ATOMIC STRUCTURE

THE PERIODIC TABLE AND PERIODIC PROPERTIES

Information collected from

http://gaia.floyd.edu/tutor/Periodicity.htm#Radii

http://antoine.frostburg.edu/chem/senese/101/periodic/index.shtml

http://chemed.chem.purdue.edu/genchem/topicreview/

Development of the Periodic Table

Mendeleev, a Russian Chemist, was one of the first to be partially sucessful in arranging the known elements in the 1870's into a chart that would allow the prediction of properties. He arranged the elements known in those days according to increasing atomic masses. The first Periodic Law proposed by him stated:

"The properties of the elements are a periodic function of their atomic masses"

When Mendeleev arranged the known elements using this principle as a guide, he found that certain elements grouped themselves into vertical groupings sometimes called families. If one measured a property of each of the elements in the group, Mendeleev noted that the value for that property would either be high or low depending upon the group under observation. For example, measuring the atomic radius of the atoms of elements which is the distance from the outermost valence region to the nucleus of an atom, one would find that in the case of the first group of elements on the left side of the chart all the elements seemed to have characteristically high values where those elements grouped on the right side of the chart seemed to have characteristically small radii values. All of the elements in group 1 had the highest radius value of any element in its respective period. A period is a horizontal row of elements in the table. Likewise all the elements in group 17 on the far right of the chart had the lowest radius value of any element in its respective period. This is a manifestation of the periodicity exhibited by the elements between their atomic weights and their property measurements. One could conceivably show this periodicity using other properties such as boiling points, melting points, etc.

You may access a short biography of Demitri Mendeleev here.

There were some inconsistencies in the arrangement of the elements according to his law, however it wasn't until the early 1900's (1914) that a Prof Moseley, a British Physicist, was able to determine the atomic numbers of all the known elements using an experimental technique. Moseley then proceeded to rearrange the elements according to increasing atomic numbers. Moseley's arrangement seemed to clear up the contradictions and inconsistencies of Mendeleev's arrangement, but Moseley based his arrangement on atomic numbers and not atomic masses.

Moseley's periodic law is now considered the current Periodic Law. It resulted in a slight alteration of Mendeleev's arrangement, but the slight difference was enough to correct the inconsistencies that existed in Mendeleev's arrangement.

The elements are arranged in vertical columns known as Groups. The elements in each group have consistently high or low values for certain properties. The horizontal rows of elements are referred to as "periods".

Group 1 is also called the alkali metal group. These are strong metals that are unusually soft and very reactive toward Oxygen forming Oxides and water forming hydroxides of the metal. These elements are so reactive toward Oxygen and water vapor that they are stored under an inert liquid to protect them from Oxygen and water vapor.

Group 2 is called the alkaline earth metals. These metals are not as soft as Group 1 metals. They also react more mildly with Oxygen to produce oxides of the metals and only react with water at temperatures where the water is steam.

Groups 3-12 are referred to as the transition metal groups. These metals are not as predictable because of the shielding effect of the inner electrons. As for the "shielding effect" this refers to the inner electrons found in the transition state elements and the inner transition (rare earth) elements. These electrons had a tendency to block the electrical effect of the positive nucleus upon the outer valence electrons of those atoms. This shielding effect helps to partially explain the erratic placement of the electrons in the d and f orbitals relative to the s and p orbitals.

Groups 1-2 and 13-18 are referred to as the representative elements

Group 17 is referred to as the halogen group

Group 18 is referred to as the Noble gas group previously known as the inert gas group.

There are two special series of elements that occur right after the transition metal element Actinium (Actinides) and Lanthanum(Lathanides). These special inner transition state metals were first rearranged by Dr. Glen Seaborg of Univ. of Calif. at Berkeley in the 1950's. It caused quite a stir until it was pointed out and demonstrated that this arrangement seemed to predict the properties of several newly synthesized man-made elements. I would call Dr. Seaborg's work the third milestone in our quest to make order out of the behavior of elemental substances. You may access a brief biography of Glen Seaborg here.

The metals which tend to have their atoms losing electrons during a chemical change are roughly found to the left Group 14

Non-metals which tend to have their atoms gaining electrons during chemical change are roughly found in Group16-17 with some elements in the lower parts of Groups 15.

Metalloids which tend to have their atoms sometimes losing and sometimes gaining electrons during chemical change are generally found in Groups 14-16

The Noble gases really belong to their own category since their atoms tend neither to lose or gain electrons. There are only a handful of compounds involving the Noble Gases (mostly involving Xenon).

