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A Brief History of the Development of the Periodic Table

Long before we know much about atomic structure, it was known that certain patterns existed among the elements.

Dobereiner (1829) discovered triads of elements that reacted in similar ways. For example, Li, Na and K belonged to one triad and Cl, Br and I belonged to another.

Newlands (1863) discovered a repeating pattern among some of the elements. Every eigth element reacted similarly. This was known as the law of octaves.

Mendeleev (1869) classified the elements according to atomic mass. He created a periodic table that had similar behaving elements in the same column. There were a few atoms that seemed to be out of place. Te and I seemed to be switched.

Moseley (1912) found that properties correlated better with atomic number rather than atomic mass. The modern periodic table is based on atomic number.

Electron Configuration and the Periodic Table

With our knowledge of electron configuration, we can correlate position on the periodic table with outer electron configuration. The A group number is correlated with outer electron configuration as illustrated below.

 The B groups also have group numbers that indicate how the electrons are configured. For example, IIIB has an outer electron configuration of ns² (n-1)d¹.

The period number indicates the principle quantum number of the outermost energy level. In other words, it tells us how many energy levels there are in an atom.

Chlorine is in group IIIA and period 3. This means its outer electron configuration is ...3s² 3p⁵. The 3 that appears before the two sublevels indicates the principle quantum number of the outer energy level.

Trends in Atomic Radius on the Periodic Table

The elements on the periodic table are arranged according to atomic number. There are repeating patterns in other properties besides electron configuration.  The atomic radius follows the following trends with a few exceptions:

·        Within a vertical column (group), the radius increases as you move down the group.  I.e. r_K > r_Na > r _Li

Moving down a group adds an additional energy level, increasing the size of the electron cloud.

·        Within a horizontal row (period), the radius decreases as you move from left to right.   I.e. r _Na > r_Mg > r_Al

Moving from left to right across a period increases the amount of positive charge in the nucleus. This increase pulls the negatively charged electron cloud closer to the nucleus and decreases the size of the electron cloud.

What Happens When an Atom Reacts?

What is it about an element that makes it react in a certain way? While 99.99% of an atom's mass is in the nucleus, it is the electron arrangement that determines the reactivity of an element. When an element reacts it either loses, gains or shares electrons. The easier this happens, the more reactive an element is. Conversely, an element that does not lose, accept or share electrons readily is considered inert or nonreactive.

Elements in the same group have similar electron configurations. Even so, there are significant differences in the reactivity of elements within the same group. Why would two elements with similar electron configurations behave so differently?

Losing Electrons

Some elements lose electrons when they react. We can measure the ease with which an element loses electrons.

Ionization Energy - the energy required to remove an electron from an atom in its gaseous state.

• The lower the ionization energy, the easier it is to remove the outer electron.

• The higher the ionization energy, the more difficult it is to remove the outer electron.

There are 4 factors that affect the ionization energy. They are:

These factors plus the arrangement of the elements on the periodic table result in trends in the ionization energy we can discern.

·        Within a group on the periodic table, the ionization energy decreases as you go down the group.

·        Within a period of the periodic table, the ionization energy increases as you go left to right. Deviations within a period are usually due to metastable electron configurations.

Ionization Energies of the Elements in the First 5 Periods of the Periodic Table

Multiple Ionization Energies

The first ionization energy is the lowest. Subsequent ionization energies are greater due to the attraction of the positively charged ion to the electron being removed.

Metals and Nonmetals and Ionization Energies

Metals always have a positive oxidation number which means they lose electrons when they react. This means we would expect metals to have a lower ionization energy than nonmetals.

Gaining Electrons

Under some circumstances, it is possible to get an atom to accept electrons.

Electron Affinity - the energy released when an atom accepts an electron.  Atoms that lose lots of energy (and go to a more stable state) when they accept an electron have a high electron affinity. The same factors that affect ionization energy affect electron affinity. We observe the same periodic trends in electron affinity on the periodic table as we do for ionization energy.

Metals and Nonmetals and Electron Affinity

Since metals tend to lose electrons, they have a low electron affinity. Nonmetals can accept electrons and have a higher electron affinity.

Both ionization energy and electronegativity apply to atoms by themselves. We're more interested in how atoms of one element behave in the presence of atoms of other elements.

Stable and Metastable Electron Configurations

Most elements are unstable in their uncombined state. To achieve stability, they either lose, gain or share electrons in order to make a stable electron configuration.

1) The stable octet, s²p⁶, is one particularly stable electron configuration. Elements in group VIIIA, the Noble Gases have this configuration. Elements either lose, gain or share electrons in order to reach this electron configuration.

i.e. 1    Sodium, Na has an electron configuration 1s² 2s² 2p⁶ 3s¹. When Na reacts, it loses one electron forming the sodium ion, Na⁺¹ with the stable octet configuration, 1s² 2s² 2p⁶.

i.e. 2    Fluorine, F has an electron configuration 1s² 2s² 2p⁵. When F reacts, it gains an electron forming the fluoride ion, F^-1 with the stable octet configuration 1s² 2s² 2p⁶.

Notice that both Na⁺¹ and F^-1 both have an outer electron configuration similar to Ne, a Noble Gas.

2) A full or empty sublevel electron configuration gets an atom into a metastable (somewhat stable) electron configuration. In this case, all of the orbitals of that sublevel are either completely filled with electrons or empty.

i.e.

A filled s sublevel would be s². A full d sublevel would be d¹⁰.

An empty p sublevel would be p⁰ and an emptied d sublevel would be d⁰.

3) A half filled orbital (or half emptied for you pessimists out there!) is metastable. All of the orbitals in the sublevel are filled with one electron each.

i.e.

An s sublevel half filled would be s¹. A p sublevel half filled would be p³. A half filled d sublevel would be d⁵.

These metastable electron configurations explain:

1)     deviations from the expected electron configuration in Cr ...4s¹ 3d⁵ as opposed to ...4s² 3d⁴.
                                                                                                   

2)      Cu ...4s¹ 3d¹⁰ as opposed to ...4s² 3d⁹

       3) Some oxidation states like C⁺²        1s² 2s² 2p⁰
                                                           Cl⁺⁵       1s² 2s² 2p⁶ 3s² 3p³

Revisiting Radii: Ionic Vs. Atomic

If an atom loses or gains electrons, the resulting ion will have a different radius.

If an atom loses electrons, electrons are removed from the electron cloud and in some cases, the outer electrons are completely removed. The resulting positive ion will have a smaller radius than the atom. (r₊ < r₀)

i.e. r_Na+1 < r_Na

If an atom gains electrons, the additional electrons will cause even more repulsion and the electron cloud will increase in size. (r_- > r₀)

i.e. r_Cl-1 > r_Cl

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