Chapter 7 - Periodicity Flashcards
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 of 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 the trend across period 2 ?
Across period 2, the 2s sub shell fills with two electrons, followed by the 2p sub shell with six electrons
What is the trend across period 3 ?
Across period 3, the 3s sub shell fills with two electrons, followed by the 3p sub shell with six electrons
What is the trend across period 4 ?
- Across period 4, although the 3d sub shell is involved, the highest shell number is n=4.
- From the n=4 shell, only the 4s and 4p sub shells are occupied
What are blocks ?
The elements in the periodic table can be divided into blocks corresponding to their highest energy sub shell
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 other electrons
What is shielding ?
It is when inner shell electrons repel outer shell electrons
How does electron shielding affect ionisation energy ?
This reduces the attraction between the nucleus and the outer electrons
Why does ionisation energy decrease going down a group ?
More electron 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 therefore 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
- 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 so 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 from nitrogen to oxygen of ionisation energies
Nitrogen’s electrons in the 2p sub shell are unpaired and oxygen’s 8th electron is paired, causing repulsion and a lower ionisation energy.
Explain the trend of atomic radii across a group
- Nuclear charge increases
- Nuclear attraction increases
- Atomic radius decreases
What is metallic bonding ?
Metallic bonding is when each atom donates an outer shell electron, which then becomes delocalised. This creates cations.
What causes metallic bonding ?
Strong electrostatic force of attraction that occurs between 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 theres a 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)
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