Periodic Table of the Elements

lanthanides

actinides

PERIODIC PROPERTIES

►Quantum numbers and the periodic table

An element's location on the periodic table reflects the quantum numbers of the last orbital filled
The period indicates the value of principal quantum number for the valence shell

the lanthanides and actinides are in periods 6 and 7, respectively.

The block indicates value of azimuthal quantum number () for the last subshell that received electrons in building up the electron configuration.

blocks are named for subshells (s, p, d, f)
Each block contains a number of columns equal to the number of electrons that can occupy that subshell

§ The s-block (in orange) has 2 columns, because a maximum of 2 electrons can occupy the single orbital in an s-subshell.

§ The p-block (in purple) has 6 columns, because a maximum of 6 electrons can occupy the three orbitals in a p-subshell.

§ The d-block (in green) has 10 columns, because a maximum of 10 electrons can occupy the five orbitals in a d-subshell.

§ The f-block (in dark blue) has 14 columns, because a maximum of 14 electrons can occupy the seven orbitals in a f-subshell.

questions to ponder

What would the periodic table look like in a hypothetical universe where:

§ there were 3 possible values of m_s, instead of 2?

§ the angular momentum quantum number could take on values from 1 to n-1 only?

§ values of m = 0 were not allowed?

§ the maximum value of n were 5?

►Factors affecting the valence shell

Factors affecting the valence shell. Anything that influences the valence electrons will affect the chemistry of the element.
	Factors	Effect
1.	valence principal quantum number n	Larger n means a larger valence shell (because n controls the size of orbitals)
2.	nuclear charge Z	Larger Z means a smaller valence shell (because higher positive charge on the nucleus attracts the valence electrons, and pulls them inward)
3.	number of core electrons	More core electrons means a larger valence shell (because highly penetrating core electrons repel valence electrons, and push them farther from the nucleus)

►Effective nuclear charge

Electrons moving across the nucleus do not experience the same nuclear attraction; those electrons closer to the nucleus experience a greater force than those that are farther away. The nuclear charge actually "felt" by an electron is called the effective nuclear charge, Zeff. Zeff for a given electron is given by the true nuclear charge, Z, less the amount by which electrons closer to the nucleus screen it, S, Zeff = Z - S. Therefore, effective nuclear charge increases from left to right in a period and from top in a group on the periodic table.

Zeff To the right.

Li: 1s² 2s¹ => Zeff = +3 - 2 = 1

Cs > Li > Cl > F

The greater the effective nuclear charge, the greater the attractive force between the nucleus and its electrons (F q q/d²)

►Atomic Radii

To the left. As the attractive force between a nucleus and its electrons increases, the average distance between the nucleus and its electrons decreases. The average distance between the nucleus and its outermost electron is expressed as the atomic radius of the atom. It can be said that atomic radius decreases from left to right in a period and from bottom to top in a group on the periodic table. The greater the force of attraction, the smaller the radius.

trend	valence n	Z	# core electrons	net effect on atomic radius
going right across main group rows...	no change	increases	no change	the increase in Z causes a decrease in radius
going right across transition series...	no change	increases	increases	the increase in Z causes a decrease in radius, but the increase in the number of core electrons causes an increase. The two competing effects cause a small decrease, then small increase!
going down groups...	increases	increases	increases	three competing effects; but n is strongest, so radius increases.

►Sizes of Ions

Recall that atoms increase in size going from right-to left on a period and top-to-bottom in a group.

Cations are smaller than their parent atom because the effective nuclear charge on the outermost electrons is greater in the cation. The number of protons remains the same but the number of screening electrons decreases. Li > Li⁺

Z_li = 3 - 2 = 1

Z_li+ = 3 - 0 = 3

Anions are larger than their parent atoms because the effective nuclear charge on the outermost electrons in smaller in the anion. The number of protons remains the same but the number of screening electrons increases.

For ions of the same charge, size increases going down a group.

Isoelectronic series are groups of atoms and ions which have the same electronic configuration. Within isoelectronic series, the more positive the charge, the smaller the species and the more negative the charge, the larger the species.

