Atomic Structure Flashcards
Trend of metallic character in a group
Increases down a group
Chemical properties of elements depend on ______ _______
Electron arrangement
Spectroscopy define
Study the absorption and emission of light to obtain info about electron energies
Greater jump of energy levels =
(In terms of energy, f and λ
Greater energy = greater frequency = lower wavelength
Schrodinger’s mathematical model - energy and position of electrons
Specific about electron energy, less specific about position
Define orbital
Regions where electrons are likely to be found
What are the 4 quantum numbers describing an electron?
- Principle, n (shell)
- Azimuthal, l (shape - sub shells)
- Magnetic, m (orbital orientation)
- Spin, ms (+-1/2)
Principle quantum number
The shell in which the orbital is located
Integer values
Azimuthal quantum number
Specifies shape.
Principle n shell has n sub shells (l)
eg. n=2 has l=0, l=1 subshells
Also s, p, d, f
Magnetic quantum number
Orientation
Allowed integer values from -l to l inclusive
Eg. p sub shell (l=1) has three orbitals, m=-1,0,1
s, p, d sub-shell lobes, number (types) of orbitals, max electron number
Type = l x2 +1
Max e- = type x2
s: 1 lobe, 1 type, max 2 e-
p: 2 lobes, 3 orbitals, 6 e-
d: 4 lobes, 5 orbitals, 10 e-
max number of electrons in each orbital
magnetic spin +1/2 or -1/2
Pauli’s exclusion principle
No two electrons may have the same four quantum numbers
Limits # electrons per shell and sub shell
Define electron configuration
List of occupied orbitals
Orbitals are filled in order of increasing energy (increasing n and l)
Orbitals of equal energy have the same _ and _
n and l
= same sub shell
s-block on periodic table
groups 1 and 2
Eg. Li = 2s1, Mg = 3s2
p-block on table
Groups 13-18
G13 = p1, … G18 = p6
eg. Al = 3s2p1
Define valence and core electrons
Valence = outermost infilled shell electrons
Involved in bond formation, determine chem properties
Core= inner shells, not involved in bonding
d-block on periodic table
Group 3 to 12, transition block
Where do the 3d orbitals begin on the periodic table?
At Sc (4s2 3d1) 4s is filled BEFORE 3d (eg K has 4s1 valence configuration) due to LOWER energy
Transition metals general valence configuration
ns2 (n-1)dx
Sum of superscripts of electron configuration =
Total number of electrons on the atom
Transition metals usual valence configuration and exceptions
Usually 3dn 4s2
EXCEPT Cr and Cu
Because having more 3d orbitals is of lower energy for Cr and Cu
Cr, 3d5 4s1, has max number of UNPAIRED electrons (lower energy) for a d sub shell.
Cu, 3d10, 4s1
Energy of 3d orbitals decreases across the row, at the end it is lower than 4s.
Transition metal ions orbital occupation
e- in 3d sub shell = total # valence e- (G# - charge)
For transition ions of charge >= +2, the 4s orbital is NOT occupied.
Hund’s Rule
Electrons occupy orbitals of equal energy (same sub shell) SINGLY before a second electron enters any of them.
Effective nuclear charge at outer electrons is determined by…
The number of electrons in the outer shell
AND
Number of electrons in shells between
Atomic radius trends
Determined by size of e- cloud
Increases DOWN a group
Same period, atoms at the right are smaller, because valence e- in shape shell but greater nuclear charge.
Decreases —> and ^
Define ionisation enthalpy
Endo/exo?
A measure of the energy required to remove an electron from a GASEOUS atom
Always ENDOTHERMIC (require separation of charge)
E(g) —> E+(g) + e-(g)
Trends in ionisation enthalpy
Increases —> and ^
Reflects atomic size - less energy required to remove e- from outer shells of larger atoms.
Metal atoms lose e- more easily than non metal.
Some inconsistencies
Define electron affinity (two ways)
Energy required for removal of an electron from a gas phase ANION
Also a measure of the element’s tendency to act as an oxidising agent (electron acceptor)
Which ions have the highest electron affinity?
Halide ions require the most energy to remove as they have full s and p valence shells.
Which atoms have 0 electron affinity?
Group 2 atoms.
Full valence s-subshell
adding another e- would have to be added to the higher np sub shell.
Same with Noble Gases - completely filled shell, incoming e- added to higher n level
Define electronegativity
Combination of… ?
The ability of a bonded atom to attract electrons to itself when it is part of a compound.
Combination of electron affinity and ionisation energy
Trends in electronegativity
Non-metals attract e- more
Increase —> and up (F most electroneg)
How is electronegativity used to predict bonding type?
Difference in electroneg
Similar = covalent
Difference 2.0+ = ionic
The sum of superscripts of electron configuration =
Total number of electrons on atom
Characteristics of transition metal valence electrons (subshells)
Incomplete inner d or f subshell
Valence e- in 2 different subshells
Transition metal valence configuration general
3dn 4s2
Exceptions to the general transition metal valence configuration
Chromium (Cr) 3d5 4s1
Copper (Cu) 3d10 4s1
Because 3d and 4s orbitals very close energy
Having for 3d is lower energy for Cr and Cu
Explain relative atom and ion sizes with structure
Nonmetallic element —> anion
MORE e- = larger than parent atom
Metallic atom —> cation
LESS e- = smaller than atom
The BIGGER ION arises from the smaller atom
(Li>F, but F->Li+)
Cation is SMALLER than anion of the same row
Determine size of ions with the same number of e-
Larger ion has fewer protons to attract e-.
Factors affecting ionisation energy
Protons in the nucleus
(More protons, e- require higher ionisation energy)
Core electrons
increase orbital energy of e- due to shielding, less energy required
Same shell e-
Repulsive forces, increase orbital energy
General trend of ionisatjon energy
Increases across the row
Extra proton lowers orbital energy more than extra electron repulsive forces
Inconsistency in the general increase of ionisation energy across row
When effect of additional e- is more significant than the additional proton.
Dip at the first p-block atom - p-orbitals on average further from nucleus (lower ionisation energy)
Dip between 3rd and 4th p-block elements, as orbital energies are lower if each orbital in the subshell has an unpaired e-
Characteristics of elements with lower ionisation enthalpies (below 1000 kJ mol-1)
Be electrical conductors
Exist in compounds as cations
(Metals)
Successive ionisations
Each removal of e- requires more energy.
Transition metal characteristics
Ox states, radius, ionisation energy changes
Multiple ox states
Relatively gradual decrease in radius (3d shielding outer 4s, vs s or p block where e- added in same shell)
No abrupt change of ionisation energy from Ca to Sc.
Highest energy e- for d-block is 4s
All transition metals form compounds in which they exist as what ox state?
M2+, +2 state
EXCEPT Sc
Maximum common ox state for Ni, Cu, Zn
+2
M3+ ions are commkn for transition metals up to…
Co
+3 is the max common ox state for which transition metals?
Fe, Co
And Sc
What is the maximum oxidation state of the first 5 d-group elements?
Same as column number.
Sc3+ Ti(IV) V(V) Cr(VI) Mn(VII)
what is partial charge and how do you find the partial charge of an atom in a diatomic molecule?
covalent bonding, uneven sharing of e-.
The more ELECTRONEGATIVE atom has a NEGATIVE partial charge (pulls e- towards it)
delta- or d+
