1. Atomic Structure
1.1. Subatomic Particles
● Protons (p+): +1 charge, 1 atomic mass unit (a.m.u.)
○ Atomic number (Z) = # protons
● Neutrons (n0): 0 charge, 1 a.m.u.
○ Mass number (A) = # protons + # neutrons
● Electrons (e–): –1 charge, negligible mass
○ Electron shells: closer to nucleus = lower energy, stronger electrostatic
interaction w/ nucleus
○ Valence electrons: involved in bonding
○ Cations (+), anions (–)
1.2. Atomic Mass, Weight
● Atomic mass ≈ mass number
○ Isotopes: same element w/ different mass numbers (e.g., carbon-12, 126C)
● Atomic weight: weighted average of all naturally occurring isotopes of an element
○ Avogadro’s number: NA = 6.02 × 1023
○ Mass of 1 mol atom ≈ atomic weight
1.3. Bohr Model
● Planck relation: E = hf = hc/λ
○ Planck’s constant: h = 6.626 × 10–34 J·s
● Bohr model: e–’s are in discrete shells around nucleus
○ Angular momentum of e– orbiting H+: L = nh / 2π
○ Energy of e– in H: E = –(RH) / n2
■ Rydberg unit of energy: RH = 2.18 × 10–18 J
● Atomic emission spectrum: e– moving to lower energy level emits energy (light)
○ E = hc/λ = RH (1/ni2 – 1/nf2)
■ Lyman series: to n = 1
■ Balmer series: to n = 2
■ Paschen series: to n = 3
● Atomic absorption spectrum: e– moving to higher energy level absorbs energy
○ Complement of emission spectrum
1.4. Quantum-Mechanical Model
● Quantum-mechanical (QM) model: e–’s are delocalized in orbitals
○ Heisenberg uncertainty principle: σxσp ≥ h / 4π
● Quantum numbers
○ Pauli exclusion principle: 2 e– cannot share same quantum numbers
○ Principal (n): energy level (shell in Bohr model)
■ n≥1
■ Each shell fits 2n2 e–
○ Azimuthal (l): subshell
■ 0≤l≤n–1
■ Each subshell fits 4l + 2 e–
■ Spectroscopic notation: s (l = 0), p (l = 1), d (l = 2), f (l = 3)
○ Magnetic (ml): orbital
■ –l ≤ ml ≤ +l
○ Spin (ms): 2
, ■ ms = ±½
● Electron configuration
○ Aufbau principle: fill lower-energy orbitals first
■ 1s → 2s → 2p, 3s → 3p, 4s → 3d, 4p, 5s → …
■ Shorthand w/ periodic table: start counting from last noble gas
○ n + l rule: higher n + l = higher energy when comparing 2 orbitals, n is tiebreaker
○ Anion: add e–’s to configuration
○ Cation: remove e–’s from highest n first, then highest l
● Hund’s rule: half-fill orbitals first, w/ unpaired e– of parallel spin
○ Half-filled, full orbitals are much more stable
■ Cr is [Ar] 4s1 3d5, not [Ar] 4s2 3d4
■ Cu is [Ar] 4s1 3d10, not [Ar] 4s2 3d9
○ Paramagnetism: atoms w/ many unpaired e– (parallel spins) are attracted to
magnetic fields
○ Diamagnetism: atoms w/ only paired e– are repelled by magnetic fields
● Valence electrons: farthest from nucleus, greatest PE, participate in reactions
○ Highest s, p in valence shell
○ Transition metals: highest s, d subshells
○ Lanthanides/actinides: highest s, f subshells
○ ≥ Na elements can accept e– into d valence shell, violate octet rule
2. Periodic Table
2.1. Periodic Table
● Periods (rows), groups/families (columns)
○ Elements in same group have same e– configuration in valence shell
● A (representative) elements: valence e – in s, p
○ Group # = # valence e–
● B (nonrepresentative) elements: valence e– in s/d, s/f
○ Transition metals, lanthanides, actinides
2.2. Element Types
● Metals: left side of periodic table
○ Lustrous, malleable, ductile, solid (except Hg), high melting point/density (except
Li)
○ High electropositivity (ability to donate e–)
■ Large atomic radius, small ionic radius
■ Low ionization energy, low electron affinity, low Zeff
■ Good heat, electrical conductors
○ Transition metals: multiple oxidation states
● Nonmetals: right side of periodic table
○ No luster, brittle
○ High electronegativities (ability to attract e–)
■ Small atomic radius, large ionic radius
■ High ionization energy, high electron affinity
■ Poor heat, electrical conductors
● Metalloids/semimetals: B, Si, etc.
