Periodicity - Period 3 Flashcards
When discussing period 3 elements, you can compare atomic properties and physical properties.
What are the 4 topics you can talk about relating to atomic properties?
What are the 3 topics you can talk about relating to physical properties?
Atomic properties: a) electronic structures b) first ionisation energy c) atomic radius d) electronegativity Physical properties: a) structures of the elements b) electrical conductivity c) melting and boiling points
List which blocks the period 3 elements are in, and why.
sodium and magnesium are in the s-block because their outermost electron is held in an s orbital
the remaining elements (al, si, p, s, cl, ar) are in the p-block because their outermost electrons are held in p orbitals
What is first ionisation energy?
first ionisation energy is the energy that is required to remove `1 electron from every atom in a mole of gaseous atoms to produce 1 mole of electrons and 1 mole of ions each with a single positive charge
What is the general trend of ionisation energy across period three?
first ionisation energy increases across the period
Explain the trend of ionisation energy across a period, with a few words for each point
- increased nuclear charge
- same distance
- same screening
Explain in detail, why the trend for ionisation energy generally decreases or increases
Increases because
- across a period the atomic radii is not hugely different so this factor does not influence the trend
- all electrons are in the same energy level (energy level 3 for period 3) so the amount of screening across a period does not change
- but the amount of protons increases as you go along so there is a greater attraction between nucleus and electrons
Where is the general trend broken in period 3, and why?
- dips at group 3, aluminium
because: previous elements are in s block, aluminium however holds outermost electron in p orbital. p orbitals are more remote, the electron is held at greater distance from nucleus and there is additional shielding - dips at group 6, sulphur
because: previous element, phosphorus, has unpaired electrons in p orbital. In sulphur, the extra electron means an orbital is occupied by a pair, this causes extra repulsion so it is easier to remove the electron
Which period 3 element has the lowest IE?
Sodium
Which period 3 element has the largest IE? Mention why.
Argon (it is a noble gas with a full outer shell of electrons)
Summarise the trends of IE across period 3
- general increase across period 3
- dips at aluminium and sulphur
Across period 3, what happens to the atomic radii of the atoms?
- it decreases between sodium through to chlorine
- argon is much bigger however
What causes atomic radii to decrease between Na and Cl?
- there is no difference in shielding
- but nuclear charge is increasing so the attraction between the nucleus and outermost electrons is greater which pulls them closer together
Why does increasing proton number effect the atomic radii in the elements of period three (excluding argon)?
- Na, Mg, Al are metallic structures
- Si, P, S and Cl are molecular substances with covalent bonds
- metallic and covalent radii are the measure of distance from the nucleus to the bonding pair of electrons
- increasing proton number pulls the bonding electrons more tightly to it hence decreasing the atomic radius
Why is argon’s atomic radius much larger than chlorine, when you might expect it to be smaller due to the increased proton?
Argon only has van der Waals attractions, there are no covalent bonds and bonding pairs of electrons to pull closer.
What is electronegativity?
Electronegativity is the measure of the tendency of an atom to attract a bonding pair of electrons towards it.
What scale is electronegativity measured on?
The pauling scale that runs 0-4.0
Which element in period 3 is the most electronegative?
Chlorine (think: FONCl)
What is the trend in electronegativity across a period and why?
There is a steady increase as each element has an extra proton
Which element is not assigned an electronegativity value at all and why?
Argon because it does not form compounds
What are the three types of structure you see in period 3? Name the elements that belong to each group.
Metallic - Na, Mg, Al
Giant covalent - Si
Simple Molecular - P4, S8, Cl2, Ar
Give a brief description of the general structures of Na, Mg and Al
- metallic structures
- closely packed metal ions surrounded by a sea of delocalised electrons that have been delocalised from the atoms outer shell of electrons
Na: ions have 1+ charge, 1 electron per atom in the sea of electrons, 8-co-ordinated structure
Mg: ions have 2+ charge, 2 electrons per atom in the sea of electrons, 12-co-ordinated structure
Al: ions have 3+ charge, 3 electrons per atom in the sea of electrons, 12-co-ordinated structure
What is meant by 8-co-oridinated and 12-co-ordinated structure, and which is more effective and why
8-co-ordinated structure means that each atom is touched by 8 other atoms. 12-co-ordinated structure means that each atom is touched by 12 other atoms.
12-co-ordinated is more effective because there is less space wasted
Which period 3 element has a giant covalent structure? Give a description of this element.
- it is silicon
- 3 dimensional covalent bonds like diamond
- silicon is a metalloid (properties lie intermediate between metal and non metal)
In the simple molecular structures, how are the atoms held together?
- by covalent bonds (in P4, S8 and Cl2)
- argon is monatomic so there are no covalent bonds
- in the liquid or solid state molecules are held together by van der waals attractions
