Chapter 7 - Periodicity Flashcards
How did Mendeleev arrange elements in the periodic table?
- In order of atomic mass
- Elements with similar properties were put into groups
- Left gaps for undiscovered elements
How are periods in the periodic table arranged?
Elements with the same number of electron shells
How are groups in the periodic table arranged?
Elements with the same number of electrons in their outer shell
What groups are in s-block?
Groups 1 and 2
What groups are in p-block?
Groups 3 to 0
What groups are in d-block?
Transition metals
What groups are in f-block?
Lanthanides and Actinides
Why are elements classified into blocks in the periodic table?
Elements in the same block have their outer electrons in the same orbital
Define ‘Periodicity’
A repeating trend in properties of elements
Why does atomic radius decrease across a period?
- There is an increased nuclear charge for the same umber of electron shells
- The outer electrons are pulled in closer to the nucleus as increased charge produces a greater attraction
- Therefore reducing the atomic radius
What does the atomic radius increase down the groups?
- As you go down the group there are more electron shells
- This increases the distance between the nucleus and the outer electrons reducing the power of attraction
- More shells increases the electron shielding , increasing the radius of the nucleus
What are group 1 metals called?
Alkali metals
What are group 2 metals called?
Alkaline earth metals
What are group 7 elements called?
Halogens
What are group 0 elements called?
Noble gases
Define ‘First ionisation energy’
The energy required to remove one mole of electrons from one mole of atoms in their gaseous state to form one mole of +1 ions
Define ‘Second ionisation energy’
The energy required to remove one mole of electrons from one mole of +1 ions in their gaseous state to form one mole of +2 ions in their gaseous state
What are the factors that affect ionisation energy?
- Atomic radius
- Nuclear charge
- Shielding
How does the atomic radius impact ionisation energy?
- As the atomic radius increases, the force of attraction between the positive nucleus and the outer electrons decreases
- This means that less energy is required to overcome the weak forces of attraction between the nucleus and the outer electrons to remove the electron
How does the nuclear charge impact ionisation energy?
- The greater the number of protons in the nucleus, the greater the force of attraction between the nucleus and the outer electrons
- This requires more energy to overcome the strong forces of attraction between the nucleus and the outer electrons to remove the electron
How does shielding impact ionisation energy?
- Electrons in the outer shell are repelled by electrons in inner shells
- This shielding effect reduces the attraction between the outer electrons and the nucleus
- Less energy is required to overcome the weak forces of attraction, to remove the electron
How does first ionisation energy vary as you go down a group?
- The first ionisation energy decreases as you do down the group
- As you go down a group the atomic radius increases
- This means that the outer electron shell is further away from the nucleus
- Also, there is an increases in electron shielding between the nucleus and outer electrons
- The attraction between the positive nucleus and outer electrons is weak
- So less energy is required to overcome the weak force of attraction and remove an electron
How does first ionisation energy vary across a period?
- The first ionisation energy increases as you go across a period
- As you move across a period the nuclear charge increases as the number of protons increases which increases the forces of attraction between the nucleus and the electrons
- The atomic radius decreases causing an increase in the forces of attraction between the outer electron and the nucleus
What are the exceptions to the ionisation energy across a period?
- Boron and Oxygen do not fit the pattern
- Berylliums electron is being removed from the 2s sub-shell whereas Borons is being removed from the 2p sub-shell
- The 2p sub-shell has a higher energy than the 2s sub-shell meaning less energy is required to remove the outer electron of Boron
- Nitrogens electron is in a seperate 2p orbital, whereas in oxygen one oribitals contains a pair of electrons
- This means that less energy is required to remove one of these electrons rather than in seperate orbitals
Define ‘Metallic bonding’
The strong electrostatic force of attraction between positive ions and delocalised electrons
What are the 3 factors that affect metallic bonding?
- Nuclear charge
- The number of delocalised electrons
- Size of the ion
Why do metals conduct electricity?
Metals contain delocalised electrons that are free to move and carry charge through the structure as both a solid and liquid
Why do metals have high melting and boiling points?
Metals contain strong electrostatic forces of attraction between the cations and delocalised electrons. This means that large amounts of energy through high temperatures are needed to overcome the forces of attraction
What structure do metal bonds make?
Giant metallic lattice
Why are metals malleable?
Contain uniform layers of positive ions that are able to slide over eachother
Why do giant covalent lattices have high melting and boiling points?
They contain very strong covalent bonds which require high amounts of energy to break the bonds
Why are giant covalent lattices insoluble?
The covalent bonds in them are too strong to be broken by solvents such as water
