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CHEMISTRY-FY > S1- L5-6: The Electronic Configuration of the Atom > Flashcards

S1- L5-6: The Electronic Configuration of the Atom Flashcards

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1
Q

What is “Quantum theory”?

A
  • theoretical basis of modern physics

- ->explains nature AND behavior of matter + energy on atomic plus subatomic level

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2
Q

Quantum theory explains some observations classical mechanics cant. Outline the 3 in this lecture

A

1-Spectra of light emitted by atoms
-e-‘s in atom stable + stay in their orbitals
2-wave-particle duality

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3
Q

Explain the first and third observations in appropriate detail

A

1- quantisation of energy AND energy levels (put energy in numbers like 400KJ
3-light able to behave as if made up of photons with energy depending on frequency
–>amount of light can see OR example depends on amount of photons coming from source which determines energy of light

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4
Q

Outline the following important concepts of quantum theory:

1-Uncertainity principle
2- e-‘s described as waves by Schrodinger equation

A

1-can never know exact location PLUS velocity of subatomic particle at same time (as constantly moving)
2-a mathematical function which describes e-‘s properties in quantum number terms

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5
Q

Describe the “Hydrogen spectrum”

A
  • shows existence of energy levels within atoms
  • e-‘s in atom are in stable energy levels
  • absorption of photon of light lets e- jump to higher level
  • jumping to lower level emits photon of light
  • measuring energy of photon lets energy difference between levels to be determined
  • ->showed energy levels not equally spaced
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6
Q

Define “ionisation energy”

A

-measures amount of energy needed to remove e-‘s from atoms/ions

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7
Q

What is “first ionisation energy”?

A

-energy needed to remove 1 mole of e-‘s (to infinity) form 1 mole of gaseous atoms to form 1 mole of gaseous (+) ions

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8
Q

Write the first ionisation energies of the following two elements:

1-Na/ 2-Al

A

1- Na(g) –> Na+ + e-

1- Al(g) –> A+ + e-

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9
Q

What is the effect of the charges of the particles involved?

A
  • e-‘s (-) charged AND protons in nucleus (+) charged
  • ->will be attraction between them
  • ->must add energy to system to pull e-‘s away
  • ionisation energies are (+)
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10
Q

Briefly explain the effect of the pull of the nucleus on ionisation energy

A
  • greater pull of nucleus
  • ->harder to pull e- away from atom
  • ->higher nuclear charge has higher ionisation energy
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11
Q

What do first ionisation energy patterns across a period and successive I. Energies for an element prove?

A

-give evidence for e-‘s being in subshell AND shells

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12
Q

Define “electron affinity”

A

-amount of energy needed to add e-‘s to atoms/ions

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13
Q

Outline what “First e- affinity” is and give an example

A
  • energy needed to add 1 mole of e-‘s to 1 mole of gaseous atoms to form 1 mole of gaseous (-) ions
  • Cl(g) + e- –> Cl- (g)
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14
Q

What does first electron affinities indicate?

A
  • indicate energy released on electron addition

- ->more (-) electron affinity the more stable negative ion formed

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15
Q

Explain the trend in atomic size in the periodic table (figure 1)

A
  • Increases as go down group as e- in higher energy levels
  • decreases as go across- as e- in same shell BUt increasing nuclear charge
  • ->so stronger attraction/pull to nucleus
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16
Q

Outline and explain the trend in first ionisation energy as you move down a group (figure 1)

A
  • Decreases down group as outer e- weaker attraction to nuclear charge (shielding) as atom bigger
  • increases across as shielding stays same BUT nuclear charge increases
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17
Q

What happens to the first electron affinity as you move down and across a group? (figure 1)

A
  • becomes less negative as go down group
  • ->as e- further away (larger atom) so more shielding and less attraction
  • more negative as go across periodic table
  • ->because incoming e- has stronger attraction to nucleus
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18
Q

Which key factors must be considered when explaining periodic table trends?

A
  • nuclear charge (protons in nucleus)
  • e- shells
  • shielding–> effect between outer e-‘s AND nucleus
  • ionisation energy
  • electronegativity
  • atomic size
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19
Q

State the purpose of quantum numbers

A
  • identify various energy levels available with atom in which e- can reside
  • identification numbers addresses each e- in atom
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20
Q

What do identification numbers specifically address for each e-?

