Periodic Trends

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Atomic Radii

1) As you move down a group, atomic radius increases.

WHY? - The number of energy levels increases as you move down a group as the number of electrons increases.  Each subsequent energy level is further from the nucleus than the last.  Therefore, the atomic radius increases as the group and energy levels increase.

2) As you move across a period, atomic radius decreases.

WHY? - As you go across a period, electrons are added to the same energy level.  At the same time, protons are being added to the nucleus.  The concentration of more protons in the nucleus creates a "higher effective nuclear charge."  In other words, there is a stronger force of attraction pulling the electrons closer to the nucleus resulting in a smaller atomic radius.

 

Ionic Radii

1) Anions (negative ions) are larger than their respective atoms.

WHY?

Electron-electron repulsion forces them to spread further apart.
Electrons outnumber protons; the protons cannot pull the extra electrons as tightly toward the nucleus.

2) Cations (positive ions) are smaller than their respective atoms.

WHY?

There is less electron-electron repulsion, so they can come closer together.
Protons outnumber electrons; the protons can pull the fewer electrons toward the nucleus more tightly.
If the electron that is lost is the only valence electron so that the electron configuration of the cation is like that of a noble gas, then an entire energy level is lost.  In this case, the radius of the cation is much smaller than its respective atom.

First Ionization Energy

Definition:  The energy required to remove the outermost (highest energy) electron from a neutral atom in its ground state.

1) As you move down a group, first ionization energy decreases.

WHY?

Electrons are further from the nucleus and thus easier to remove the outermost one.
"SHIELDING" - Inner electrons at lower energy levels essentially block the protons' force of attraction toward the nucleus.  It therefore becomes easier to remove the outer electron

2) As you move across a period, first ionization energy increases.

WHY? - As you move across a period, the atomic radius decreases, that is, the atom is smaller.  The outer electrons are closer to the nucleus and more strongly attracted to the center.  Therefore, it becomes more difficult to remove the outermost electron.

Exceptions to First Ionization Energy Trends

1) Xs2 > Xp1  e.g. 4Be > 5B WHY? - The energy of an electron in an Xp orbital is greater than the energy of an electron in its respective Xs orbital.  Therefore, it requires less energy to remove the first electron in a p orbital than it is to remove one from a filled s orbital.  2) Xp3 > Xp4  e.g.  7N > 8O WHY? - After the separate degenerate orbitals have been filled with single electrons, the fourth electron must be paired.  The electron-electron repulsion makes it easier to remove the outermost, paired electron. (See Hund's Rule)

Veiw a periodic table with first ionization energies.

Second and Higher Ionization Energies

Definition:  Second Ionization Energy is the energy required to remove a second outermost electron from a ground state atom.
Subsequent ionization energies increase greatly once an ion has reached the state like that of a noble gas.  In other words, it becomes extremely difficult to remove an electron from an atom once it loses enough electrons to lose an entire energy level so that its valence shell is filled.

Ionization Energies (kJ/mol)
Element Na Mg Al		1st IE 495.8 737.7 577.6		2nd IE 4562.4 1450.6 1816.6		3rd IE 6912 7732.6 2744.7		4th IE 9543 10,540 11,577

Electron Affinity

Definition:  The energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion.

1) As you move down a group, electron affinity decreases.

2) As you move across a period, electron affinity increases.

	Exceptions Among nonmetals, however, the elements in the first period have lower electron affinities than the elements below them in their respective groups.  Elements with electron configurations of Xs2, Xp3, and Xp6 have electron affinities less than zero because they are unusually stable.  In other words instead of energy being given off, these elements actually require an input of energy in order to gain electrons.  e.g.  Be, N, Ne  WHY? - Electron affinities are all much smaller than ionization energies.  Xs2 < 0:  Stable, diamagnetic atom with no unpaired electrons.  Xp3 < 0: Stable atom with 3 unpaired p-orbital electrons each occupying its own subshell.  Xp6 < 0: Stable atom with filled valence (outermost) shell.

Lattice Energy

Definition:  The energy given off when oppositely charged ions in the gas phase come together to form a solid.

The strength of a bond between ions of opposite charge can be calculated using Coulomb's Law.

Coulomb's Law - The force of attraction between oppositely charged particles is directly proportional to the product of the charges of the particles (q₁ and q₂) and inversely proportional to the square of the distance between the particles.

1) As you move down a group, lattice energy decreases.

WHY? - The atomic radius increases as you move down a group.  Since the square of the distance is inversely proportional to the force of attraction, lattice energy decreases as the atomic radius increases.

2) As you increase the magnitude of the charge (becomes more positive or more negative), lattice energy increases.

WHY? - The product of the charges of the particles is directly proportional to the force of attraction. Therefore, lattice energy increases as the charges increase.

Lattice Energies of Alkali Metals with Halides (kJ/mol) Li+ Na+ K+ Rb+ Cs+ F- 1036 923 821 785 740 Cl- 853 787 715 689 659 Br- 807 747 682 660 631 I- 757 704 649 630 604

Lattice Energies of Salts of OH- and O2- with Cations of varying charge (kJ/mol) Na+ Mg2+ Al3+ OH- 900 3006 5627 O2- 2481 3791 15916