Chapter 7 - Periodicity Flashcards
What order are the elements arranged in?
Increasing atomic number
What do the groups in the periodic table tell you?
- Number of electrons in the outer shell
- Elements that have similar chemical properties
What do the periods in the periodic table tell you?
Number of the highest energy electron shell
What is periodicity?
Repeating trend in properties of the elements across a period
Which groups are the s-block elements?
- Group 1 and 2
- Helium
Which groups are the d-block elements?
Transition metals
Which groups are the p-block elements?
Groups 3, 4, 5, 6, 7, 8
What is the first ionisation energy?
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
What is the first ionisation energy of Na(g) ?
Na(g) -> Na+ (g) + e-
First electron lost will be the _________ energy level and will experience the _________ attraction from the nucleus
- Highest
- Least
What happens to first ionisation energy when atomic radius increases? Explain why
Greater distance between nucleus and outer electrons
Less nuclear attraction
Lower ionisation energy
What happens to first ionisation energy when nuclear charge increases? Explain why
More protons
Greater attraction between nucleus and outer electron
Higher ionisation energy
What happens to first ionisation energy when there is more electron shielding Explain why
Shielding effect increases when number of inner shells increase
Less attraction between nucleus and outer electrons
Lower ionisation energy
Give the general relationship between nuclear attraction and first ionisation energy?
Less attraction between nucleus and outer electrons
Lower ionisation energy
Greater attraction between nucleus and outer electrons
Higher ionisation energy
What is the shielding effect?
Inner shell electrons repel outer shell electrons
Element has as many ionisation energies as there are ___________
Electrons
What is second ionisation energy?
Energy required to remove one electron from each atom in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions
Why are second ionisation energies greater than the first?
After first electron is lost, remaining electron is pulled closer to nucleus, stronger nuclear attraction = more energy needed to remove electron = higher ionisation energy
Electrons with _________ ionisation energies are ___________ to the nucleus
Higher
Closer
What does the large difference/jump in ionisation energies represent?
Change from one shell to another
What is the trend in first ionisation energies down a group?
- Atomic radius increases
- More inner shells so shielding increases
- Nuclear attraction decreases
- First ionisation energy decreases
s.n. nuclear charge increasing is outweighed by the effects of increased radius and (to a lesser extent) increased shielding
What is the trend in first ionisation energies across a period?
- Nuclear charge increases
- Same shell numbers so similar shielding
- Nuclear attraction increases
- Atomic radius decreases
- First ionisation energy increases
Why does boron have a lower first ionisation energy than beryllium? (hint: sub-shell is the different)
- Start of filling of 2p sub shell
- 2p sub shell is higher energy than 2s sub shell
- 2p electron easier to remove
Why does oxygen have a lower first ionisation energy than nitrogen? (hint: sub-shell is the same)
- Start of electron pairing in the p-orbitals of 2p sub shell
- In O the paired electrons in one of the 2p orbitals repel each other
- Easier to remove one of the paired electrons from O
Why does aluminium have a lower first ionisation energy than magnesium? (hint: sub-shell is the different)
- 3p sub-shell in Al has a higher energy level than 3s sub-shell in Mg
- Easier to remove 3p electron from Al
Why does sulphur have a lower first ionisation energy than phosphorus? (hint: sub-shell is the same)
- P has 3 electrons in 3p sub shell which are all unpaired
- S has 4 electron in 3p sub shell which has one pair in an orbital
- Paired electrons repel so easier to remove an electron from S
Describe metallic bonding?
Strong electrostatic attraction between cations and delocalised electrons
Describe the electrical conductivity of metallic structures?
Conduct electricity in solid and liquid states since the delocalised electrons are able to carry charge
Describe the melting and boiling points of metals?
High M.P. and B.P. as lots of energy needed to break strong electrostatic attraction
Describe the solubility of metals?
Do not dissolve in water
Describe giant covalent structures?
Boron, carbon and silicon atoms held together by covalent bonds to form giant covalent structures
What is the shape of the structures that carbon and silicon atoms form?
Tetrahedral
Describe the electrical conductivity of giant covalent structures?
Non conductors of electricity (except graphene and graphite)
Why can diamond not conduct electricity?
All four outer shell electrons involved in bonding so no delocalised electrons to carry charge
Why can graphene (single layer of graphite) conduct electricity?
Only three of the carbon outer shell electrons are involved in bonding so there are delocalised electrons to carry charge
Describe the melting and boiling points of giant covalent structures?
High M.P. and B.P. as lots of energy needed to break strong covalent bonds
Describe the solubility of giant covalent structures?
Insoluble in almost all solvents as covalent bonds are too strong to be broken
What is the general trend in melting points across period 2 and 3?
- Melting point increases from group 1-4
- Sharp decrease in melting point from group 4-5
- Melting point lower from group 5-0
Why is there a sharp decrease in melting point from group 4-5?
Marks change from giant to simple molecular structures
What structure to metals form?
Giant metallic lattice structures
