Atomic Structure Flashcards by Joanna Li

Trend of metallic character in a group

Increases down a group

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Chemical properties of elements depend on ______ _______

Electron arrangement

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Spectroscopy define

Study the absorption and emission of light to obtain info about electron energies

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Greater jump of energy levels =

(In terms of energy, f and λ

Greater energy = greater frequency = lower wavelength

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Schrodinger’s mathematical model - energy and position of electrons

Specific about electron energy, less specific about position

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Define orbital

Regions where electrons are likely to be found

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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)

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Principle quantum number

The shell in which the orbital is located

Integer values

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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

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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

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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-

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max number of electrons in each orbital

magnetic spin +1/2 or -1/2

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Pauli’s exclusion principle

No two electrons may have the same four quantum numbers

Limits # electrons per shell and sub shell

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Define electron configuration

List of occupied orbitals

Orbitals are filled in order of increasing energy (increasing n and l)

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Orbitals of equal energy have the same _ and _

n and l

= same sub shell

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s-block on periodic table

groups 1 and 2

Eg. Li = 2s1, Mg = 3s2

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p-block on table

Groups 13-18
G13 = p1, … G18 = p6
eg. Al = 3s2p1

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Define valence and core electrons

Valence = outermost infilled shell electrons
Involved in bond formation, determine chem properties

Core= inner shells, not involved in bonding

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d-block on periodic table

Group 3 to 12, transition block

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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

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Transition metals general valence configuration

ns2 (n-1)dx

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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

M3+ ions are commkn for transition metals up to…

+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+