Chem unit 3 electrons

PERIODIC TRENDS:

Oxidation number: A number assigned to an atom to represent how many electrons are lost/gained

Electrostatic Attraction: Attraction between charged particle - ELECTRONS LOVE PROTONS

Electron Shielding: inner electrons shield outer electrons - ELECTRONS HATE OTHER ELECTRONS

Atomic Radius: Size of an atom

DECREASE AS YOU GO ACROSS TOP OF PERIODS INCREASE AS YOU GO DOWN PERIODS

Electrostatic Attraction: The stronger the attraction between protons (positive nucleus) and electrons (negative charge), the smaller the atomic radius because electrons are pulled closer.

Electron Shielding: Inner electrons shield outer electrons from the full nuclear charge, reducing the pull on the valence electrons and increasing the atomic radius. This is why atomic size increases down a group in the periodic table.

Ionization Energy: Energy required to remove an electron

INCREASE AS YOU GO ACROSS TOP OF PERIODS DECREASE AS YOU GO DOWN PERIODS

Electrostatic Attraction: Stronger attraction means higher ionization energy because more energy is needed to overcome the pull of the nucleus and remove an electron.

Electron Shielding: More shielding decreases the effective nuclear charge felt by the outermost electron, making it easier to remove (lower ionization energy). This explains why ionization energy decreases down a group.

Electronegativity: Ability to attract electrons

INCREASE AS YOU GO ACROSS TOP OF PERIODS DECREASE AS YOU GO DOWN PERIODS

Electrostatic Attraction: A stronger attraction between nucleus and electrons means a higher ability to attract shared electrons (higher electronegativity).

Electron Shielding: More shielding weakens the nucleus’ ability to attract bonding electrons, decreasing electronegativity. This is why electronegativity decreases down a group but increases across a period (more protons with minimal additional shielding).

Ionic Radius: Size of an ion

CATIONS(METALS) ON LEFT ANIONS (NON METALS) ON RIGHT

Cations: Lose e- (positive) lose energy level, decrease e- shield, ionic radius decreases

Anions: Gains e- (negative) more e- shielding, ionic radius increases

Electrostatic Attraction: Cations (positively charged ions) have fewer electrons than their neutral atoms, leading to stronger electrostatic attraction and a smaller radius. Anions (negatively charged ions) gain electrons, increasing electron-electron repulsion and making the ionic radius larger.

Electron Shielding: More shielding in larger ions weakens nuclear attraction, contributing to a larger radius, especially for anions.

Why does atomic radius decrease across a period?

As you move across a period, more protons are added to the nucleus, increasing electrostatic attraction between the nucleus and electrons. This pulls electrons closer, making the atom smaller.

Why does ionization energy increase across a period?

Since atoms have more protons across a period, they have a stronger pull on their electrons. This makes it harder to remove an electron, so more energy is required (higher ionization energy).

Why does atomic radius increase down a group?

As you move down a group, atoms gain more energy levels (electron shells). Even though the nucleus has more protons, the extra electron shells increase electron shielding, weakening the pull on outer electrons and making the atom larger.

Why does ionization energy decrease down a group?

Moving down a group, atoms have more electron shielding, which weakens the nuclear pull on outer electrons. Since the nucleus holds electrons less tightly, less energy is required to remove an electron, so ionization energy decreases.

Why does electronegativity increase across a period?

Across a period, atoms gain more protons but don’t get much extra shielding. This stronger nuclear attraction pulls electrons in closer and makes atoms more eager to attract additional electrons in bonds, so electronegativity increases.

Why does electronegativity decrease down a group?

As you move down a group, more electron shielding weakens the pull of the nucleus on bonding electrons. Since the nucleus has less ability to attract shared electrons, electronegativity decreases.

ELECTRON CONFIGURATION:

Hund's rule: Electrons hate other electrons so they fill the orbitals first by themselves

Pauli Exclusion Principle: Electrons have opposite spin and repel each other

Aufbau Principle: E- fill lower energy levels first

D starts at 3d F starts at 4f

EMR AND ELECTRONS:

Moves energy level = excited

Spectrum: All forms of electromagnetic radiation

Ionizing radiation: Enough energy to rip an electron away; atom = unstable

Non-ionizing radiation: Not enough energy to rip electron away atom = stable