Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109	General Chemistry	Summer 2002
Lecture Notes:: 1 July	© R. Paselk 2002
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Atomic Structure & Chemical Periodicity, cont.

Chemical Periodicity

Trends: Note the trends for

atomic size: decreases going from left Æ right and from bottom Æ top.

• Size goes up with atomic number for any individual group.
• Size decreases irregularly as atomic number increases for any given period (more charge pulls electrons in to nucleus, but shielding reverses as subshells [s or p orbital sets] fill.
• ionization energy: increases from left Æ right and from bottom Æ top.

• Ionization energy goes down with atomic number for any individual group.
• Ionization energy increases irregularly as atomic number increases for any given period (more charge pulls electrons in to nucleus, but shielding reverses as subshells [s or p orbital sets] fill.
• electronegativity increases from left Æ right and from bottom Æ top.

• note hydrogen combining ratios (LIH, BeH₂, BH₃, CH₄, H₃N, H₂O, HF) and acid/base properties of oxides (basic for metals, acidic for non-metals)

What is the basis of the periodicity of properties?

Electrons are held in shells.

• The first shell holds only 2 electrons in what is called the 1s orbital. Thus helium has its first shell filled. There is no more room for electrons, so it can't react by picking up another electron. On the other hand, as a crude thought model, we can consider that each electron is held by both charges in the He nucleus, so they are much more tightly held than the electron in H, so He won't give up an electron either - its inert.
• The second shell is larger (its out further from the nucleus) so holds 2 electrons in a 2s orbital, but there is now room for an additional three 2p orbitals. Thus 8 electrons can be accommodated in the second shell. Note the consequences:
  • for lithium (Li) the inner two electrons of the 1s shell cancel the attraction of two of the three protons, so the outer 2s electron "sees" only a single charge. But its out further than the electrons in the 1s shell were, so its not held as strong, so Li loses its outer electron more readily than H and is more reactive.
  • for Fluorine on the other side of the chart we can think of the outer shell electrons being attracted to the nucleus by 9 - 2 = 7 charges, so the last open space in an orbital will be super attractive to an outside electron, so F will be be very reactive, but in an opposite way to Li - it wants to steal electrons instead of giving them up.
  • for neon all of the orbitals will be filled, and the electrons will be strongly attracted to the nucleus and there is no for additional electron in the ground state, so Ne will again be inert like He above it.
• The third shell is larger yet (further from the nucleus), but still crowded, so initially it can only accommodate another eight electrons.
  • Of course the electrons in the 3s orbitals are even farther out from the nucleus, so we would expect Na to be even more reactive than Li, and so on for K, Rb, etc. each giving up its outermost electron more readily than the element above it in the Periodic table.
  • On the other hand Cl will also attract electrons less than F, so it will be less reactive etc. for the halogens.
• So we will expect the most reactive elements to be on the opposite corners of the table - lower left and upper right.

Electronic Configurations & Periodicity

There are a number of different notation conventions for electronic configurations.

• Spectroscopic notation: in this convention we indicate shells (main energy levels) by numbers, orbitals within these shells by letters (s, p, d, or f), and the number of electrons in each orbital type by superscript. For example:
  • H: 1s¹
  • He: 1s²
  • B: 1s² 2s² 2p¹
  • P: 1s² 2s² 2p⁶ 3s² 3p³
  • V: 1s² 2s² 2p⁶ 3s² 3p³ 4s² 3d³

Note that when we get to the d electrons they are added to the next inner orbital - they are added inside the atom and are not outermost! Note also that you may write them in the order they show up on the Periodic Table, or, if you prefer, you may group them by shell

• Spectroscopic notation using the Noble gas core convention. Notice that at the end of each period the outermost shell (s &  p) is filled, and when you go to the next element (e.g. Na) its as if you are adding onto the electronic configuration of the Noble gas. So we can save a lot of writing if we substitute its symbol for these inner electrons (note we are not really assuming an inner noble gas, we are just creating a type of short-hand). For example:
  • P: [Ne] 3s² 3p³ (instead of 1s² 2s² 2p⁶ 3s² 3p³)
  • V: [Ar] 4s² 3d³ (instead of 1s² 2s² 2p⁶ 3s² 3p³ 4s² 3d³)
  • U: [Rn] 7s²5f⁴ The savings here is really obvious! Note that if you use the Periodic chart that comes on exams you will fill the f's before the d's of a given period (this is not the case for the wall chart).

Aufbau Principle: Pattern of electron addition to atoms. Electrons fill atoms by sequentially filling hydrogen-like orbitals in order of energy. Use to predict electron patterns in atoms. But there are some variations, since in fact we are also adding charge to nucleus.

Aufbau Pattern: This is a useful aid for remembering order - draw diagonal arrows with the points on the upper left through the diagram below parallel to the diagonal through 2p and 3s. Following the arrows in order starting with the 1s gives the predicted filling order:

7s 7p

6s 6p 6d

5s 5p 5d 5f

4s 4p 4d 4f

3s 3p 3d

2s 2p

1s

 

• Orbital Filling Diagrams: In orbital filling diagrams we provide a little more information - noting that the electrons come in two different spins and that they fill into orbitals with their spins paired in opposite directions. I have no preference for the direction of unpaired arrows other than that they should be the same. For example:
  • 
  • 
  • 
  • Notice how the electrons first fill into empty orbitals before they pair up
  • Now when we add one more electron it goes back to pair up with the first p electron.

• Orbital Filling Diagrams using the Noble gas core convention. You can also use the Noble gas core convention with orbital filling diagrams, just like the spectroscopic notation, but using arrows etc.

• Ions: When an atom loses electrons we would expect it to lose its outermost electrons first. But which are outermost? Remember the "last added" electrons in the transition elements are in the d orbitals of the next outermost shell. The the d orbital electrons should not be the outermost electrons in an atom. Thus we will lose the s & p electrons first then the d electrons if any are present. If additional electrons are lost then we can go into the d shell. Examples:
  • Na⁺ = 1s² 2s² 2p⁶ 3s⁰ or [Ne] 3s⁰ (In both cases the 3s⁰ is usually not shown, I've shown it here for clarity.)
  • Cu²⁺ = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s⁰ 3d⁹ or [Ar] 4s⁰ 3d⁹ (In both cases the 4s⁰ is usually not shown, I've shown it here for clarity.)
  • Fe³⁺ = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s⁰ 3d⁵ or [Ar] 4s⁰ 3d⁵ (In both cases the 4s⁰ is usually not shown, I've shown it here for clarity.)

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Last modified 1 July 2002