6.5 - Periodic Trends

Atomic Radii

• Half the distance between nuclei in a molecule of identical atoms (S8, N2)

• Increases down a group (vertical)

  • Larger n = larger size

• Decreases across a period (left to right)

As Z (atomic number/# of protons) increases…

• Electron-nucleus attraction increases

• Outer electrons are “shielded” by inner electrons as increasing the number of electrons causes more e-e repulsion

Shielding and Penetration

• Shielding:The reduction of the attractive positive charge of the nucleus (called the nuclear charge or (Zeff) down to a smaller, net positive charge

  • Core electrons (inner shell electrons) sit closer to the positively charged nucleus and have the same negative charge as outer valence electrons. Because like charges repel, these inner electrons push outward on outer electrons and partially "block" or screen them from feeling the full pull of the nucleus.

• Penetration: The relative probability that an electron in a given subshell can approach very close to the nucleus.

  • Because of the shapes of different atomic orbitals, some electrons have a small amount of electron density very close to the nucleus (inside the path of inner shell electrons).

Why is a 1s orbital lower in energy than a 7s?

• 1s electrons are closer to the nucleus

• 1s electrons experience no shielding

• 7s electrons have lower energy electrons between them and the nucleus

In which orbital do electrons have a greater probability of being closer to the nucleus, a 2s or 2p electron?

• 2p, even though its higher in energy

Exceptions

• 2s electrons are usually further from the nucleus than 2p but 2s electrons penetrate, moving close to the nucleus

• 2S ORBITAL HAS LOWER ENERGY OVERALL THAN 2P BECAUSE IT PENETRATES (THE SMALL PEAK)

Effective Nuclear Charge (Zeff)

• The nuclear charge an electron feels considering shielding effects

  • core electrons electrons strongly shield valence electrons (allows valence electrons to have higher energies)

  • Valence electrons poorly shield each other

Zeff = Z-S

Z=nuclear charge, S=Number of lower energy electrons

How to calculate Zeff of every electron in an element

$$Z_{\text{eff}} = Z - (\text{Number of core electrons below that shell})$$

1. Group electrons by principle shell (n)

2. n = 1:1s2 (2 electrons)

   n = 2: 2s 2p6 (8 electrons) n = 3:35 3p6 (8 electrons)

   n = 4:4s2 (2 valence electrons)

There are no electrons inside the 1s shell (0 core electrons)

Radius of Ions Trend

• Ionic radii also increase going down a group

• Cations are smaller than its neutral atom (less e)

• Anions are bigger than its neutral atom (more e)

Isoelectric Series Trend

• Isoelectric = same number of electrons

  • Larger positive charge = smaller cation

  • Larger negative charge = larger anion

Ionization Energy

• The energy required to remove an electron from an atom or ion/measures how strongly an atom holds its electrons

  • Removing an electron is always endothermic

Ionization Energy Trend

• Increases up a group

  • Because valence electrons are closer to the nucleus (takes less energy to remove an electron farther from the nucleus)

• Increases to the right as Zeff increases

  • The larger an electron’s Zeff, the more energy it takes to remove it

Ionization Energy Trend Exceptions

• Which is easier to remove an electron from?

  • B vs Be

    • B

    • Be has its valence e in the 2s orbital while B has its outermost e in a 2p orbital

    • The 2p orbital is slightly higher in energy than the 2s orbital, meaning the 2p electron is further from the nucleus and less tightly bound.

    • The single 2p electron in boron is shielded by the inner 2s and 1s electrons, reducing the pull of the nucleus and making it easier to pull away.

  • N vs O

  • O

  • N has a stable, half filled 2p subshell with 3 unpaired e, O has one paired e in its 2p subshell

  • In oxygen, putting two electrons into the same 2p orbital creates electron-electron repulsion. This repulsion makes it easier to pull one of those paired electrons away,

Successive Ionization Energies

• Outer electrons get closer to the nucleus and are harder to remove.

• There is a smaller increase in energy for each

successive valence electron removal and a large increase for core electrons.

Electron Affinity (EA)

• The energy change when an atom gains an electron

• Usually exothermic, sometimes endothermic

• No strong EA trend

  • Noble gases have the most endothermic EA

  • Halogens have the most exothermic EA