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?
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 one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
How does atomic radius affect ionisation energy?
Greater distance between the nucleus and outer electrons, the attraction is lower.
How does nuclear charge affect ionisation energy?
More protons creates a greater attraction between the nucleus and the outer electrons. therefore higher ionisation energy
How does electron shielding affect ionisation energy?
Inner shell electrons repel outer shell electrons, called shielding. This reduces the attraction between the nucleus and the outer electrons. therefore lower ionisation energy
Why does ionisation energy decrease going down a group?
More electrons shells so the outer electrons are further away and there is a greater shielding effect.
Why does ionisation energy increase across a period?
The number of protons in the nucleus increases so nuclear charge increases causing atomic radius to decrease, whilst shielding stays the same.
Why do successive ionisation energies increase?
There are less electrons so the nuclear attraction on the remaining electrons will be greater.
Define second ionisation energy
The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.
What causes the large jumps in successive ionisation energies?
Moving down to a closer shell, as these electrons are closer so experience a greater nuclear attraction.
What predictions can be made from a graph of successive ionisation energies?
The number of electrons in the outer shell, the group of the element in the periodic table and thus the identity of the element.
Explain the trend of first ionisation energy down a group
Atomic radius increases,
More inner shells so shielding increases,
Nuclear attraction on outer electrons decreases,
First ionisation energy decreases.
Explain the general trend of first ionisation energy across a period
Nuclear charge increases, Same shell: similar shielding, Nuclear attraction increases, Atomic radius decreases, First ionisation energy increases.
In period 2, explain the fall from beryllium to boron of first ionisation energies
The new electron enters the 2p sub shell, which is slightly further away from the nucleus than the 2s sub shell.
In period 2, explain the fall of ionisation energy from nitrogen to oxygen
Nitrogen’s electrons in the 2p sub shell are unpaired so oxygen’s 8th electron is paired, causing repulsion and a lower ionisation energy.
Explain the trend of atomic radii decreasing across a group
Nuclear charge increases,
Nuclear attraction increases,
Atomic radius decreases.
What is metallic bonding?
Each atom donates an outer shell electron, which becomes delocalised. This creates cations.
What causes metallic bonding?
Strong electrostatic attraction between the fixed cations and the 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?
Number of delocalised electrons per atom and charge on cation increase, so stronger electrostatic attraction.
In period 3, why does silicon have the highest melting point?
Forms a giant covalent lattice, where each atom is covalently bonded to four others.
What are the properties of giant covalent lattices?
High melting and boiling points
Insoluble in almost all solvents
Do not conduct (except for graphite and graphene)
formed from carbon, silicon anf boron
Why are giant covalent lattices insoluble?
Covalent bonds holding it together are too strong to be broken by interaction with solvents.
Why do most giant covalent lattices not conduct electricity?
All four outer shell electrons are involved in covalent bonding.
Why do simple molecules have low melting points?
Weak induced dipole-dipole forces between molecules are easy to break.
