Periodic Table Flashcards
What factors affect radius and first IE ? With explanation
Nuclear charge
- increases when number of protons increase
Shielding effect
- increases when number of inner electron shells increase
Effective nuclear charge
- accounts both NC and SE, if both increase ENC cancels out
- only NC increase, ENC increase
How does atomic radius change across a period ?
Across a period
- NC increases as number of protons increases
- SE relatively constant as electrons are added to valence shell and there is no change in number of inner electron shells
- ENC increases leading to stronger attraction between the nucleus and valence electrons, valence electrons are pulled closer to nucleus causing atomic radius to DECREASE
How does atomic radius change down a group ?
Down a group
- NC increase as number of protons increase
- SE increases as number of inner shell electrons increase
- ENC remains relatively constant
- the distance between the nucleus and valence electrons increase due to increase in number of electron shells, resulting in weaker attraction between nucleus and valence electrons, thus atomic radius INCREASES
How does cation radius change across third period ?
From Na+ to Si4+
- NC increases as number of protons increase
- SE constant as cations have same number of electrons
- ENC increases leading to stronger attraction between the nucleus and valance electrons so the valence electrons are pulled closer to the nucleus causing ionic radius to DECREASE
How does anion radius change across the third period ?
From P3- to Cl-
- NC increases as number of protons increase
- SE same as anions have same number of electrons
- ENC increases leading to stronger attraction between valence electrons and nucleus so valence electrons are pulled closer to nucleus casing ionic radius to DECREASE
How is the ionic radius different for cation and anions across the third period ?
Anions have larger radii as there is an extra shell of electrons present
How does ionic radius change down a group ?
Down a group
- NC increases as number of protons increase
- SE increases as number of inner shell electrons increase
- ENC remains constant
- the distance between the nucleus and valence electrons increases due to increase in number of electron shells resulting in weaker attraction between the nucleus and valence electrons, causing ionic radius to INCREASE
How does first IE change across a period ?
Across a period
- NC increases as number of protons increase
- SE remains constant as number of inner shell electrons are the same
- ENC increases leading to stronger attraction between the nucleus and valence electrons so more energy is required to remove the valence electron, hence first IE INCREASES
What are the deviations in first IE across third period ? With explanation
First IE for Aluminium is lower than Mg
- first electron to be removed from AL is from 3p subshell while first electron to be removed from Mg is 3s subshell
- electron in 3p subshell is further away from nucleus compared to 3s subshell, the 3p subshell is of higher energy level than 3s subshell and its electron is less strongly attracted to nucleus
- less energy is required to remove an electron from 3p subshell than a 3s subshell
First IE for S is lower than P
- first electron to be removed in S is a paired 3p electron compared to unpaired 3p electron for P
- easier to remove paired 3p electron in S than unpaired 3p in P due to presence of inter-electronic repulsion between electrons in the same orbital
- less energy required to remove paired 3p electron in S
How does first IE change between periods ?
Huge dip in first IE after each period due to extra shell of inner electrons in atom of next period
How does first IE change down a group ?
Down a group
- NC increases as number of protons increase
- SE increases as number of inner shell electrons increase
- ENC remains constant
- distance between nucleus and valence electrons increase due to increase in number of electron shells, weaker attraction between nucleus and valance electrons thus less energy required to remove valence electrons, thus first IE DECREASES
How does reactivity of group 1 metals change down the group ?
(Same as first IE)
Down group
- NC increases as number of protons increase
- SE increases as number of inner shell electrons increase
- ENC remains unchanged
- distance between valence electrons and nucleus increases due to an increase in number of electron shells, resulting in weaker attraction between valence electrons and nucleus thus less energy is required to remove the valence electrons
- ease of losing electrons increases, thus reducing strength/reactivity INCREASES
How does electronegativity change across a period ?
Across period
- NC increases as number of protons increase
- SE constant as number of inner shell electrons same
- ENC increases leading to stronger attraction between nucleus and electrons so the ability of the atom to attract electrons to itself increases, hence electronegativity INCREASES
How does electronegativity change down a group ?