○ Intermediate electronegativities, ionization energies
○ Variable physical properties
, 2.3. Periodic Trends
● Effective nuclear charge (Zeff): net (+) charge felt by outermost e–
○ ↑ from left to right: more (+) nucleus, same valence shell
○ Constant w/in group: more (+) nucleus, more shielding
● Noble (inert) gases: full octet
● Atomic radius: size of neutral atom
○ ↓ from left to right: more (+) nucleus, same valence shell
○ ↑ from top to bottom: more shells
● Ionic radius: size of cation (metal), anion (nonmetal)
○ < atomic radius for metals (lose e–)
○ > atomic radius for nonmetals (gain e–)
○ ↓ from left to right: metals lose more e–, nonmetals gain fewer e–
● Ionization energy (IE)/ionization potential: energy needed to remove e– from gaseous
atom
○ ↑ from left to right: harder to remove e– from higher Zeff
○ ↓ from top to bottom: easier to remove e– farther from nucleus
○ 1st IE < 2nd IE < …
○ Active metals (Groups 1–2): losing 1–2 e– creates filled valence shell (stable)
○ Very high for noble gases
● Electron affinity (EA): energy released when adding e– to gaseous atom
○ ↑ from left to right: easier to add e– to higher Zeff
○ ↓ from top to bottom: less energetic to add e– farther from nucleus
○ Almost zero for noble gases
● Electronegativity (EN): ability to attract bonding e–
○ Same trend as IE, except ≈ 0 for noble gases
● Periodic trends
○ Left to right: atomic radius ↓, IE ↑, EA ↑, EN ↑, Zeff ↑
○ Top to bottom: atomic radius ↑, IE ↓, EA ↓, EN ↓
2.4. Groups
● Alkali metals (1, IA): active metals, low density, 1 valence e–, low Zeff
● Alkaline earth metals (2, IIA): active metals, similar to alkali metals, 2 valence e–
● Transition metals (3–12, B): metals, low IE, multiple oxidation states
● Pnictogens (15, VA): nonmetals/metalloids/metal, 5 valence e–
● Chalcogens (16, VIA): nonmetals/metalloids, 6 valence e–
● Halogens (17, VIIA): highly reactive nonmetals, 7 valence e–, diatomic or bond w/ active
metals
● Noble gases (18, VIIIA): gases, unreactive, high IE
3. Bonding, Interactions
3.1. Bonding
● Molecules: atoms joined by covalent bonds (valence e– interactions)
● Octet rule: 8 e– in valence shell = stable noble gas configuration
○ Incomplete octet: H, He, Li (2), Be (4), B (6)
○ Expanded octet: period 3+ elements
○ Odd # e– in molecule
1.1. Subatomic Particles
● Protons (p+): +1 charge, 1 atomic mass unit (a.m.u.)
○ Atomic number (Z) = # protons
● Neutrons (n0): 0 charge, 1 a.m.u.
○ Mass number (A) = # protons + # neutrons
● Electrons (e–): –1 charge, negligible mass
○ Electron shells: closer to nucleus = lower energy, stronger electrostatic
interaction w/ nucleus
○ Valence electrons: involved in bonding
○ Cations (+), anions (–)
1.2. Atomic Mass, Weight
● Atomic mass ≈ mass number
○ Isotopes: same element w/ different mass numbers (e.g., carbon-12, 126C)
● Atomic weight: weighted average of all naturally occurring isotopes of an element
○ Avogadro’s number: NA = 6.02 × 1023
○ Mass of 1 mol atom ≈ atomic weight
1.3. Bohr Model
● Planck relation: E = hf = hc/λ
○ Planck’s constant: h = 6.626 × 10–34 J·s
● Bohr model: e–’s are in discrete shells around nucleus
○ Angular momentum of e– orbiting H+: L = nh / 2π
○ Energy of e– in H: E = –(RH) / n2
■ Rydberg unit of energy: RH = 2.18 × 10–18 J
● Atomic emission spectrum: e– moving to lower energy level emits energy (light)
○ E = hc/λ = RH (1/ni2 – 1/nf2)
■ Lyman series: to n = 1
■ Balmer series: to n = 2
■ Paschen series: to n = 3
● Atomic absorption spectrum: e– moving to higher energy level absorbs energy
○ Complement of emission spectrum
1.4. Quantum-Mechanical Model
● Quantum-mechanical (QM) model: e–’s are delocalized in orbitals
○ Heisenberg uncertainty principle: σxσp ≥ h / 4π
● Quantum numbers
○ Pauli exclusion principle: 2 e– cannot share same quantum numbers
○ Principal (n): energy level (shell in Bohr model)
■ n≥1
■ Each shell fits 2n2 e–
○ Azimuthal (l): subshell
■ 0≤l≤n–1
■ Each subshell fits 4l + 2 e–
■ Spectroscopic notation: s (l = 0), p (l = 1), d (l = 2), f (l = 3)
○ Magnetic (ml): orbital
■ –l ≤ ml ≤ +l
○ Spin (ms): 2
, ■ ms = ±½
● Electron configuration
○ Aufbau principle: fill lower-energy orbitals first
■ 1s → 2s → 2p, 3s → 3p, 4s → 3d, 4p, 5s → …
