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Chem Oly Pre-Camp > Atomic Structure > Flashcards

Atomic Structure Flashcards

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1
Q

Trend of metallic character in a group

A

Increases down a group

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2
Q

Chemical properties of elements depend on ______ _______

A

Electron arrangement

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3
Q

Spectroscopy define

A

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

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4
Q

Greater jump of energy levels =

(In terms of energy, f and λ

A

Greater energy = greater frequency = lower wavelength

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5
Q

Schrodinger’s mathematical model - energy and position of electrons

A

Specific about electron energy, less specific about position

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6
Q

Define orbital

A

Regions where electrons are likely to be found

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7
Q

What are the 4 quantum numbers describing an electron?

A
  1. Principle, n (shell)
  2. Azimuthal, l (shape - sub shells)
  3. Magnetic, m (orbital orientation)
  4. Spin, ms (+-1/2)
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8
Q

Principle quantum number

A

The shell in which the orbital is located

Integer values

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9
Q

Azimuthal quantum number

A

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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10
Q

Magnetic quantum number

A

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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11
Q

s, p, d sub-shell lobes, number (types) of orbitals, max electron number

A

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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12
Q

max number of electrons in each orbital

A

magnetic spin +1/2 or -1/2

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13
Q

Pauli’s exclusion principle

A

No two electrons may have the same four quantum numbers

Limits # electrons per shell and sub shell

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14
Q

Define electron configuration

A

List of occupied orbitals

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

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15
Q

Orbitals of equal energy have the same _ and _

A

n and l

= same sub shell

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16
Q

s-block on periodic table

A

groups 1 and 2

Eg. Li = 2s1, Mg = 3s2

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17
Q

p-block on table

A

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

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18
Q

Define valence and core electrons

A

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

Core= inner shells, not involved in bonding

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19
Q

d-block on periodic table

A

Group 3 to 12, transition block

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20
Q

Where do the 3d orbitals begin on the periodic table?

A 
At Sc (4s2 3d1) 
4s is filled BEFORE 3d (eg K has 4s1 valence configuration) due to LOWER energy
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21
Q

Transition metals general valence configuration

A

ns2 (n-1)dx

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22
Q

Sum of superscripts of electron configuration =

A

Total number of electrons on the atom

23
Q

Transition metals usual valence configuration and exceptions

A

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.

24
Q

Transition metal ions orbital occupation

A

e- in 3d sub shell = total # valence e- (G# - charge)

For transition ions of charge >= +2, the 4s orbital is NOT occupied.

25
Q

Hund’s Rule

A

Electrons occupy orbitals of equal energy (same sub shell) SINGLY before a second electron enters any of them.

26
Q

Effective nuclear charge at outer electrons is determined by…

A

The number of electrons in the outer shell

AND

Number of electrons in shells between

27
Q

Atomic radius trends

A

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 ^

28
Q

Define ionisation enthalpy

Endo/exo?

A

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)

29
Q

Trends in ionisation enthalpy

A

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

30
Q

Define electron affinity (two ways)

A

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)

31
Q

Which ions have the highest electron affinity?

A

Halide ions require the most energy to remove as they have full s and p valence shells.

32
Q

Which atoms have 0 electron affinity?

A

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

33
Q

Define electronegativity

Combination of… ?

A

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

34
Q

Trends in electronegativity

A

Non-metals attract e- more

Increase —> and up (F most electroneg)

35
Q

How is electronegativity used to predict bonding type?

A

Difference in electroneg
Similar = covalent
Difference 2.0+ = ionic

36
Q

The sum of superscripts of electron configuration =

A

Total number of electrons on atom

37
Q

Characteristics of transition metal valence electrons (subshells)

A

Incomplete inner d or f subshell

Valence e- in 2 different subshells

38
Q

Transition metal valence configuration general

A

3dn 4s2

39
Q

Exceptions to the general transition metal valence configuration

A

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

40
Q

Explain relative atom and ion sizes with structure

A

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

41
Q

Determine size of ions with the same number of e-

A

Larger ion has fewer protons to attract e-.

42
Q

Factors affecting ionisation energy

A

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

43
Q

General trend of ionisatjon energy

A

Increases across the row

Extra proton lowers orbital energy more than extra electron repulsive forces

44
Q

Inconsistency in the general increase of ionisation energy across row

A

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-

45
Q

Characteristics of elements with lower ionisation enthalpies (below 1000 kJ mol-1)

A

Be electrical conductors
Exist in compounds as cations
(Metals)

46
Q

Successive ionisations

A

Each removal of e- requires more energy.

47
Q

Transition metal characteristics

Ox states, radius, ionisation energy changes

A

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

48
Q

All transition metals form compounds in which they exist as what ox state?

A

M2+, +2 state

EXCEPT Sc

49
Q

Maximum common ox state for Ni, Cu, Zn

A

+2

50
Q

M3+ ions are commkn for transition metals up to…

A

Co

51
Q

+3 is the max common ox state for which transition metals?

A

Fe, Co

And Sc

52
Q

What is the maximum oxidation state of the first 5 d-group elements?

A

Same as column number.

Sc3+ 
Ti(IV) 
V(V)
Cr(VI) 
Mn(VII)
53
Q

what is partial charge and how do you find the partial charge of an atom in a diatomic molecule?

A

covalent bonding, uneven sharing of e-.
The more ELECTRONEGATIVE atom has a NEGATIVE partial charge (pulls e- towards it)
delta- or d+

Chem Oly Pre-Camp (14 decks)

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