Trends in the Periodic Table

TRENDS IN THE PERIODIC TABLE

            Though the Periodic Law states that the periodicity of elements is based on increasing atomic number, the physical and chemical properties of elements are ultimately based on the electron configuration of its atoms. Some of the key properties of elements that can be studied in the periodic table are Electronegativity, Ionization energy, Electron Affinity, and Atomic radius.

                Two important factors must be discussed first before venturing into the periodic trend. These factors should be understood first in order for us to better understand why such trends occur. These are the (1) Principal Quantum Number (n) of the valence electrons and the (2) Effective Nuclear Charge (Z_eff) “felt” by the valence electrons.

                The Principal Quantum Number (n) tells us the average distance of the electron from the nucleus. As n increases, the outermost electrons called valence electrons gets farther away from the nucleus. The Effective Nuclear Charge (Z_eff) is the attraction of the valence electrons to the nucleus. This attraction is affected by the inner electrons that screens or shields the valence electrons from the nucleus. This is called the Shielding Effect. Because of the inner electrons, the valence electrons do not feel the full nuclear charge.

                                Z_eff< Z

              Z_eff = Z – S

                where: S is the Screening or Shielding factor;

                                 Z is the Nuclear Charge; and

                                Z_eff is the effective Nuclear Charge

The valence electron does not feel the full nuclear charge because the inner electrons screen or shield it from the nucleus

               The Effective Nuclear Charge can be likened to Family A with one child and to Family B with seven or eight children. The child in Family A can have more attention coming from his parents compared to the children in Family B.

                Electronegativity is defined as the ability of an atom to attract electrons towards itself. In other words, it is the electron appeal of an atom. The higher the electronegativity of an element, the more it can pull its own valence electrons and the valence electrons of other elements towards its nucleus. The trends for electronegativity of elements in the periodic table are as follows:

§  Increases from Left to Right (across the period)

Ø  This is due to the fact that the average distance of the valence electrons from the nucleus (n) are the same while the number of charge of the elements increases as you go along the period. In effect, the nuclear charge becomes stronger from left to right of the periodic table, making it attract more electrons towards itself.

§  Decreases from Top to Bottom (down the group)

Ø  This is because the average distance of the valence electrons from the nucleus (n) becomes greater, hence, making the effective nuclear charge of the element weaker. Because the effective nuclear charge is weak, its electronegativity becomes weaker also.

                 Noble gases have weak or no electronegativity values because they have completely filled s- and p-sublevels and attracting electrons toward itself would make it unstable. This explains why Noble Gases are unreactive to most substances. This also means that the most electronegative element is Fluorine and the least electronegative is Francium. Furthermore, this makes Fluorine the most reactive nonmetal and why most nonmetals are negatively charged.

                Ionization Energy is the energy absorbed by an atom to dislodge or remove its valence electrons. The trends for ionization energy of elements in the periodic table are as follows:

§  Increases from Left to Right (across the period)

Ø  Because the effective nuclear charge increases as you move from left to right of the periodic table, it would require less energy to dislodge electrons for the elements in the Alkali Metal Group while more energy is needed for elements in the Halogen Group.

§  Decreases from Top to Bottom (down the Group)

Ø  As n increases, the effective nuclear charge decreases. As effective nuclear charge decreases down the group, the lesser energy is required to dislodge electrons

This trend explains why metals are positively charged, they tend to give away their valence electrons because of weak effective nuclear charge and large atomic size. This also makes Francium the most reactive metal.

As more electrons are removed or dislodged from the atom, the greater the ionization energy becomes. This is because the number of negative charge becomes lesser, thus, the nucleus has a greater pull on the electrons, like in a tug-o-war where one team loses a member and the whole team gets pulled.

   

  

                Electron Affinity is defined as the energy released by an atom after gaining an electron. Generally, the greater the electron affinity of an atom, the more stable it becomes after gaining an electron. Electron Affinity can be likened to the excitement of atoms to have the same electron configuration as the noble gases, which are stable atoms. The greater their excitement, the more energy they release after gaining an electron. The trends for electron affinity of elements in the periodic table are as follows:

§  Increases from Left to Right (across the period)

Ø  A metal’s tendency is to give away their electrons so their excitement to receive electrons is low. A nonmetal’s tendency is to accept electrons (due to high effective nuclear charge), thus, it has high excitement to receive electrons, hence, high energy is release after gaining an electron.

§  Decreases from Top to Bottom (down the group)

Ø  Effective nuclear charge decreases down the group, thus, lesser excitement to receive electrons an lesser energy is released after gaining an electron.

Ø  Effective nuclear charge decreases down the group, thus, lesser

excitement to receive electrons and lesser energy is released after gaining an electron

This trend is only applicable for the first electron gained by a nonmetal. For the

second electron, electron affinity is positive, thus, energy is absorbed because, like in the case of O^-, energy would be needed to attach another electron to a negatively charged ion due to the electrostatic repulsion between the two.