1s² 2s²p⁶

N^3- O^2- F^- Na⁺ Mg²⁺ Al³⁺

Isoelectronic - same number of electrons

►Ionization Energy

ionization energy is the minimum amount of energy required to remove an electron from an atom or ion in the gas phase
normally, ionization removes valence electrons first

valence electrons are farthest from nucleus on average, so they feel the least attraction for the nucleus and are easiest to remove
end of valence electrons is marked by a big jump in ionization energies

Na(g) Na⁺(g) + e^-	H = +496 kJ	first ionization energy
Na⁺(g) Na⁺²(g) + e^-	H = +4560 kJ	second ionization energy

core orbitals have lower n, and are much more penetrating than valence orbitals
proximity to nucleus makes core electrons much more difficult to remove
core noble gas configurations have high stability

factors affecting ionization energy

atomic radius

smaller atoms hang on to valence electrons more tightly, and so have higher ionization energy

charge

the higher the positive charge becomes, the harder it is to pull away additional electrons
second ionization energy is always higher than the first

orbital penetration

It's easier to remove electrons from p orbitals than from s orbitals

electron pairing

within a subshell, paired electrons are easier to remove than unpaired ones
reason: repulsion between electrons in the same orbital is higher than repulsion between electrons in different orbitals
example
On the basis of gross periodic trends, one might expect O to have a higher ionization energy than N. However, the ionization energy of N is 1402 kJ/mol and the ionization energy of O is only 1314 kJ/mol. Explain.

Taking away an electron from O is much easier, because the O contains a paired electron in its valence shell which is repelled by its partner.

There are periodic trends in the ionization energies, also tied to the effective nuclear charge. As the effective nuclear charge increases, it requires more energy to remove the outermost electron from an atom. Consequently, ionization energy is also related to the atomic radius, with ionization energy increasing as atomic radius decreases. Therefore, the first ionization energy increases from left to right in a period and from bottom to top in a group.

Na > Al > Mg > Si

Diagram

Exceptions to the General Pattern of First Ionization Energies

The figure below shows the first ionization energies for elements in the second row of the periodic table. Although there is a general trend toward an increase in the first ionization energy as we go from left to right across this row, there are two minor inversions in this pattern. The first ionization energy of boron is smaller than beryllium, and the first ionization energy of oxygen is smaller than nitrogen.

These observations can be explained by looking at the electron configurations of these elements. The electron removed when a beryllium atom is ionized comes from the 2s orbital, but a 2p electron is removed when boron is ionized.

Be: [He] 2s²

B: [He] 2s² 2p¹

The electrons removed when nitrogen and oxygen are ionized also come from 2p orbitals.

N: [He] 2s² 2p³

O: [He] 2s² 2p⁴

But there is an important difference in the way electrons are distributed in these atoms. Hund's rules predict that the three electrons in the 2p orbitals of a nitrogen atom all have the same spin, but electrons are paired in one of the 2p orbitals on an oxygen atom.

Diagram

Hund's rules can be understood by assuming that electrons try to stay as far apart as possible to minimize the force of repulsion between these particles. The three electrons in the 2p orbitals on nitrogen therefore enter different orbitals with their spins aligned in the same direction. In oxygen, two electrons must occupy one of the 2p orbitals. The force of repulsion between these electrons is minimized to some extent by pairing the electrons. There is still some residual repulsion between these electrons, however, which makes it slightly easier to remove an electron from a neutral oxygen atom than we would expect from the number of protons in the nucleus of the atom.

Second, Third, Fourth, and Higher Ionization Energies

By now you know that sodium forms Na⁺ ions, magnesium forms Mg²⁺ ions, and aluminum forms Al³⁺ ions. But have you ever wondered why sodium doesn't form Na²⁺ ions, or even Na³⁺ ions? The answer can be obtained from data for the second, third, and higher ionization energies of the element.

The first ionization energy of sodium, for example, is the energy it takes to remove one electron from a neutral atom.

Na(g) + energy Na⁺(g) + e^-

The second ionization energy is the energy it takes to remove another electron to form an Na²⁺ ion in the gas phase.

Na⁺(g) + energy Na²⁺(g) + e^-

The third ionization energy can be represented by the following equation.

Na²⁺⁽g) + energy Na³⁺(g) + e^-

The energy required to form a Na³⁺ ion in the gas phase is the sum of the first, second, and third ionization energies of the element.

First, Second, Third, and Fourth Ionization Energies of Sodium, Magnesium, and Aluminum (kJ/mol)

chart of ionization energies

It doesn't take much energy to remove one electron from a sodium atom to form an Na⁺ ion with a filled-shell electron configuration. Once this is done, however, it takes almost 10 times as much energy to break into this filled-shell configuration to remove a second electron. Because it takes more energy to remove the second electron than is given off in any chemical reaction, sodium can react with other elements to form compounds that contain Na⁺ ions but not Na²⁺ or Na³⁺ ions.