A
  • specifically position (location) of e- in atom
  • predict direction of spin/rotation of e-
  • determine energy AND angular momentum of e-
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21
Q

Outline the 4 quantum numbers and their symbols

A
  • Principal Quantum Numbers (n)
  • Azimuthal Subsidiary Quantum Numbers (l)
  • Magnetic Quantum Numbers (m)
  • Spin Quantum Numbers (s)
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22
Q

Describe what an “orbital” is

A
  • region where approx 95% probability of finding particular e-
  • ->can not specify exact location of e-
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23
Q

What is a “node”?

A

-place less likely to find e- in orbital

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24
Q

Explain what principal energy levels are

A
  • In ply-electron atoms-principal energy levels called shells
  • ->referred to by letters K/L/M
  • ->these correspond to principal quantum numbers n. of “n” (n= number greater than 0)
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25
Q

What is the value of n for the different shells?

A
  • K–> n= 1
  • L–> n=2
  • shells contain multiple orbitals except n=1
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26
Q

Outline what happens to energy levels as n increases in complex atoms

A

-where more than 1 e- energy levels get closer together as n increases

27
Q

What does the Principal Quantum Numbers (n) indicate?

A
  • distance of e- from nucleus: higher n is–> further away e- from nucleus
  • energy of e-: higher n–>higher e- so less tightly held so easier to remove
  • n. of e- shell able to hold: n can have 2n^2 e-‘s
28
Q

Hoe many e-‘s do each the K/ L/ M/ N shell hold?

A
  • K shell: n=1 holds max 2x1^2= 2 e-‘s
  • L shell: n=2 holds 2x2^2= 8 e-‘s
  • M shell: n=3 holds 2x3^2= 18 e-‘s
  • N shell: n= 4 holds 2x4^2= 32 e-‘s
29
Q

What are subshells (L) and how are they described?

A
  • group of orbitals with same energy called subshell

- subshells described by azimuthal OR subsidiary quantum number l

30
Q

What is the value of (L) dependent on/determined by?

A
  • value of “L” depends on value of n
  • ->L can have value from 0 to (n-1)
  • ->i.e: l=0/1/2/2…(n-2), (n-2)
  • ->these correspond to different subshells which designated letters s/p/d/f
31
Q

Outline what the letters s/p/d/f stand for

A

-sharp/ principal/ diffuse/ fundamental

32
Q

State what the Azimuthal Quantum Numbers (L) indicates

refer to figure 2

A

L indicates:

  • which subshell e- is in
  • energy of subshell increases with increasing L
  • orbital shape in that subshell
  • max n. of e-‘s given subshell able to hold–> 2(2L+1)
33
Q

What are the shapes of the s/p/d/f orbitals? (refer to figure 3)

A
  • s- spherical
  • p- dumbbell
  • d- more complex
  • f- still more complex
34
Q

Outline what n and l are used to identify (refer to figure 4)

A
  • combination of n and l (nl) identify particular subshell

- ->so describe e- location in atomic energy levels

35
Q

What are Magnetic Quantum Numbers (m) and their significance? (figure 5)

A
  • represents orbitals in given sub-shell
  • indicates direction of particular orbital relative to magnetic field/axes
  • not indicate energy
  • “m” can have integral vale ranging from -l –>0 to +l
  • ->so for given l value total n. of m values is (2l +1)
36
Q

Figure 6-shape and orientation of orbitals

A

-Analyse and be able recognise

37
Q

Explain what Spin Quantum Numbers (S) are

A
  • e- moving around nucleus rotates/spins about it’s own axis clockwise OR anti-clockwise direction
  • ->directions described by spin quantum number
  • ->can have 2 values clockwise spinning of e- indicates (+1/2) AND anticlockwise (-1/2)
  • 2 e-‘s in same orbital must have opposite spins (+1/2 AND -1/2) also described as anti-parallel spins
38
Q

What is the significance of the clockwise and anti-clockwise spins? (refer to figure 7)

A
  • spins also known as up and down spins

- the opposite spinning of e- produces opposite magnetic field

39
Q

Quantum number’s summary:

1-Principal
2-Azimuthal
3-Magnetic
4-spin

Quantum number–> Restriction–> Range

A

1- n –> positive integer–> 1,2,3…..
2- l –> 0 AND positive integers less than n –> 0,1,2…(n-1)
3- m –> integers between -l and +l –> -l,..-1,0,+1,..+l
4- s –> +1/2 OR -1/2 –> +1/2 OR -1/2

40
Q

What is the Heisenberg’s Uncertainty Principle?

