Periodic Properties - QUICK notes for IITJEE and AIEEE

19 Mar 2009

An element’s location on the periodic table reflects the quantum numbers of the last orbital filled
The period indicates the value of principal quantum number for the valence shell
- the lanthanides and actinides are in periods 6 and 7, respectively.
The block indicates value of azimuthal quantum number () for the last subshell that received electrons in building up the electron configuration.
- blocks are named for subshells (s, p, d, f)
- Each block contains a number of columns equal to the number of electrons that can occupy that subshell
  - The s-block (in orange) has 2 columns, because a maximum of 2 electrons can occupy the single orbital in an s-subshell.
  - The p-block (in purple) has 6 columns, because a maximum of 6 electrons can occupy the three orbitals in a p-subshell.
  - The d-block (in green) has 10 columns, because a maximum of 10 electrons can occupy the five orbitals in a d-subshell.
  - The f-block (in dark blue) has 14 columns, because a maximum of 14 electrons can occupy the seven orbitals in a f-subshell.
questions to ponder
- What would the periodic table look like in a hypothetical universe where:
  - there were 3 possible values of m_s, instead of 2?
  - the angular momentum quantum number could take on values from 1 to n-1 only?
  - values of m = 0 were not allowed?
  - the maximum value of n were 5?

Factors affecting the valence shell. Anything that influences the valence electrons will affect the chemistry of the element.
	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)

what does atomic radius really mean?

Explaining periodic trends in atomic radius. See the section on factors affecting the valence shell above.
trend	valence n	Z	# core electrons	net 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.

using the trends
- to compare atoms in different groups and different periods, look for atoms that must be intermediate in size
  - this isn’t always possible!
  - example: Which is larger, a silicon atom, or a selenium atom?

periodic trends parallel those of atomic radius
cations are always smaller than the parent atom
- removing an electron decreases electron-electron repulsion, so the electron clouds contract
- emptying the valence shell completely leaves only electrons with lower n value
anions are always larger than the parent atom
- adding an electron to an atom increases electron-electron repulsion and swells the electron cloud
comparing radii for isoelectronic ions and atoms
- size within isoelectronic series is affected only by Z
- example
  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 is the minimum amount of energy required to remove an electron from an atom or ion in the gas phase

normally, ionization removes valence electrons first

valence electrons are farthest from nucleus on average, so they feel the least attraction for the nucleus and are easiest to remove

end of valence electrons is marked by a big jump in ionization energies

Na(g) Na⁺(g) + e^-	H = +496 kJ	first ionization energy
Na⁺(g) Na⁺²(g) + e^-	H = +4560 kJ	second ionization energy

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

Periodic Properties - QUICK notes for IITJEE and AIEEE

The properties of the alkali metals ultimately result from their ns¹ valence configuration.
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.