Inorganic Chemistry Flashcards
Electromagnetic spectrum: order
RMIVUXG
Electromagnetic radiation in order of frequency
GXUVIMR
Electromagnetic spectrum: names
Radio, Microwave, Infrared, UV, X-Ray, Gamma
Finding energy using wavelength
Energy is in kj. So wavelength (nm) is x10 to power of -9. Answer x10 to power of -3 (milli to 1)
how are lines on a spectrum formed?
An electron is promoted to a higher energy level. When the excited electron drops back to a lower level, energy is given off to a corresponding wavelength. DO NOT MENTION ABSORBING LIGHT.
absorption spectroscopy
electromagnetic radiation is directed at an atomised sample. radiation is absorbed as electrons are promoted to higher energy levels
emission spectroscopy
high temperatures are used to excite electrons, making them drop to lower energy levels, emitting a photon as they do so
atomic spectroscopy: conc. of an element
concentration of an element is related to the intensity of light emitted or absorbed
number of electrons found in an orbital
max of 2
n
principal quantum number, main energy level for an electron
shapes of orbitals
s, p, d, f
l
angular momentum quantum number. shape of subshell (s, p, d)
ml
orientation of the orbital. -l to +l values
ms
direction of spin. +1/2 of -1/2
the aufbau principle
electrons fill orbitals in order of increasing energy
hund’s rule
electrons fill degenerate orbitals singly and with parallel spins before pairing occurs
the pauli exclusion principle
no two electrons in one atom have the same set of four quantum numbers. therefore, no orbital can hold more than two electrons and those electrons must have parallel spins
division of periodic table into s, p, d, f
corresponds to the outer electronic configurations of the elements within these blocks
special stability
associated with half-filled and full subshells
stability and ionisation energy
the more stable the electronic configuration, the higher the ionisation energy
VSEPR
valence shell electron pair repulsion theory: used to predict the shapes of molecules and polyatomic ions. electron pairs are negatively charged so repel each other, therefore they are arranged to minimise repulsion and maximise separation
electron pair repulsion strength
decrease in order of: non-bonding pair/ non-bonding pair> non-bonding pair/ bonding pair> bonding pair/ bonding pair
transition metal
metals with an incomplete d subshell in at least one of their ions
copper and chromium
have a complete d subshell due to the special stability associated with having the d subshell being half-filled or completely filled
transition metals becoming ions
the 4s electrons are lost before the 3d electrons
oxidation state
is the oxidation number
oxidation number
the sum of all the oxidation numbers of all the atoms in a neutral compound must add up to zero. the sum of all the oxidation numbers of all the atoms in a polyatomic ion must be equal to the charge on the ion
oxidation
increase in oxidation number
reduction
decrease in oxidation number
oxidising agents
compounds containing metals in high oxidation states
reducing agents
compounds with metals in low oxidation states
ligands
negative ions or molecules with non-bonding pairs of electrons that they donate to the central metal atom or ion, forming dative covalent bonds
ligand classification
monodentate up to hexadentate
coordination number
total number of bonds from the ligands to the central transition metal
complex of a transition metal
d orbitals no longer degenerate
splitting of d orbitals
splitting to higher and lower energies occurs when the electrons present in approaching ligands cause the electrons in the orbitals lying along the axes to be repelled
strong field ligands
create a large difference in energy between subsets of d orbitals
weak field ligands
create a small energy difference between subsets of d orbitals
colours of transition metals
explained in terms of d-d transitions. light is absorbed when electrons in a lower energy d orbital are promoted to a d orbital of higher energy
heterogeneous catalysis
presence of unpaired d electrons or unfilled d orbitals allows activated complexes to form as reactive molecules are adsorbed, providing a reaction pathway with a lower activation energy. therefore, transition metals are often catalysts
homogeneous catalysis
changing oxidation states with the formation of intermediate complexes
Ionic bonds
Will form when atoms can rearrange their electrons to produce an arrangement of lower energy. Ionic lattices are more stable
Minimising electron repulsion
Lone pairs of electrons are always positioned equatorially to minimise repulsion. The shape of the molecule is also supposed to minimise this
Bonding pairs
Shape is dependent on number of bonded electrons, so answer to formula may not always be right
Octahedral
6 pairs. 90
Trigonal bipyramidal
5 pairs. Down 90, across 120
Tetrahedral
4 pairs. 109.5
Trigonal planar
3 pairs. 120
Linear
2 electron pairs. 180
Number of electron pairs, formula
Number of outer electrons on central atom + number of bonded atoms — charge. DIVIDE ALL THIS BY 2
Resonance structure
More stable than normal structure. A double and single bond will flip around e.g. ozone
Dative covalent bond
When one atom provides both of the electrons that form the covalent bond
Orbital boxes
1s, 2s, 2p, 3s, 3p, 4s, 3d. 4s comes before 3d if it is an ion or transition metal
Oxidation number
How many electrons have been lost
Coordination compound
Cation/ anion can’t exist on their own so must have charges balanced. The complex ion will bond with oppositely charged ions to form a complex
