chem 1

Electron Shells and Orbitals
- Each electron in an atom is assigned quantum numbers characterizing its state.
- Principal quantum number (n): Indicates the shell level (1, 2, 3,…).
- Azimuthal quantum number (l): Indicates the subshell type (s, p, d, f). The values are:
  - s: l = 0
  - p: l = 1
  - d: l = 2
  - f: l = 3
- Magnetic quantum number (m_l): Can range from -l to l.
- For example, for d (l = 2): values range from -2 to +2.
- Sequence for m_l values for d: -2, -1, 0, +1, +2.
- Spin quantum number (m_s): Either +1/2 or -1/2 for electron spin orientation.

All electron addresses in an atom create its overall electron configuration.
Current knowledge includes 118 different elements.
Simplified Notation:
- Starts with the previous noble gas to represent inner electrons.
- Example: Cobalt (Co), located in period 4.
- Argon (Ar) is the previous noble gas.
- Electron configuration of Co: [Ar] 4s² 3d⁷.
Atoms typically end in p⁶ of their period, except helium.

Follow the periodic table to keep track of electron filling order.
Order of filling:
- 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
S, P, D, and F designations:
- S starts in row n, P starts in row n-1 (exceptions for row 1), D starts in n-1, and F starts in n-2.

Periodic Table Design:
- Created based on atomic mass and properties long before electron behavior was understood, allowing strong prediction of atomic behavior based on electron configuration.
- The periodic table positively correlates electron behavior with element properties.
Valence Electrons:
- Tied to bonding behavior.
- Main/Representative elements: main group valence electrons = s and p of their period.
- Example: Period 3 elements involve 3s and 3p.
- D-transition metals: electrons used in bonding are from the outermost s and one behind d.
- F-block involves s electrons in the same period and d minus 2f.

Atomic Radius:
- Distance from the nucleus to the valence electrons.
Ionization Energy:
- Energy required to remove an electron from an atom.
- More tightly the atom holds onto electrons, larger energy required.
- Deduction: Increases left to right in a period, decreases top to bottom in a group.
Electron Affinity:
- Energy released when an atom gains an electron. More negative value indicates a stronger affinity.
- Trend: Increases (becomes more negative) left to right, decreases down a group.
Ionic Radius:
- Distance from the nucleus to the outer shell in ions.
- Trends: Cations (positively charged) < neutral atoms < anions (negatively charged).

The concept that explains how tightly electrons are held by the nucleus.
Z_eff = Z - S
- Where Z = atomic number (number of protons).
- S = shielding electrons; calculated as:
  $S = 0.85 imes ext{(inner electrons)} + 0.35 imes ext{(valence electrons - 1)}$
Shielding Effect:
- Inner electrons partially shield outer electrons from the nuclear charge.
- Example: Lithium has 3 protons and 2 inner electrons, shielding the outer electron.
- Zeff calculation for Lithium: 3 - (0.85)2 = 1.3.

Ionization Energy:
- Increases across a period (more protons, more Zeff), decreases down a group (more shielding).
Electron Affinity:
- More negative moving to the right, less negative moving down.
Ionic Radius Trends:
- Cation radii shrink as electrons are removed.
- Anion radii grow with additional electrons.
Electronegativity:
- Tendency to attract electrons; follows similar trends to electron affinity.

Not all trends follow expected patterns; unique configurations can lead to lower energy requirements in certain transitions.
- Example: Group 2 to Group 3 (Beryllium to Boron) - stabilization from moving from full to half-filled orbitals can lead to lower ionization energy.
- Oxygen's ionization energy drops between groups 5 and 6 (half-filled vs. filled orbital dynamics).

Describes a characteristic reaction energy behavior when atoms accept electrons into vacant orbitals.
Greater Electronegativity:
- Indicates a higher negative electron affinity.
- Group 8 elements: Electron affinities tend to be zero or positive due to stable full orbitals.

A measure of the tendency of an atom to attract a bonding pair of electrons.
General Trends:
- Increases across a period (stronger pull for bonding) but decreases down a group (hold on to electrons less tightly).

Transition elements behave irregularly in terms of electronegativity and ionic charges due to d-block electron configurations impacting Z_eff dynamically.
Each of these nuances is vital for understanding chemical bonding, periodic properties, and the fundamental behavior of atoms in chemical reactions.