Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109 - General Chemistry - Spring 2015

Lecture Notes 24: 27 March

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Atomic Structure & Chemical Periodicity

The Periodic Table

Look at the Periodic Chart. The pattern arises due to a repetition or periodicity of chemical properties. The vertical columns of the charts are called groups, while the rows are referred to as periods.

Note the numbering of the groups. The numbers from 1 - 18 are the internationally accepted numbers. We will also use the I - VIII "American" numbering system. Note that the "tallest" columns comprise what are referred to as the "representative elements" (IA - VIIIA).

Terms:

• Period: the rows of elements showing a repeating pattern of properties (e.g. Na - Ar).
• Group: a vertical column of elements on the table sharing a family resemblance of properties (e.g. Li - Fr).
• Representative elements: the elements of the s-block and p-block (blue and green on the table above).
• Transition metal elements: the elements of the d-block (yellow in the table above).
• Inner-transition metal elements: The f-block or Lanthanides and Actinides (not shown on the table above)
• Groups:
  • IA = alkali metals;
  • IIA = Alkaline earth metals;
  • VIIA = Halogens (note the generic symbol of X standing for any halogen);
  • VIIIA = Noble gases (older = inert gases).

You should know the terminology above.

Recall the introduction to Chemical Periodicity in Lecture 6, including hydrogen combining ratios (LIH, BeH₂, BH₃, CH₄, H₃N, H₂O, HF) and acid/base properties of oxides (basic for metals, acidic for non-metals)

Let's look at some of the elements and see what their properties are like:

• Group IA, on the left side of the chart, is known as the alkali metals because they react with water to produce strong bases (a base is alkaline). Note that all of them are soft (cut with a butter knife), low density (Li floats on oil, Na and K float on water), very reactive metals. All of them react with water with Li<Na<K<Rb<Cs. In each case the metal gives its electron to water leaving hydroxide ion (OH^- a base) and hydrogen gas. For example with sodium:

2 Na + 2 H₂O 2 Na⁺ + 2 OH^- + H₂

• Group VIIA, on the right side of the chart, is known as the halogens. The halogens form acids with water, are gases at the top of the Periodic Chart and high vapor pressure liquids, then solid going down; exist as diatomic molecules (X₂), and are very reactive towards metals. For example sodium reacts violently with chlorine gas to give table salt (NaCl):

2 Na + Cl₂ 2 NaCl

• Group VIII is known as the Noble Gases, or sometimes the Inert Gases because until the 1960's they had no known compounds. Very unreactive. The only known compounds of the Noble gases are with very reactive elements like F and O, and even they don't form compounds with smaller Noble gases such as He and Ne.

Trends:Note the trends for:

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

Next we continue with the trend for first ionization energy (the energy needed to strip the outermost electron from a free atom) and electronegativity (an indication of how electrons are shared by atoms in bonded atoms).

First ionization energy: increases from left right and from bottom 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 right and from bottom top.

Electronegativity is a measure of how electrons are shared between two interacting atoms. It is an empirical (experimental) measure, ranging in value from <1 (Cs) to 4.0 (F). Note that the truely inert Noble gases such as He and Ne will have NO electronegativty values since they don't bond

Group 1 elements have the least tendency to attract bonding electrons, while F has the greatest attraction for electrons in bonds.

Note and memorize the electronegativities for H (2.1) and the elements of the second Period (Li {1.0}, Be {1.5}... F {4.0})

Highest Densities

Note, elements are near "center" of each Period. Due to a combination of nuclear mass, size and packing in crystals.

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.

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

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© R A Paselk

Last modified 27 March 2015