Periodic Properties - QUICK notes for IITJEE and AIEEE
Periodic Properties - QUICK notes for IITJEE and AIEEE Factors affecting the valence shell. Anything that influences the valence electrons will affect the chemistry of the element. Explaining periodic trends in atomic radius. See the section on factors affecting the valence shell above.
Quantum numbers and the periodic table for IITJEE and AIEEE
indicates the value of principal quantum number
for the valence shell
and actinides
are in periods 6 and 7, respectively.
indicates value of azimuthal quantum number
(
) for the last subshell
that received electrons in building up the electron configuration.
= 0 were not allowed?
Factors affecting the valence shell
Factors (in order of decreasing importance)
Effect
1.
valence principal quantum number n
Larger n means a larger valence shell (because n controls the size of orbitals)
2.
nuclear charge Z
Larger Z means a smaller valence shell (because higher positive charge on the nucleus attracts the valence electrons, and pulls them inward)
3.
number of core electrons
More core electrons means a larger valence shell (because highly penetrating core electrons repel valence electrons, and push them farther from the nucleus)
Atomic radius
trend
valence
nZ
# core
electronsnet effect on atomic radius
going right across main group rows…
no change
increases
no change
the increase in Z causes a decrease in radius
going right across transition series…
no change
increases
increases
the increase in Z causes a decrease in radius, but the increase in the number of core electrons causes an increase. The two competing effects cause a small decrease, then small increase!
going down groups…
increases
increases
increases
three competing effects; but n is strongest, so radius increases.
Ionic radius
ions and atoms
F-, Ne, and Na+ are isoelectronic, with Z = 9, 10, and 11, respectively. All have identical valence n and identical numbers of electrons, so the larger Z is, the smaller the atom or ion. Na+ is the smallest and F- is the largest.Ionization energy
Na(g)
Na+(g) + e-
H = +496 kJ
first ionization energy
Na+(g)
Na+2(g) + e-
H = +4560 kJ
second ionization energy
- atomic radius
- smaller atoms hang on to valence electrons more tightly, and so have higher ionization energy
- charge
- the higher the positive charge becomes, the harder it is to pull away additional electrons
- second ionization energy is always higher than the first
- orbital penetration
- It’s easier to remove electrons from p orbitals than from s orbitals
- electron pairing
- within a subshell, paired electrons are easier to remove than unpaired ones
- reason: repulsion between electrons in the same orbital is higher than repulsion between electrons in different orbitals
- example
On the basis of gross periodic trends, one might expect O to have a higher ionization energy than N. However, the ionization energy of N is 1402 kJ/mol and the ionization energy of O is only 1314 kJ/mol. Explain.
Taking away an electron from O is much easier, because the O contains a paired electron in its valence shell which is repelled by its partner.
Why metals are metals
- the ionization energy of metallic elements is very low
- valence electrons are easily lost, and shared among all atoms in the metal
- this ‘sea’ of valence electrons binds together the metal cations and gives metals their characteristic properties
- mobility of electrons in the sea explains metal’s ability to conduct electricity and heat
- metals are workable because cations can slide past each other but still be bound by the electron sea
- comparing metals
- more valence electrons means stronger metal
- higher positive charge on cations, higher negative charge on sea = stronger bonding
Explaining elemental properties: the s block elements
Periodic Properties - QUICK notes for IITJEE and AIEEE
property of alkali metals | explanation |
metallic |
very low ionization energy; the electron sea model works well for alkali metals |
soft |
ns1 valence configuration contributes just 1 electron to the electron sea. The sea is weak. Metal cations aren’t tightly bound and it’s easy to slide them past each other. |
low densities |
Alkali metals have the largest radii and lowest atomic weight in each period. Low mass in high volume = low density. |
highly reactive |
very low ionization energies make alkali metals good electron donors in redox reactions. |
- the alkaline earth metals (Group IIA)
- soft, but harder than alkali metals
- ns2 valence configuration = more electrons in the sea = more tightly bound metal cations
- reactive, but not as reactive as alkali metals
- ionization energies are not as low as alkali metals
- salts less soluble than those of the alkali metals
- higher cation charge concentrated on smaller cations makes it hard to pull apart ionic lattices
- soft, but harder than alkali metals