■ Shorthand w/ periodic table: start counting from last noble gas
○ n + l rule: higher n + l = higher energy when comparing 2 orbitals, n is tiebreaker
○ Anion: add e–’s to configuration
○ Cation: remove e–’s from highest n first, then highest l
● Hund’s rule: half-fill orbitals first, w/ unpaired e– of parallel spin
○ Half-filled, full orbitals are much more stable
■ Cr is [Ar] 4s1 3d5, not [Ar] 4s2 3d4
■ Cu is [Ar] 4s1 3d10, not [Ar] 4s2 3d9
○ Paramagnetism: atoms w/ many unpaired e– (parallel spins) are attracted to
magnetic fields
○ Diamagnetism: atoms w/ only paired e– are repelled by magnetic fields
● Valence electrons: farthest from nucleus, greatest PE, participate in reactions
○ Highest s, p in valence shell
○ Transition metals: highest s, d subshells
○ Lanthanides/actinides: highest s, f subshells
○ ≥ Na elements can accept e– into d valence shell, violate octet rule
2. Periodic Table
2.1. Periodic Table
● Periods (rows), groups/families (columns)
○ Elements in same group have same e– configuration in valence shell
● A (representative) elements: valence e – in s, p
○ Group # = # valence e–
● B (nonrepresentative) elements: valence e– in s/d, s/f
○ Transition metals, lanthanides, actinides
2.2. Element Types
● Metals: left side of periodic table
○ Lustrous, malleable, ductile, solid (except Hg), high melting point/density (except
Li)
○ High electropositivity (ability to donate e–)
■ Large atomic radius, small ionic radius
■ Low ionization energy, low electron affinity, low Zeff
■ Good heat, electrical conductors
○ Transition metals: multiple oxidation states
● Nonmetals: right side of periodic table
○ No luster, brittle
○ High electronegativities (ability to attract e–)
■ Small atomic radius, large ionic radius
■ High ionization energy, high electron affinity
■ Poor heat, electrical conductors
● Metalloids/semimetals: B, Si, etc.
○ Intermediate electronegativities, ionization energies
○ Variable physical properties
, 2.3. Periodic Trends
● Effective nuclear charge (Zeff): net (+) charge felt by outermost e–
○ ↑ from left to right: more (+) nucleus, same valence shell
○ Constant w/in group: more (+) nucleus, more shielding
● Noble (inert) gases: full octet
● Atomic radius: size of neutral atom
○ ↓ from left to right: more (+) nucleus, same valence shell
○ ↑ from top to bottom: more shells
● Ionic radius: size of cation (metal), anion (nonmetal)
○ < atomic radius for metals (lose e–)
○ > atomic radius for nonmetals (gain e–)
○ ↓ from left to right: metals lose more e–, nonmetals gain fewer e–
● Ionization energy (IE)/ionization potential: energy needed to remove e– from gaseous
atom
○ ↑ from left to right: harder to remove e– from higher Zeff
○ ↓ from top to bottom: easier to remove e– farther from nucleus
○ 1st IE < 2nd IE < …
○ Active metals (Groups 1–2): losing 1–2 e– creates filled valence shell (stable)
○ Very high for noble gases
● Electron affinity (EA): energy released when adding e– to gaseous atom
○ ↑ from left to right: easier to add e– to higher Zeff
○ ↓ from top to bottom: less energetic to add e– farther from nucleus
○ Almost zero for noble gases
● Electronegativity (EN): ability to attract bonding e–
○ Same trend as IE, except ≈ 0 for noble gases
● Periodic trends
○ Left to right: atomic radius ↓, IE ↑, EA ↑, EN ↑, Zeff ↑
○ Top to bottom: atomic radius ↑, IE ↓, EA ↓, EN ↓
2.4. Groups
● Alkali metals (1, IA): active metals, low density, 1 valence e–, low Zeff
● Alkaline earth metals (2, IIA): active metals, similar to alkali metals, 2 valence e–
● Transition metals (3–12, B): metals, low IE, multiple oxidation states
● Pnictogens (15, VA): nonmetals/metalloids/metal, 5 valence e–
● Chalcogens (16, VIA): nonmetals/metalloids, 6 valence e–
● Halogens (17, VIIA): highly reactive nonmetals, 7 valence e–, diatomic or bond w/ active
metals
● Noble gases (18, VIIIA): gases, unreactive, high IE
3. Bonding, Interactions
3.1. Bonding
● Molecules: atoms joined by covalent bonds (valence e– interactions)
● Octet rule: 8 e– in valence shell = stable noble gas configuration
○ Incomplete octet: H, He, Li (2), Be (4), B (6)
○ Expanded octet: period 3+ elements
○ Odd # e– in molecule
