Chemistry Lecture Notes Flashcards

electron clouds and energy levels

• electrons are in probability clouds around the nucleus.

• electrons have specific energy levels.

• simple model: electrons in rings around the nucleus.

• electrons fill the cloud in a specific order.

orbitals

• electrons are in orbitals within the electron cloud.

• an orbital is where an electron is likely to be.

• atoms have multiple orbitals.

• orbital shapes:

  • s orbital: spherical

  • p orbital: dumbbell-shaped

  • d orbital: cloverleaf-shaped

  • f orbital: complex shapes
    quantum numbers

• quantum numbers describe electron location.

• four main quantum numbers:

  • principal quantum number (n): orbital size. non-zero positive integer.

  • angular momentum quantum number (l): orbital shape. integer between 0 and n-1.

  • magnetic quantum number (m): orbital orientation in space. integer between -l and +l.

  • spin quantum number (m_s): electron spin. +1/2 or -1/2.

• analogy: like an electron's address.

• quantum numbers have specific integer values.

• if l = 1, m can be -1, 0, or 1.

• orbitals show probable electron locations, not exact paths.

shells

• electron shells: orbitals with the same n.

• shells fill consecutively, closest to the nucleus first (n=1).

• lowest energy: electrons in lowest energy shells.

• each shell holds a specific number of electrons:

  • n = 1 shell: 2 electrons

  • n = 2 shell: 8 electrons

  • n = 3 shell: 18 electrons

  • n = 4 shell: 32 electrons

• full shells are more stable.

• noble gases are unreactive due to full shells.

• shells are sometimes shown as circles with dots, but this is simplified.

subshells

• electron shells are divided into subshells.

• subshells: orbitals with the same values of n and l.

  • l = 0: s subshell

  • l = 1: p subshell

  • l = 2: d subshell

  • l = 3: f subshell

• values of m indicate the number of orbitals for each type:

  • s subshell: 1 orbital (m = 0)

  • p subshells: 3 orbitals (m = -1, 0, 1)

  • d subshells: 5 orbitals (m = -2, -1, 0, 1, 2)

  • f subshells: 7 orbitals (m = -3, -2, -1, 0, 1, 2, 3)

calculating maximum number of electrons

• example: if n = 3, then l can be 0, 1, or 2.

  • l = 0: 3s subshell

  • l = 1: 3p subshell

  • l = 2: 3d subshell

• determine values of m for each subshell:

  • l = 0: m = 0

  • l = 1: m = -1, 0, 1

  • l = 2: m = -2, -1, 0, 1, 2

• calculate the number of orbitals by counting combinations of m:

  • 3s: 1 orbital

  • 3p: 3 orbitals

  • 3d: 5 orbitals

• total orbitals: 1 + 3 + 5 = 9

• each orbital holds 2 electrons, so the total number of electrons is 9 \times 2 = 18.

• formula: 2n^2

  • for n=3, 2(3^2) = 2(9) = 18

electron configuration

• electron configuration: arrangement of electrons in an atom.

• orbital notation: lines and arrows show shells, subshells, and orbitals.

  • lines: orbitals.

  • numbers and letters: orbital name (e.g., 1s).

  • arrows: electrons.

pauli exclusion principle

• no two electrons can have the same set of quantum numbers.

• fourth quantum number: electron spin quantum number (m_s).

  • values: +1/2 (up arrow) or -1/2 (down arrow)

• an orbital holds two electrons with opposite spins.

filling orbitals

• lithium (3 electrons): 1s orbital (2 electrons), 2s orbital (1 electron).

• beryllium (4 electrons): 1s orbital (2 electrons), 2s orbital (2 electrons).

• electrons fill lowest energy states first (max 2 electrons per orbital).

• boron (5 electrons): 1s (2 electrons), 2s (2 electrons), 2p (1 electron).

• carbon (6 electrons): 1s (2 electrons), 2s (2 electrons), 2p (2 electrons).

hund's rule

• electrons in the same sublevel (p, d, f) occupy individual orbitals before pairing up.

• electrons won't occupy a filled orbital if an empty one is available in the same subshell.

• electrons added to orbitals have the same spin until half-full orbitals are achieved.

• boron: one electron in the 2p orbitals

• carbon: two electrons in separate 2p orbitals with the same spin.

• fluorine: fills the first, second and third 2p orbitals with the same spin, then adds fourth and fifth electrons with opposite spins.

aufbau principle

• electrons fill orbitals from lowest to highest energy.

• smaller n values fill before larger n values (1s before 2s, 2s & 2p before 3s & 3p, etc.).

• 3d subshell is higher in energy than 4s, so 4s fills before 3d.

• potassium (19 electrons):

  • fills 1s, 2s, 2p, 3s, 3p, then 4s with one electron.

• vanadium (23 electrons):

  • fills 1s, 2s, 2p, 3s, 3p, 4s, then 3d with three electrons in different orbitals with the same spin.

diagonal rule

• diagram to determine the order of filling electron configurations: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 5f, 6d, etc.

writing electron configurations

• determine the number of electrons (same as the atomic number for a neutral atom).

• fill subshells according to the aufbau principle, indicating the number of electrons as a superscript.

• example: cobalt (27 electrons)

  • 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^7

• check work by summing superscripts to match the total number of electrons.

dot structures