S1- L5-6: The Electronic Configuration of the Atom Flashcards
What is “Quantum theory”?
- theoretical basis of modern physics
- ->explains nature AND behavior of matter + energy on atomic plus subatomic level
Quantum theory explains some observations classical mechanics cant. Outline the 3 in this lecture
1-Spectra of light emitted by atoms
-e-‘s in atom stable + stay in their orbitals
2-wave-particle duality
Explain the first and third observations in appropriate detail
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
Outline the following important concepts of quantum theory:
1-Uncertainity principle
2- e-‘s described as waves by Schrodinger equation
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
Describe the “Hydrogen spectrum”
- 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
Define “ionisation energy”
-measures amount of energy needed to remove e-‘s from atoms/ions
What is “first ionisation energy”?
-energy needed to remove 1 mole of e-‘s (to infinity) form 1 mole of gaseous atoms to form 1 mole of gaseous (+) ions
Write the first ionisation energies of the following two elements:
1-Na/ 2-Al
1- Na(g) –> Na+ + e-
1- Al(g) –> A+ + e-
What is the effect of the charges of the particles involved?
- 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 (+)
Briefly explain the effect of the pull of the nucleus on ionisation energy
- greater pull of nucleus
- ->harder to pull e- away from atom
- ->higher nuclear charge has higher ionisation energy
What do first ionisation energy patterns across a period and successive I. Energies for an element prove?
-give evidence for e-‘s being in subshell AND shells
Define “electron affinity”
-amount of energy needed to add e-‘s to atoms/ions
Outline what “First e- affinity” is and give an example
- 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)
What does first electron affinities indicate?
- indicate energy released on electron addition
- ->more (-) electron affinity the more stable negative ion formed
Explain the trend in atomic size in the periodic table (figure 1)
- 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
Outline and explain the trend in first ionisation energy as you move down a group (figure 1)
- 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
What happens to the first electron affinity as you move down and across a group? (figure 1)
- 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
Which key factors must be considered when explaining periodic table trends?
- nuclear charge (protons in nucleus)
- e- shells
- shielding–> effect between outer e-‘s AND nucleus
- ionisation energy
- electronegativity
- atomic size
State the purpose of quantum numbers
- identify various energy levels available with atom in which e- can reside
- identification numbers addresses each e- in atom
What do identification numbers specifically address for each e-?
- specifically position (location) of e- in atom
- predict direction of spin/rotation of e-
- determine energy AND angular momentum of e-
Outline the 4 quantum numbers and their symbols
- Principal Quantum Numbers (n)
- Azimuthal Subsidiary Quantum Numbers (l)
- Magnetic Quantum Numbers (m)
- Spin Quantum Numbers (s)
Describe what an “orbital” is
- region where approx 95% probability of finding particular e-
- ->can not specify exact location of e-
What is a “node”?
-place less likely to find e- in orbital
Explain what principal energy levels are
- 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)
What is the value of n for the different shells?
- K–> n= 1
- L–> n=2
- shells contain multiple orbitals except n=1
Outline what happens to energy levels as n increases in complex atoms
-where more than 1 e- energy levels get closer together as n increases
What does the Principal Quantum Numbers (n) indicate?
- 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
Hoe many e-‘s do each the K/ L/ M/ N shell hold?
- 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
What are subshells (L) and how are they described?
- group of orbitals with same energy called subshell
- subshells described by azimuthal OR subsidiary quantum number l
What is the value of (L) dependent on/determined by?
- 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
Outline what the letters s/p/d/f stand for
-sharp/ principal/ diffuse/ fundamental
State what the Azimuthal Quantum Numbers (L) indicates
refer to figure 2
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)
What are the shapes of the s/p/d/f orbitals? (refer to figure 3)
- s- spherical
- p- dumbbell
- d- more complex
- f- still more complex
Outline what n and l are used to identify (refer to figure 4)
- combination of n and l (nl) identify particular subshell
- ->so describe e- location in atomic energy levels
What are Magnetic Quantum Numbers (m) and their significance? (figure 5)
- 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)
Figure 6-shape and orientation of orbitals
-Analyse and be able recognise
Explain what Spin Quantum Numbers (S) are
- 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
What is the significance of the clockwise and anti-clockwise spins? (refer to figure 7)
- spins also known as up and down spins
- the opposite spinning of e- produces opposite magnetic field
Quantum number’s summary:
1-Principal
2-Azimuthal
3-Magnetic
4-spin
Quantum number–> Restriction–> Range
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
What is the Heisenberg’s Uncertainty Principle?
-Can not determine position AND momentum of e- at same time
Outline the Aufbau Principle
-e-‘s enter lowest available energy level first
Describe Hund’s Rule to Maximum Capacity
- when in orbitals of equal energy e-‘s will try to remian unpaired
- ->minimises repulsion between like charges so system more stable
What is the Pauli Exclusion Principle?
- 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
Briefly explain what Madelung’s rule is
- energy of atomic orbitals increases as (n+1) increases
- ->for identical values of n+1 energy increases with increasing n
Analyse the figures of the max n. of e-‘s in subshells AND the max n. of e-‘s in energy levels
-Refer to figure 8 and figure 9
What happens to energy levels as you get further from the nucleus?
-the energy levels get closer together
Outline the order of filling orbitals up to 4f
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f
Order of filling orbitals
-analyse figure 10
Why do high values of n fill first before lower n values?
-due to interactions between multi-electron energy level atoms not equally being spaced out
Periodic table and electron configuration
-analyse figure 11
How would you write the electronic configuration of a Helium atom?
- 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
Outline the method to write the e- configuration of an N atom
- 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
How may the electron configuration of N be written in more detail?
-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)
What are the special cases in the e- configuration rules?
- Chromium- atomic number is 24
- ->1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5
- Copper
- ->1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10
How are positive and negative ions formed?
- (+ cations): formed by removing e-‘s from atoms
- (- anions): formed by adding e-‘s to atoms
In what order must e-‘s be removed in the formation of ions?
-e-‘s first removed from highest occupied orbitals except transition metals
Outline the formation of Na and Cl ions
- 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)
How are ions formed from transition metals?
-e-‘s in 4s orbital removed before any e-‘s in 3d orbital
Example: Titanium
-refer to figure 13
Figure 14 shows the successive ionisation energies of Calcium. What does a large jump in I. Energy indicate?
-large jump between successive I. Energy indicates change in energy level from which e- been removed
Analyse the trend in 1st variation of I. Energy across a period
-Figure 15- shows the trend
Why is GD3+ commonly used for MRI scanning image enhancement?
- GD3+ has 7 unpaired e-‘s
- ->better image of structures inside body
Summary of lecture
- 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
