron configuration?
The sub shells of n energy level fill up
What is ionisation?
The removal of one or more electrons from an atom
Define first ionisation energy:
The energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form one mole of gaseous 1+ ions
How (and why) does atomic radius affect ionisation energy?
• Greater distance between the nucleus and outer electrons
• Attraction between them is reduced
• So ionisation energy decreases
How (and why) does nuclear charge affect ionisation energy?
• More protons increases the nuclear charge
• There is a greater attraction between the nucleus and outer electrons
• Atomic radius decreases
• So ionisation energy increases
How (and why) does electron shielding affect ionisation energy?
• Inner electrons repel outer electrons
• Which “makes outer electrons easier to remove”
• So ionisation energy decreases
Why does ionisation energy decrease going down a group?
• N° electron shells increases down a group
• Greater effect of electron shielding
• So ionisation energy decreases
Why does ionisation energy increase across a period?
• N° protons in the nucleus increases across a group
• Nuclear charge increases
• Atomic radius decreases
• (electron shielding is the same)
• So ionisation energy increases
Why do successive ionisation energies increase?
• Electron-electron repulsion decreases as electrons are removed
• Remaining electrons are drawn closer to the nucleus
• So effective nuclear charge increases
• So successive ionisation energies increase
Define second ionisation energy:
The energy required to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form one mole of gaseous 2+ ions
(Add 1 to each ion for third, fourth, etc ionisation energies)
What causes the large jumps in successive ionisation energies?
• Moving into a lower energy level
• The now outer electrons are closer to the nucleus
• So effective nuclear charge increases
• So the successive ionisation energy increases with a large jump
What predictions of electron configuration can be made from a graph of successive ionisation energies?
• How many electron shells are present (per jumps in ionisation energy)
• Hence, how many outer electrons there are (per how many ionisation energies before the first jump)
• Thus, the element’s group
In period 2, explain the fall from beryllium to boron of first ionisation energies
• Beryllium’s outermost electrons are both in the 2s orbital (1s2 2s2)
• Boron has an outermost electron in the 2p orbital (1s2 2s2 2p1)
• (The 2p orbital is farther from the nucleus)
• So, alongside greater electron shielding from the 2s orbital, the 2p1 electron experiences less nuclear attraction and first ionisation energy is lower
Explain the trend of atomic radii decreasing across a group
• Nuclear charge increases
• So nuclear attraction increases, pulling electrons closer to the nucleus
• Atomic radius decreases
What is metallic bonding?
The strong electrostatic attraction between fixed metal cations and a “sea” of delocalised electrons
What are common properties of metals?
• Strong metallic bonds
• High electrical conductivity
• High melting and boiling points
Why does the melting point and boiling point increase across the metals of a period?
• Nuclear charge increases
• Thus a decreased atomic radius
• More delocalised electrons
• So overall there is a greater attraction between cations and delocalised electrons, requiring more energy to overcome
• So an increase in melting/boiling point
In period 3, why does silicon have the highest melting point?
• Forms a giant covalent structure
• Strong covalent bonds throughout the structure require lots of energy to break
• So has a very high melting point
What are the properties of giant covalent lattices?
• High melting and boiling points
• Insoluble in almost all solvents
• Do not conduct electricity (exc. graphite/graphene)
Why are giant covalent lattices insoluble?
The covalent bonds present are too strong to be broken by the interactions with a solvent
Why do most giant covalent lattices not conduct electricity?
All four outer shell electrons are involved in covalent bonding
Why can graphite/graphene conduct electricity?
• Graphite is comprised of parallel layers of hexagonally arranged carbon atoms
• Where only 3 of 4 carbon’s outer electrons are used in covalent bonding
• So there is 1 delocalised electron per atom to carry charge throughout the structure
