Structure IB Flashcards
(59 cards)
Principal energy levels
N=1 closest to nucleus has lowest energy
Electrons in each energy level =2n^2
Main energy levels are split into sub levels
•n=1 (1s)
•n=2 (2s, 2p)
Isotope properties
Same chemical properties cuz e- count stays the same
Diff physical properties cuz neutron count is diff
Aufbau principle
Fill lower energy levels before higher ones
Pauli exclusion principle
Atomic orbital can hold 2e- with opp spins
Have to pair electrons in one level before putting electrons in next
Hund’s rule
In sub level fill each degenerate orbital with electrons with single parallel spin electrons before pairing them in that degenerate orbital
S orbital
Spherical has max 2 electrons
P orbital
Has 3 degenerate orbitals (Px Py Pz)
Each degenerate can hold 2 so 6 total electrons
Dumbbell shaped
Flame test
Identified metal ions in solution
Heat causes e- to absorb energy and transition to higher energy level. As they emit energy they go back to their og state and the energy emitted corresponds to a wavelength
Line Spectra’s
Continuous- all wavelengths
Absorption- black lines on coloured background
Emission- coloured lines on black background
Ionization energy
Energy needed for electrons to move from n=1 to n=infinity
First IE is energy needed to remove 1mol of e- from 1 mol of gaseous atom to produce a cation and e- (endothermic)
Increases across period (nuclear charge increases and atomic radii decreases so more energy needed)
Decreased down group (with more energy levels, valence e- are farther from nucleus and easier to remove
Electron Sheilding
Inner e- shield outer valence e- from full attraction to nuclear
Constant across period
Increases down a group
Effective nuclear charge
Net positive charge experienced by valence e-
Atomic number - shielding electrons
Constant across group
Increases left to right
Trends in atomic radii
Decreases across period
Increases across group cuz more energy levels occupied
Nuclear charge increased across period
As atomic radii decreases attraction between nuclear and outer electrons increases
Trends in ionic radii
If attraction between nucleus and outer e- increases than ionic radii decreases
Cations are smaller than parent atom cuz they lose e- while anions are bigger cuz they gain
Electronegativity
Increased across period
Decreased down group as bonding electrons are farther from the nucleus
Metallic character and metallic binding
MC- how easily atom can lose e- (increases down group decrease left to right)
MB-electrostatic attraction between lattice e of cations and delocalized e-
Electron affinity
Opposite of ionization energy so it’s the energy released adding electrons to atom to make it anion (exothermic)
Greater the atomic radii means greater number of shielding e- =decrease in affinity
Transition element
Atom w incomplete d-sub level (zinc is not transition metal cuz it has complete d sub level)
Complex ions
Contains central metal ion and ligands (have lone pairs that can form coordinate covalent bonds (both electrons in bond come from same atom))
Determining charge if complex ion
If ligands are neutral than complex ion has same charge as central metal
If negative than subtract total charge of ligands from the charge of the metal
Coordination componds
Contain the complex ion and a counter ion to balance the charge of complex ion
Resonance structures
More than 1 position for double bind in molecule
Allotropes
Diff physical forms of the same element (carbon can be diamond graphite )
Difference IMF’s
LDF- movement of electrons
DD- electrostatic attraction between partial positive and negative
H bonding- H to NOF
