3.1 periodicity Flashcards
(44 cards)
How are the elements arranged in the periodic table?
In the modern periodic table elements are arranged in the order of increasing atomic number.
What is meant by periodicity?
It is the trend in properties across the periodic table.
Across each period elements change from metals to non-metals.
Proton number increases
Ionisation energy generally increases.
Atomic radius decreases
Define first ionisation energy.
Energy required to remove one electron from each atom in one mole of gaseous atoms
Li(g)– Li+ +e-
Define second ionisation energy.
Energy required to remove a second mole of electrons from a gaseous atom
Li+ (g)—- Li2+ + e-
What factors affect ionisation energy?
Atomic radius
Nuclear charge
Electron shielding
How does atomic radius affect ionisation energy?
The larger the atomic radius the smaller the nuclear attraction experienced by the outer electron.
How does nuclear charge affect ionisation energy?
The higher the nuclear charge the larger is the attraction force on the outer electron.
How does shielding affect ionisation energy?
The inner shell electrons repel the outer shell electrons and create a smaller nuclear attraction on the outer electron
Why does first ionisation energy decrease between group 2 and 3?
Group 3 elements have their outermost electron in the p-orbital, whereas group 2 have theirs in the s.
P- orbitals have a slightly higher energy level than S-orbitals and so are marginally further away from the nucleus do they are easier to remove.
Why does first ionisation energy decrease between group 5 and 6?
In both groups the outer electrons are in the P-orbital. However, group 6 outer electron is paired up in the Px orbital so it experiences electron-electron repulsions making it slightly easier to remove.
What is the trend of first ionisation energies down a group?
Shielding increases and so does the atomic radius causing it to decrease.
How does the trend in ionisation energy look across a group?
Number of protons increases (nuclear charge), but the shielding remains the same so atomic radius decreases. So, first ionisation energy increases
Describe the properties of giant metallic lattices.
High melting and boiling points
Electrical conductivity
Malleability and ductility
Why do metallic lattices have high melting and boiling points?
Strong attraction between positive ions and negative, delocalised electrons.
Explain why giant metallic lattices have electrical conductivity?
celocalised electrons are free to mave and carry charge.
Explain why giant metallic lattices are malleable and ductile?
The delocalised electrons can move giving the structure a certain degree of ‘give’ allowing atoms or layers to slide past each other.
Describe the structure, forces and bonding of the elements across period 2.
Li,Be— giant metallic, strong forces between cations and negative electrons, metallic bonding.
B,C– giant covalent, strong covalent bonds between atoms
N2,O2,F2,Ne — simple molecular, weak van der waals between molecules
Describe the structure, forces and bonding of the elements across period 3.
Na, Mg, Al — giant metallic lattices, metallic bonding
Si – giant covalent, strong forces between atoms
P4,S8,Cl2,Ar — simple molecular with weak van der Waals between molecules
Describe melting points across period 3
g1,2,3 melting points increase steadily due to their giant structure and increasing nuclear charge.
G4 has a large increase due to its giant covalent structure and strong covalent bonds
G5,6,7,8, is relatively low due to weak van der Waals forces between molecules (S8 is slightly higher due to it being a larger molecule than P4)
What are the physical properties of group 2 elements?
Reasonably high melting points
Light metals with low densities.
Form colourless(white ) compounds
Describe and explain the configuration down group 2.
Mg - 1s2 2s2 2p6 3s2 [Ne] 3s2
Describe and explain the ionisation energies down group 2.
1st and 2nd ionisation energies decrease
They are strong reducing agents so they are oxidised in reactions and form 2+ ions.
Describe the reaction between group 2 elements and oxygen.
They react vigorously with oxygen.
2Ca +02 — 2CaO
OXIDISED - Ca – Ca2+ + 2e-
REDUCED - 0 +2e- — O2-
Describe the reactions between group 2 and dilute acids.
All group two elements except beryllium react with acids to form salts and hydrogen. (compound with ionic ionic assembly of anions and cations)
Ca+2HCl — CaCl2 +H2
OXIDISED- Ca— Ca2+ +2e-
REDUCED - 2H+ + 2e- —- H2
