Chem 109 - General Chemistry - Spring 2011
Lecture Notes 24: 23 March
Trends in Chemical Periodicity
(plots ©1994 Hanson, Harper, Paselk, & Russell)
Last time we looked at the trends for atomic radii (atomic size) and first ionization energy. Today we continue with the trend for electronegativity, which is an indication of how electrons are shared by atoms in bonds.
Electronegativity increases from left 
right and from bottom top.
Highest Densities
Highest Melting Points
Due to multiple strong covalent bonds in Representative elements, and strong "covalent-metallic" bonding via unfilled d orbital electrons in Transition elements.
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 room 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 3p6 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. You may also wish to use a tool to help you remember this pattern of filling:
The 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
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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 where [Ne] = 1s2 2s2 2p6)
- V: [Ar] 4s2 3d3 (instead of 1s2 2s2 2p6 3s2 3p6 4s2 3d3 where [Ar] = 1s2 2s2 2p6 3s2 3p6)
- 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).
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.
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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.
© R A Paselk
Last modified 23 March 2011
