Humboldt State University ® Department of Chemistry

Richard A. Paselk

	Science 331
Fall 2004	Lecture/Activity	Office: SA560a
Notes: 8 September		Phone: x 5719 Home: 822-1116
		e-mail: rap1

Chemical Periodicity

Let's look at the Periodic Table of the Elements again. Last time we looked at some examples of element properties, particularly for the alkali metals (group IA) and the halogens (group VIIA), as well as the variation in properties for the elements of Period 3 (Na - Cl). Now let's look at the last group of elements, the Noble Gases and then some terms. We will then look at more general properties.

• 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 and only known compounds 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.

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 below).
• Transition metal elements: the elements of the d-block (yellow in the table below).
• Inner-transition metal elements: The Lanthanides and Actinides (not shown on the table below)
• 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 and memorize the names and symbols for the elements shown in the table below.

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 then 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 (cancelling all nuclear charges, so no charge is visible from the outside), and the electrons of the atom are strongly attracted to the nucleus (they can all "see" many proton charges), so Ne will be 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.

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] file.
• 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] file.

Lewis Structures

Lewis strucutres are a very simple model for representing some chemical properties of atoms. For Lewis structures we assume all atoms above Be want eight electrons in their outermost shells. We represnt the nucleus and all inner electrons as the symbol or "kernal." Outer electrons are then represented as dots for individual electrons (electron bonding pairs may be represented as lines).

Lewis Structures for Atoms: Just show inner "kernel" where symbol stands for nucleus and all inner shell electrons. Only good for Representative elements. For ions the charge is always shown. Thus for metal ions such as calcium the Lewis Structure simply becomes the symbol for the ion. For negative ions such as we see for oxygen (2-) we enclose the ion and its electrons in brackets to indicate that the electrons are all "owned" by the oxygen - it does not share. Examples:

• Sodium:
• Sodium ion:
• Phosphorus:
• Bromine:
• Bromide ion:
• Sodium chloride:

Chemical Bonds

Atoms and molecules can be held together by Strong bonds or Weak bonds. We are first going to look at strong bonds.

Covalent vs. Ionic compounds: There are two kinds of strong bonds: ionic bonds and covalent bonds.

• In covalent compounds atoms have a definite relationship to each other, they are "married." Thus for water, H₂O, the smallest particle is a water molecule containing one oxygen and two hydrogens.
• In ionic compounds ions of opposite charge attract each other, but there is no definite attachment. Thus in sodium chloride crystals each sodium ion is surrounded by six chloride ions and each chloride ion is surrounded by six sodium ions and they are equally attracted by each - there is no one-to-one relationship.

We will begin our discussion with ionic bonds since they are easier to understand.

With the representative elements bond formation generally results in the formation of "octets" of electrons in the outermost shell.

Ionic bonds and Ionic Compounds:

• Formation: 2 Na + Cl₂ Æ NaCl. Can think of as composed of two equations, the oxidation of Na and the reduction of Cl:
  • Na + energy Æ Na⁺ + e^-
  • Cl + e^- Æ Cl^- + energy
• Lewis Structure
• Crystal Structure (model)
• Ionic Compound Lewis Structure Examples:
  • Potassium bromide
  • Aluminum chloride

Covalent Compounds and Covalent Bonds:

• Formation: H₂ + Cl₂ Æ 2 HCl. In this case can consider that we get two equations each involving a homo dissociation to give radicals, that is atoms with unpaired electrons:
  • H₂ Æ 2 H^.
  • Cl₂ Æ 2 Cl^. These radical then combine to form a bond with these two electrons shared between the two atoms.
• Lewis Structure:
• Covalent Compound Lewis Structure Examples:
  • Water
  • Ammonia
  • Ammonium ion
  • Methane
  • Hydrogen sulfide
  • Carbon dioxide
  • Carbon monoxide

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© R A Paselk

Last modified 8 September 2004