Periodic Properties Flashcards
IUPAC nomenclature (100+)
0 - nil
1 - un
2 - bi
3 - tri
4 - quad
5 - pent
6 - hex
7 - sept
8 - oct
9 - enn
Nuclear Charge
Measure of ability of protons to attract negative electrons
Periods
Horizontal rows
There are 7
Groups
Vertical columns
There are 18
Which elements begins the 3d transition series
Scandium (Z = 21)
Which element begins the 4d transition series
Yttrium (Z = 39)
When does filling of 4f electrons begin
Z = 58 - Cerium
Properties of s - block elements
- ns1 or ns2 outer config
- Reactive metals with low ionization enthalpies
- Lose 1+ or 2+ ions
S block elements as we go down the group
- Metallic character increases
- Reactivity increases
Which s-block elements are not ionic
Lithium and Beryllium
P block elements outer config
ns2np1 and ns2np6
P block character down the group and across
- Non - metallic character increases left to right
- Metallic increases as we go down
D block elements outer config
(n-1)d (1-10) ns (0-2)
Covalent Radius
Size of an atom that is part of a single - covalent bond
Metallic Radius
Measurement of size of atom in relation to other metal elements
Atomic Radius
Measures size of the atom; mean / typical distance from the center of the nucleus to the outermost isolated electron
What happens to size across the period and why?
- Size usually decreases
- In the same outermost shell, added electrons do not screen each other.
- Hence, effective nuclear charge increases pulling electrons closer to the nucleus
Factors affecting atomic radii
- Increase in effective nuclear charge = Increase in attractive force
- Increase in ionic character = Shorter bond = Lower atomic radii
What happens to size down the group?
- Down the group the size increases
- The addition of the extra valence shell has a more pronounced effect on the size compared to the effect of the increase in nuclear charge
Why is anion > parent > cation
- Anion is larger as addition of one or more electrons leads to increased repulsion among electrons and decrease in effective nuclear charge
- Cation is smaller as it has fewer electrons while nuclear charge is the same
Isoelectronic species
Atoms and ions which contain the same number of electrons
Ionization enthalpy
Energy required to remove an electron from an isolated gaseous atom in ground state
Factors affecting ionization enthalpy
- Larger size of atom = smaller ionization enthalpy
- Higher screening effect = lower value of ionization enthalpy
- Higher nuclear charge = Higher ionization enthalpy
- Penetration effect - s > p > d > f
Why do alkali metals have low ionization enthalpy
They have high reactivity
Trend of ionization enthalpy and explanation
- Across period (it increases) - electrons are added in the same principal quantum number –> shielding does not increase that much to compensate for increased attraction
- Down the group (decreases) - electron added further from nucleus causing shielding to outweigh nuclear charge and removal of electron requires less energy
Exceptions in trend in ionization enthalpy
- Boron < Berrylium - Boron electron removed from 2s where 2s penetration > 2p
- Oxygen < Nitrogen - Nitrogen has 2p3 stability and due to hund’s rule and symmetry it requires more energy to separate (Same reason for S < P)
- Al < Mg
Explain the deviation in ionization enthalpy with respect to
B, Al, Ga, In, Ti
- Decrease in B to Al is due to larger size of Al
- In Ga, as there are more 3d electrons, there is less screening causing an increase in effective nuclear charge
- Ti has very poor shielding as well (hence greater than In)
Electron Gain Enthalpy
When an electron is added to a neutral gaseous atom to convert it into a negative ion
Electron Gain Enthalpy across period
- Becomes more negative - Effective nuclear charge increases and harder to add an atom to a larger atom
Electron Gain Enthalpy down the group
Becomes less negative as atom size increases and added electron is further from nucleus
Exceptions of EA (Electron affinity / electron gain enthalpy)
- Cl > F –> Cl has a larger size, meaning less force of repulsion between electrons
- In group 16 –> S > Se > Te > Po > o
- Noble gases have positive EA
Factors affecting EA
- Higher atomic size = Lower EA
- Higher effective nuclear charge = higher EA
- EA decreases with increasing shielding effect
- Half filled or completely filled orbitals
Why is EA positive for noble gases
- Completely filled outer shell (highly stable)
- To gain another electron, they reequire energy
- Energy required to overcome electrostatic repulsion between added electrons and existing electron
Electronegativity
Ability of an atom to attract shared electrons to itself
Trend of electronegativity
- Across the period - Increase in non-metallic character = increase in electronegativity
- Down the group - Increase in metallic character = decrease in electronegativity
On the basis of quantum numbers justify that the sixth period should have 32 elements
- The sixth period can have 6s, 6p, 5d, 4f as they are lesser energy than 7s
- Total number of orbitals found using (2l + 1) –> 1 + 3 + 5 + 7 = 16.
- Total electrons accommodated is multiplied by 2 => 32
Diagonal Relationship
Behaviour of beryllium and lithium similar to Mg and Al
Size of isoelectronic species is affected by
Nuclear charge
Why does neon have a larger radius than flourine?
- due to the way atomic radii are measured for noble gases like neon and covalent radii for non-noble gases like fluorine.
- Noble gases, including neon, have van der Waals radii, which are typically larger than covalent radii. Van der Waals radii measure the distance between the center of the atom and the outermost point of the electron cloud, whereas covalent radii measure the distance between the nucleus and the bonding electrons
