test 4

• we need 3 numbers to describe the system in hydrogen atoms and 4 numbers for non-hydrogen atoms
• the quantum numbers: principal (n), orbital angular momentum (l), magnetic angular momentum (ml), electron spin (ms)
• the principal quantum number relates to the energy and probable distance of the electron from the nucleus
• angular momentum describes the shape of the orbital where the electron is located
• the angular momentum number cannot be greater then the n value
• the number of subshells in the principal shell = the number of possible l values
• s subshell: L = 0
• p subshell: L = 1
• d subshell: L = 2
• f subshell: L = 3
• the magnetic quantum number can have a negative or positive whole-number value ranging from -L to L
• the number of orbitals in a subshell is equal to the number of allowed ml values for that subshell (2l + l)
• s orbitals have a spherical shape
• p orbitals have 2 lobes with a node in the middle
• electrons cannot be in a node
• d orbitals have multiple lobes and nodes
• the number of nodes is equal to the n value minus 1
• electrons with a higher n value will generally be higher in energy
• in a single electron atom, subshells (L) with the same n value are degenerate (have the same energy)
• in a multi electron atom, the l and n values have to be the same to be degenerate
• electrons close to the nucleus can shield the electrons further out from the full nuclear attraction, causing them to experience a lower effective nuclear charge (Zeff) and have a higher energy
• s orbitals (because they have no nodes at the nucleus) are able to get closer to the nucleus and experience a greater attractive force (penetration), and be at a lower energy
• electrons fill electron configuration to minimize the overall energy of the atom
• only two electrons can be in the same orbital, and they must have opposite spins
• instead of having a full ns orbital, chromium, copper, molybdenum, silver, and gold move one electron up to the nd orbital from ns
• when an electron is in an excited state, the electron will be missing from a lower energy subshell and moved into a higher energy subshell
• when an atom loses electrons (forms a cation) it loses the highest principal energy level (n) first
• the smaller the sum of the n value and the l value, the less energy present in the subshell
• electron configuration of transition metals: ns^2 (n-1)d^(1-10)
• electron configuration of f-block element: ns^2 (n-2)f^(1-14) (n-1)d^1
• group number for an element refers to the amount of valence electrons it has
• electrons in the same period have the same number of core electrons
• cations will cease to lose electrons once the outermost, highest energy subshell present in the neutral species is emptied
• anions will add enough electrons to fully fill their highest energy subshell and no more
• effective nuclear charger (Zeff) increases across the periodic table
• Zeff = number of valence electrons in an element
• ionization energy is always endothermic
• second and higher ionization energies are almost always larger (more endothermic) than the first ionization energies
• ionization energy increases across a row and up a group
• exceptions of ionization energy: group 2 > group 13 (3A) and group 15 > group 16