Atomic Structure and Physical Periodicity Flashcards
electron, proton, neutron charges and mass
electron: -1, 1/1836
proton: +1, 1
neutron: 0, 1
atomic orbital definition
an atomic orbital is defined as the region of space with 90% probability of finding an electron
relative atomic mass definition
the ratio of the average mass of one atom of an element to 1/12 of the mass of one atom of C-12
relative isotopic mass definition
the ratio of the mass of one atom of an isotope of an element to 1/12 of the mass of one atom of C-12
formula of the angle of deflection in an electric field
angle of deflection is proportional to the charge and inversely proportional to the mass
state the Aufbau principle
electrons in their ground states occupy orbitals in order of energy levels. The orbital with the lowest energy is always filled first.
state the hund’s rule of multiplicity
when filling subshells that contain more than one orbital with the same energy, each orbital must be singly occupied before electrons are paired
state the pauli exclusion principle
an orbital cannot hold more than two electrons and the two electrons sharing the same orbital must have opposite spins.
shape of the s and d orbitals
sphere, dumb-bell
what is the schematic representation of the p and d orbitals
px, py, pz, d(x^2-y^2), dz^2, dxy, dxz, dyz
what are the two exceptions of the electronic configuration of elements
Cr: [Ar] 3d5 4s1
Cu: [Ar] 3d10 4s1
definition of transition metals
a transition metal is a d-block element that form one or more stable ions with a partially-filled d-subshell
Trend of atomic radius down the group
Atomic radius increases down the group. As the number of quantum shells increases, the outermost electrons are further away from the nucleus, hence the atomic radius increases.
why do we not use effective nuclear charge to explain atomic radius down the group?
Both the nuclear charge and shielding effect increase down the group, hence the effective nuclear charge differs little down the group.
Trend of atomic radius across periods 2 and 3
atomic radius decreases across the period. Nuclear charge increases due to the increase in the number of protons in the nucleus. Shielding effect remains relatively constant as the electrons are added to the same outermost shell. Effective nuclear charge increases, resulting in stronger electrostatic forces of attraction between the nucleus and the outermost electron. Outermost electrons are pulled closer to the nucleus and hence a decrease in the atomic radius.
why is the atomic radii of the noble gases so large
noble gases do not form compounds and do not have a covalent radius. The atomic radius of noble gases are measured through the van der waals’ radii, which is always larger than the covalent radius.
Trend of atomic radius across the first row transition elements
atomic radius is relatively invariant. Nuclear charge increases due to increasing number of protons. Electrons are added to the inner 3d subshell, which contributes to the shielding effect. Shielding effect increases thereby nullifying, to a considerable extent, the influence of each additional proton to the nucleus. Thus, effective nuclear charge remains almost constant.
Trend of ionic radius
Na+, Mg2+, Al3+, Si4+
P3-, S2-, Cl-
The cations are said to be isoelectronic as they have the same number of electrons. Similarly, anions are isoelectronic as they also have the same number of electrons. Across the two isoelectronic series, nuclear charge increases and shielding effect is the same due to the same number of electrons.
This results in an increase in the effective nuclear charge and a stronger attraction between the outermost electrons and the nucleus, hence a decrease in ionic size across each series.
There is a sharp increase in the ionic radius from the cationic series to the anionic series as the anions have one more quantum shell of electrons than the cations.
first ionisation energy definition
first ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of unipositively charged gaseous ions
second ionisation definition
second ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of unipositively charged gaseous ions to form 1 mole of gaseous ions with double positive charge
nth ionisation energy definition
nth ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of gaseous ions with (n-1)+ charge to form 1 mole of gaseous ions with n+ charge
factors affecting ionisation energy of an atom
number of quantum shells. the larger the number of quantum shells in an atom, the further the electron is from the nucleus and hence experiences weaker nuclear attraction and is easier to remove, hence the lower the ionisation energy.
effective nuclear charge. The higher the effective nuclear charge, the stronger the attractive forces between the nucleus and the electrons to be removed, hence the greater the ionisation energy.
why ionisation energies increase
IE 1 < IE 2 < IE 3…
when an electron is removed from a neutral atom, the number of protons that exert an attraction for the remaining electrons remains the same. however, the shielding effect among the remaining electrons in the outermost shell is reduced since there is now one less electron. Hence, effective nuclear charge increases. More energy is needed to remove another electron from the more positively charged ion, hence a higher ionisation energy.
trend of the ionisation energies down a group
first ionisation energy of elements decreases down a group. Down a group, the number of quantum shells of electrons increases, hence the outermost electrons are further from the nucleus. Therefore, electrostatic forces of attraction between the nucleus and the outermost electron is weaker and less energy is required to remove this electron.
Both the nuclear charge and shielding effect increase down the group, so the effective nuclear charge differs little down the group. Hence, the number of quantum shells is the more important factor.
trend of the ionisation energies across a period
across a period, nuclear charge increases and shielding effect remains relatively the same. Hence the effective nuclear charge increases and the electrostatic forces of attraction between the outermost electrons and the nucleus become stronger, so more energy is required to remove the outermost electron.
2 anomalies of ionisation energies
small dip between group 2 (Mg) and group 13 (Al) elements. The 3p subshell of Al is further away from the nucleus than the 3s subshell. There is weaker attraction between the nucleus and the outermost electron. Hence less energy is required to remove the 3p electron from Al, resulting in a lower ionisation energy for Al.
small dip between group 15 (P) and group 16 (S). All the 3p electrons in P are unpaired. In S, two of the 3p electrons are paired. There is some inter-electronic repulsion between the paired electrons in the 3p subshell in S. Thus, less energy than expected is required to remove one of these paired electrons from S.
trend of the ionisation energies across transition elements
first ionisation energies of transition elements remain relatively invariant. 1st ionisation energy involves the removal of a 4s electron. The nuclear charge increases due to increasing number of protons while additional electrons are added to the inner 3d subshell, which contributes to the shielding effect. The shielding effect increases thereby nullifying, to a considerable extent, the influence of each additional proton in the nucleus. Effective nuclear charge remains almost constant. Thus, energy required to remove the outermost electron of each succeeding element remains relatively invariant.