Down group
- NC increases as number of protons increase
- SE increases as number of inner shell electrons increase
- ENC constant as increase in NC and increase in SE cancel out
- increase in number of electron shells increases the distance between nucleus and valence electrons, resulting in weaker attraction between the nucleus and valence electrons, so ability of atom to attract electron to itself decreases and electronegativity DECREASES
How does MP change across third period ?
Na, Mg, Al
- giant metallic structures and HIGH mp
- melting involves breaking of strong metallic bonds between the metal cation and sea of delocalised electrons which requires a large amount of energy
- metallic bonds become stronger across the period due to increasing charge density of cations and increase in number of delocalised electrons, mp INCREASES
Si
- giant molecular structure made up of Si atoms held by an extensive network of strong covalent bonds
- large amounts of energy required to break the numerous strong covalent bonds, thus Si has VERY HIGH mp
P, S, Cl, Ar
- non-metals with simple molecular structure
- molecules are held by weak id-id interaction, only a small amount of energy required to overcome the weak id-id interactions thus LOW mp
- S has greatest number of electrons, electron cloud is the most polarisable thus id-id between S8 molecules are the strongest thus require greatest amount of energy to overcome
- mp DECREASES across period as number of electrons decreases and electron cloud becomes less polarisable
How does electrical conductivity change across third period ?
Na, Mg, Al
- metals with delocalised sea of electrons which can act as mobile charge carriers hence have HIGH electrical conductivity
- number of delocalised electrons increase across the period hence electrical conductivity INCREASES
Si
- semi-metal which behaves as a semiconductor hence has LOW conductivity between conductor and insulator
P, S, Cl, Ar
- non-metals with simple molecular structure
- do not have mobile charge carriers and are hence electrical insulators
How does bonding of third period oxides change across the period ?
Changes from ionic to covalent
- from Na to Al, charge density increases leading to increase in polarising power, the electron cloud of oxide experience increased distortion thus covalent character increases
How does MP change across oxides of third period ?
Na2O, MgO, Al2O3,
- giant ionic structures, HIGH mp
- a lot of energy is needed to overcome the strong ionic bonds during melting
- Mg2+ has higher charge and smaller radius than Na+, more exo LE thus stronger ionic bonds
- Al2O3 has lower mp due to mix of covalent and ionic character
SiO2
- giant molecular structure, HIGH mp (lower than Al)
- extensive network of strong covalent bonds between Si and O, a lot of energy is required to break the strong covalent bonds
P4O10, SO3, Cl2O
- simple molecular structures, LOW mp
- both P4O10 and SO3 are non-polar, small amount of energy needed to overcome id-id interaction between the molecules, mp DECREASES as number of electrons decrease which leads to weaker id-id
Structure and nature of Na2O
Reaction with water
pH
Giant ionic structure, basic oxide
Dissolves in water : Na2O (s) + H2O (l) —> 2 NaOH (aq)
pH = 13
Structure and nature of MgO
Reaction with water
pH
Giant ionic, basic oxide
Highly exo lattice energy, hard to dissolve in water, reaction is reversible :
MgO (s) + H2O (l) <==> Mg(OH)2 (aq)
pH = 9
Structure and nature of Al2O3
Reaction with water
pH
Giant ionic, amphoteric oxide
Highly exo lattice energy, strong ionic bonds with covalent characters thus insoluble in water
pH = 7
Structure and nature of SiO2
Reaction with water
pH
Giant molecular, acidic oxide
Held by numerous stronger covalent bonds making it insoluble in water
pH =7
Structure and nature of P4O10
Reaction with water
pH
Simple molecular, acidic oxide
Readily dissolves in water : P4O10 (s) + 6 H2O (l) —> 4 H3PO4 (aq)
pH = 2
Structure and nature of SO3
Reaction with water
pH
Simple molecular, acidic oxide
Reacts violently with water : SO3(g) + H2O (l) —> H2SO4 (aq)
pH = 2
