the periodic table Flashcards
why does effective nuclear charge increase across periods 2 and 3?
across the period
increase in nuclear charge
while shielding effect by inner principal quantum shells of electrons remains relatively constant
why do cations have smaller radii compared to atoms?
atoms usually lose electrons to form cations hence they usually have one less principal quantum shell
why do anions have a larger radii compared to atoms?
electrons are added to the outermost principal quantum shell of the atom hence there is greater electron-electron repulsion in the outermost principal quantum shell
describe the trend of electrical conductivity across period 3
high conductivity for metals Na,Mg and Al and increases form Na to Al = giant metallic structures with delocalised electrons = number of delocalised electrons increase from Na to Al
conductivity drops sharply at Si = giant covalent structure semi-conductor
conductivity drops to 0 at P and remains till Ar = P to Cl have simple covalent structures and Ar is monoatomic hence no mobile charge carriers to conduct electricity
can graphite conduct electricity?
yes
it has delocalised valence electrons along the layers to conduct electricity in the direction parallel to the layers
explain the first drop in 2-3-3 trend for trend of ionisation energy across a period
electron to be removed are from different sub shells therefore one requires less energy
explain the second drop in 2-3-3 trend for ionisation energy across a period
coulombic repulsion between paired electrons makes it easier to remove one of the paired electrons hence less energy is required
describe and explain the trend of ionisation energies down a group
IE decreases down a group
as atomic radius increases
shielding effect increases due to the large increase in number of principal quantum shells
this largely cancels out the increase in nuclear charge
valence electrons are less attracted to positive nucleus
less energy
explain why the melting point of the following elements are as such
sulfur > phosphorus > chlorine > argon
sulfur exists as S8, phosphorus as P4, chlorine as Cl2 and argon as Ar
hence the number of electrons increase from sulfur to argon
the greater the number of electrons the stronger the instantaneous dipole-induced dipole interactions between molecules
why does SCl6 not exist
due to steric hindrance about the central S atom
reaction of NaCl(s) with water
dissolves readily to form a neutral solution, pH 7.
hydrolysis does not take place as Na+ has low charge density and hence low polarising power
NaCl(s) –> Na+(aq) + Cl-(aq)
reaction of MgCl2(s) in water
dissolves readily to form a weakly acidic solution, pH 6.5. partial hydrolysis of Mg+(aq) ions occur
MgCl2(s) + 6H2O(l) –> [Mg(H2O)6]2+(aq) + 2Cl-(aq)
[Mg(H2O)6]2+(aq) + H2O(l) [Mg(H2O)5(OH)]+(aq) + H3O+(aq)
reaction of AlCl3(s) with water
dissolves readily to form acidic solution, pH 3. due to its high charge density, Al3+ is highly polarising and weakens the O-H bonds in water molecules causing them to break and release hydrogen ions
AlCl3(s) + 6H2O(l) –> [Al(H2O)6]3+(aq) + 3Cl-(aq)
[Al(H2O)6]3+(aq) + H2O(l) [Al(H2O)5(OH)]2+(aq) + H3O+(aq)
reaction of SiCl4(l) with water
SiCl4 dissolves in water to form a strongly acidic solution, pH 2. it undergoes hydrolysis because Si has an energetically accessible vacant 3d orbital for dative bonding with water molecules
SiCl4(l) + 2H2O(l) –> SiO2(s) + 4HCl(aq)
reaction of CCl4 with water
it doesn’t react with water as C doesn’t have energetically accessible vacant orbitals for dative bonding
reaction of PCl5(s) and *PCl3(l) with water
PCl5(s) and *PCl3(l) undergoes hydrolysis in water to give a strongly acidic solution, pH 2, as P can use its energetically accessible vacant 3d orbitals for dative bonding
In limited/cold water,
PCl5(s) + H2O(l) –> POCl3(l) + 2HCl(aq)
in excess water,
POCl3(l) + 3H2O(l) –> H3PO4(aq) + 3HCl(aq)
overall equation,
PCl5(s) + 4H2O(l) –> H3PO4(aq) + 5HCl(aq)
why are the oxidation numbers of period 3 always positive
because oxygen is more electronegative than all the period 3 elements
pH of Na2O solution
13
pHof MgO solution
8
pH of Al2O3 solution
7
pH of SiO2 solution
7
pH of P4O10 solution
2
pH of SO3 solution
2
reaction of Na2O with water
reacts vigorously with water to form NaOH which is strongly alkaline, pH 13.