A similar pattern is observed when the ionization energies of magnesium are analyzed. The first ionization energy of magnesium is larger than sodium because magnesium has one more proton in its nucleus to hold on to the electrons in the 3s orbital.

Mg: [Ne] 3s²

The second ionization energy of Mg is larger than the first because it always takes more energy to remove an electron from a positively charged ion than from a neutral atom. The third ionization energy of magnesium is enormous, however, because the Mg²⁺ ion has a filled-shell electron configuration.

The same pattern can be seen in the ionization energies of aluminum. The first ionization energy of aluminum is smaller than magnesium. The second ionization energy of aluminum is larger than the first, and the third ionization energy is even larger. Although it takes a considerable amount of energy to remove three electrons from an aluminum atom to form an Al³⁺ ion, the energy needed to break into the filled-shell configuration of the Al³⁺ ion is astronomical. Thus, it would be a mistake to look for an Al⁴⁺ ion as the product of a chemical reaction.

►Electron Affinities

The electron affinity of an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion: M₍_g) + e M_(g)^-. These reactions tend to be exothermic and so the values of E are generally negative.

Diagram In general, electron affinity tends to decrease (become more negative) from left to right in a period. Going down a group, there is little change in the electron affinities. Negative electron affinity means that the atom gains electrons easily. Cl > Na > N > Be

Several patterns can be found in these data.

Electron affinities generally become smaller as we go down a column of the periodic table for two reasons. First, the electron being added to the atom is placed in larger orbitals, where it spends less time near the nucleus of the atom. Second, the number of electrons on an atom increases as we go down a column, so the force of repulsion between the electron being added and the electrons already present on a neutral atom becomes larger.
Electron affinity data are complicated by the fact that the repulsion between the electron being added to the atom and the electrons already present on the atom depends on the volume of the atom. Among the nonmetals in Groups VIA and VIIA, this force of repulsion is largest for the very smallest atoms in these columns: oxygen and fluorine. As a result, these elements have a smaller electron affinity than the elements below them in these columns as shown in the figure below. From that point on, however, the electron affinities decrease as we continue down these columns.

At first glance, there appears to be no pattern in electron affinity across a row of the periodic table, as shown in the figure below.

When these data are listed along with the electron configurations of these elements, however, they make sense. These data can be explained by noting that electron affinities are much smaller than ionization energies. As a result, elements such as helium, beryllium, nitrogen, and neon, which have unusually stable electron configurations, have such small affinities for extra electrons that no energy is given off when a neutral atom of these elements picks up an electron. These configurations are so stable that it actually takes energy to force one of these elements to pick up an extra electron to form a negative ion.

►Electronegativity

Relative tendency of an atom to attract electrons to itself when chemically combined with another atom.

The relative ability of an atom to draw electrons in a bond toward itself is called the electronegativity of the atom. Atoms with large electronegativities (such as F and O) attract the electrons in a bond better than those that have small electronegativities (such as Na and Mg). The electronegativities of the main group elements are given in the figure below.

When the magnitude of the electronegativities of the main group elements is added to the periodic table as a third axis, we get the results shown in the figure below.

There are several clear patterns in the data in the above two figures.

Electronegativity increases in a regular fashion from left to right across a row of the periodic table.
Electronegativity decreases down a column of the periodic table.

►Why metals are metals

the ionization energy of metallic elements is very low
valence electrons are easily lost, and shared among all atoms in the metal
this 'sea' of valence electrons binds together the metal cations and gives metals their characteristic properties

mobility of electrons in the sea explains metal's ability to conduct electricity and heat
metals are workable because cations can slide past each other but still be bound by the electron sea

comparing metals

more valence electrons means stronger metal
higher positive charge on cations, higher negative charge on sea = stronger bonding

Explaining elemental properties: the s block elements

The properties of the alkali metals ultimately result from their ns¹ valence configuration.
property of alkali metals	explanation
metallic	very low ionization energy; the electron sea model works well for alkali metals
soft	ns¹ valence configuration contributes just 1 electron to the electron sea. The sea is weak. Metal cations aren't tightly bound and it's easy to slide them past each other.
low densities	Alkali metals have the largest radii and lowest atomic weight in each period. Low mass in high volume = low density.
highly reactive	very low ionization energies make alkali metals good electron donors in redox reactions.

http://gaia.floyd.edu/tutor/Periodicity.htm#Radii

http://antoine.frostburg.edu/chem/senese/101/periodic/index.shtml

http://chemed.chem.purdue.edu/genchem/topicreview/