Periodicity Flashcards
Which side of the periodic table are metals?
Left
What happens to metallic character as you go down a group?
increases
Which group are called the alkali metals?
1
Which group are called the alkaline earth metals?
2
Which group are called the halogens?
7
Which group are called the noble gases?
0
As you go down a group what happens to the number of electrons in the outer shell?
Nothing, they have the same number of electrons in their outer shell but they are in a higher energy level.
Why do elements in the same group have similar chemical properties?
Same number of electrons in their outer shells
Why do we call the block containing group 1 and 2 the s block?
The electron(s) in the highest energy level are in s orbitals
What is the atomic radius?
The distance from the centre of the nucleus to the outermost electron
How do you calculate the atomic radius?
Measure the distance between two covalently bonded atoms nuclei (must be the same element) and divide by 2.
What happens to atomic radii as you go across a period?
Decreases
Extra proton in nucleus, increase positive charge, electrons pulled closer.
What is the definition of first ionisation energy?
The energy required to remove 1 mole of outermost electrons from one mole of a gaseous element.
Describe the pattern of first ionisation energies across a period.
Rise. Then a drop. Then 2 rises. Then a drop. Then 2 rises.
Explain the pattern of first ionisation energies across a period.
Rise-increase in nuclear charge. Drop-e- removed from p orbital not s (further away). Risex2-increase nuclear charge. Drop-e- paired in orbital, repel each other so easier to remove. Risex2 increase nuclear charge.
What is the definition of electronegativity?
Power of an atom to attract electron density from a covalent bond. The more electronegative, the stronger the pull on electrons.
What happens to electronegativity as you go across a period?
Increases (increase in positive nuclear charge)
Explain the trend in electrical conductivity across a period
Increases group 1-3 as more delocalised electrons. Decreases at group 4, they form lattices but at high temps can conduct. Group 5-7 don’t at all as they form covalent bonds, no charged mobile particles.
Explain the increase in melting/boiling point across group 3 from sodium to aluminium
Held together by metallic bonds, increase boiling points as more delocalised electrons and a higher nuclear charge.
Explain the increase in melting point across group 3 from aluminium to silicon
Giant covalent structure in silicon means there’s lots of strong bonds to break as each silicon atom is bonded to 4 others.
Explain the decrease in melting/boiling points across group 3 from silicon to argon
Phosphorous-argon are all made of separate molecules which only have LDF holding them together which are easy to break, the smaller the molecules the easier it is to break the LDF.
Why is the boiling point of silicon lower than expected (lower than aluminium)?
Most the bonds were broken in the process of melting, and the bonds which are left are lower in energy than in liquid aluminium.
What 3 factors affect ionisation energy?
Atomic radius
Number of inner electrons (shielding)
Number of protons (nuclear charge)
Why does the first ionisation energy decrease down a group?
More shielding
Larger atomic radius