Na2O(s) + H2O(l) –> 2NaOH(aq)
reaction of MgO with water
reacts to form Mg(OH)2(s) which dissolves sparingly to give a weak alkaline solution, pH 8.
MgO(s) + H2O(l) –> Mg(OH)2(s)
Mg(OH)2(s) Mg2+(aq) + 2OH-(aq)
reaction of Al2O3 with water
insoluble as the energy released during hydration through formation of ion-dipole interactions is insufficient to overcome the energy required to overcome the strong electrostatic forces of attraction between Al3+ and O2- ions in a giant ionic lattice hence pH 7
Na2O reaction with acid or base
reacts with acid to form salt and water
Na2O(s) + 2H+(aq) –> 2Na+(aq) + H2O(l)
MgO reaction with acid or base
reacts with acid to form salt and water
MgO(s) + 2H+(aq) –> Mg2+(aq) + H20(l)
Al2O3 reaction with acid or base
reacts with both to form salt and water
with acid,
Al2O3(s) + 6H+(aq) –> 2Al3+(aq) + 3H2O(l)
with alkali,
Al2O3(s) + 2OH-(aq) + 3H2O(l) –> 2[Al(OH)4]-(aq)
reaction of SiO2 with water
insoluble in water as a large amount of energy is required to break the strong covalent bonds in the giant covalent structure
= doesn’t affect pH and pH remains at 7
reaction of SiO2 with acid or base
reacts with concentrated alkali to form salt and water
SiO2(s) + 2OH-(aq) –>SiO32-(aq)+H2O(l)
reaction of P4O10 with water
reacts violently to give an acidic solution, pH 2.
P4O10(s) + 6H2O(l) –> 4H3PO4(aq)
reaction of P4O10 with acid or base
reacts violently with alkali to form salt and water
P4O10(s) + 12OH-(aq) –> 4PO43-(aq) + 6H2O(l)
reaction of SO3 with water
reacts to give an acidic solution, pH 2.
SO3(l) + H2O(l) –> H2SO4 (aq)
reaction of SO3 with acid or base
reacts with alkali to form salt and water
SO3(l) + 2OH-(aq) –> SO42-(aq) + H2O(l)
*reaction of Cl2O(g) and Cl2O7(l) with water
acidic solution with pH 2
Cl2O(g) + H2O(l) –> 2HClO(aq)
Cl2O7(l) + H2O(l) –> 2HClO4(aq)
*reaction of Cl2O(g) and Cl2O7(l) with acid or base
reacts with alkali to form salt and water
Cl2O(g) + 2OH-(aq) –> 2ClO-(aq) + H2O(l)
Cl2O7(l) + 2OH-(aq) –> 2ClO4-(aq) + H2O(l)
what are the products of thermal decomposition of a group 2 carbonate
metal oxide and carbon dioxide
MCO3(s) –> MO(s) + CO2(g)
why does thermal decomposition occur in group 2 nitrates, carbonates and hydroxides?
the metal cation polarises the electron cloud of the large nitrate/carbonate/hydroxide anion to such an extent that the covalent bonds in the anion are weakened. hence these weakened bonds readily break when heat is applied to the compounds
why does the thermal stability of group 2 carbonates increase down the group?
down the group
radius of metal cation increases hence charge density decreases
ability of metal cation to polarise the electron cloud decreases
C-O bonds are weakened to a smaller extent
physical state and colour of fluorine at rtp
pale yellow gas
physical state and colour of chlorine at rtp
pale greenish-yellow gas
physical state and colour of bromine at rtp
reddish-brown liquid
physical state and colour of iodine at rtp
black solid
colour of fluorine when dissolved in non-polar solvent
colourless
colour of chlorine when dissolved in non-polar solvent
pale yellow
colour of bromine when dissolved in non-polar solvent
orange-red
colour of iodine when dissolved in non-polar solvent
purple
colour of fluorine when dissolved in water
colourless
colour of chlorine when dissolved in water
pale yellow
colour of bromine when dissolved in water
orange
colour of iodine when dissolved in water
brown
why does thermal stability of hydrogen halides decrease down the group?
size of halogen atoms increase hence their valence orbitals become more diffused resulting in less effective orbital overlap between the small H atom and less energy is required to break the weaker H-X bond