A

-Can not determine position AND momentum of e- at same time

41
Q

Outline the Aufbau Principle

A

-e-‘s enter lowest available energy level first

42
Q

Describe Hund’s Rule to Maximum Capacity

A
  • when in orbitals of equal energy e-‘s will try to remian unpaired
  • ->minimises repulsion between like charges so system more stable
43
Q

What is the Pauli Exclusion Principle?

A
  • no two e-‘s able to have same 4 quantum numbers

- ->only two e-‘s able to go in each orbital- providing they are of opposite spin

44
Q

Briefly explain what Madelung’s rule is

A
  • energy of atomic orbitals increases as (n+1) increases

- ->for identical values of n+1 energy increases with increasing n

45
Q

Analyse the figures of the max n. of e-‘s in subshells AND the max n. of e-‘s in energy levels

A

-Refer to figure 8 and figure 9

46
Q

What happens to energy levels as you get further from the nucleus?

A

-the energy levels get closer together

47
Q

Outline the order of filling orbitals up to 4f

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f

48
Q

Order of filling orbitals

A

-analyse figure 10

49
Q

Why do high values of n fill first before lower n values?

A

-due to interactions between multi-electron energy level atoms not equally being spaced out

50
Q

Periodic table and electron configuration

A

-analyse figure 11

51
Q

How would you write the electronic configuration of a Helium atom?

A
  • He atomic number is 2
  • ->start by filling 1s subshell- able to hold up to 2 e-‘s
  • ->write n. of e-‘s in each subshell as subscript
  • ->so e- configuration of He is 1s2
52
Q

Outline the method to write the e- configuration of an N atom

A
  • atomic number is 7
  • ->start filling 1s subshell which able to hold up to 2e’s
  • ->then 2s subshell (2 e-‘s) then 2p (6 e-s)
  • ->SO N e- configuration: 1s2, 2s2, 2p3
  • ->total n. of e-‘s in all orbitals must add up to atomic number
53
Q

How may the electron configuration of N be written in more detail?

A

-N–> 1s2, 2s2, 2p3
-there are 3 p orbitals (m= +1, 0, -1) in 2p sub-shell
-Hund’s rule–> when in orbitals of equal energy e-‘s will try to remain unpaired
–>so can write e- configuration in more detail:
1s2, 2s22px12py12pz1
–>or as diagram (figure 12)

54
Q

What are the special cases in the e- configuration rules?

A
  • Chromium- atomic number is 24
  • ->1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5
  • Copper
  • ->1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10
55
Q

How are positive and negative ions formed?

A
  • (+ cations): formed by removing e-‘s from atoms

- (- anions): formed by adding e-‘s to atoms

56
Q

In what order must e-‘s be removed in the formation of ions?

A

-e-‘s first removed from highest occupied orbitals except transition metals

57
Q

Outline the formation of Na and Cl ions

A
  • Na: 1s2, 2s2, 2p6, 3s1
  • -> remove 1 e- from 3s orbital to form Na+ (1s2, 2s2, 2p6)
  • Cl: 1s2, 2s2, 2p6, 3s2, 3p5
  • ->add 1 e- to 3p orbital to form Cl- (1s2, 2s2, 2p6, 3s2, 3p6)
58
Q

How are ions formed from transition metals?

A

-e-‘s in 4s orbital removed before any e-‘s in 3d orbital

59
Q

Example: Titanium

A

-refer to figure 13

60
Q

Figure 14 shows the successive ionisation energies of Calcium. What does a large jump in I. Energy indicate?

A

-large jump between successive I. Energy indicates change in energy level from which e- been removed

61
Q

Analyse the trend in 1st variation of I. Energy across a period

A

-Figure 15- shows the trend

62
Q

Why is GD3+ commonly used for MRI scanning image enhancement?

A
  • GD3+ has 7 unpaired e-‘s

- ->better image of structures inside body

63
Q

Summary of lecture

A
  • e-‘s occupy shells around nucleus
  • ->each shell only able to carry set n. of e-‘s
  • shells contain orbitals–>region of space where e- likely to be found
  • e’s enter orbitals at lowest energy level first
  • orbitals filled in increasing energy order
  • orbitals AND their e-‘s characterised by 4 parameters-Quantum Numbers
  • quantum numbers describe size/shape AND spatial orientation of e-‘s probability distribution/density AND its spin
  • orbitals; shape AND filling rules
  • how to write e- configuration
  • ionisation energy/ trends in ionisation energy AND their explanations

CHEMISTRY-FY (11 decks)

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