                Atomic radius is the average distance between the nucleus and the valence electrons.

The trends for atomic radius of elements in the periodic table are as follow:

§  Decreases from left to right (across the period)

Ø  Though the principal quantum number (n) of the valence electrons across the group is the same, the effective nuclear charge increases from left to right thus, the valence electrons get attracted towards the nucleus making the atomic radius smaller

§  lncreases from top to bottom - (down the group)

Ø  The principal quantum number (n) increases from top to bottom, thus, the distance between the nucleus and the valence electrons becomes greater.

 Atomic Radius vs Ionic Radius

Atomic radius, as studied earlier, is the average distance between the nucleus and the valence electrons. When atoms lose or gain electrons, they become positive or negatively charged and become ions. The ionic radius is an estimated radius of an ion in an ionic compound. When an atom loses or gains electrons, the size of the particle suddenly changes.

When metals lose their valence electrons, the number of protons becomes greater than the number of electrons, thus, the effective nuclear charge becomes greater and the electrons get attracted towards the nucleus, making the ionic radius smaller. The tendency of metals is to accept electrons due to their strong electronegativity. When a nonmetal accepts an electron, the electron reduces the effective nuclear charge. The additional electron causes the other electrons to move away from each to relieve the repulsive force between them, making the ionic radius larger. As more electrons are added to the nonmetal ion, the larger the ion becomes.

To summarize the main points in the trend in ionic radius:

§  Increases from top to bottom

§  Decreases with increasing positive charge for metals along a period

§  Increases a charge becomes more negative for nonmetals along a period

Enhancing Skills:

Using your knowledge on the Periodic Trends, arrange the following elements:

A. ln lncreasing Electronegativity:

1. Ca, Ba, Sr, Be, Ra

2. Sb, P, N, Bi, As

3. Br, K, Ge, Ca, Se

B. ln Decreasing lonization Energy:

1. K, Na, Li, Rb, Cs

2. Br, F, Cl, l, At

3. Po, Bi, Pb, At Tl

C. ln Decreasing Atomic Radius

1, Mo, Nb, Tc, Ru, Zr

2. Cu, Fe, Zn, Ni, Co

3. B, ln, Tl, Ga, Al

Periodic Table and the Electron Configuration

PERIODIC TABLE AND THE ELECTRON CONFIGURATION

            The atomic number also tells us the number of electron of elements. The electron configuration explains how physical and chemical properties of elements recur in the periodic table.

   

                The electron configuration of elements in Groups 1A and 2A ends with s-sublevel, thus, called s-block while the elements in Groups 3A to 8A ends with p-sublevel, thus called the p-block. The elements from Groups 1A to 7A with incomplete s- or p-sublevels of the highest principal quantum number are collectively called as the Representative Elements. Elements in Group 8A have completely filled s- or p-sublevel, and are not classified as representative elements. These elements are called the Noble Gases. The electron configuration of elements between s- and p-blocks ends with d-sublevels, thus are called Transition Elements. Elements in Group 2B are sometimes called post-transition elements because their d-sublevels are completely filled. Elements with electron configurations ending with f sublevel are called Inner Transition Elements and are grouped in the f-block.

                Comparison between the Representative and Transition Elements are as follows:

Representative Elements	Transition Elements
Constant oxidation state Ions are usually colorless Ions are diamagnetic	Variable oxidation states Often form colored ions Ions are paramagnetic

Historical Development of the Periodic Table

5.1 INTRODUCTION

 

                As scientists were trying to unravel the nature of matter through their studies on the subatomic particles and the Quantum Theory, the search for new compounds still continue. From ancient times to 1899, 83 elements were already discovered. These elements needed to be arranged in such a way that they can be easily studied. Unlike the periodic table shown above, the elements were not organized and were difficult to study, thus, paved the way to the development of the Modern Periodic Table of Elements.

The modern periodic table is a chart where the elements are arranged according to increasing atomic number and are grouped according to similar physical and chemical properties. The physical and chemical properties of elements can be observed according to columns called Groups or Families or through rows called Periods.

 

o   Group – column of elements that have similar properties

o   Period – row of elements that tells us the valence shell of the element

Periods are numbered 1 to 7 while Groups are numbered 1 to 8, classified as either A or B. Groups are labeled in two ways: the number-letter symbol, which is commonly used by chemists, and the Group numbers from 1 to 18. For convenience, groups are also given special names that depict their physical and chemical properties, which will be discussed later in the chapter.

Elements are classified in to three types: Metals, Metalloids (semimetals), and Nonmetals.

o   Metals -  are good conductors of heat and electricity

- elements on the left side of the “staircase” on the periodic table

o   Nonmetals – are poor conductors of heat and electricity

   - elements on the right side of the “staircase” on the periodic table

o   Metalloids – elements having intermediate properties of metals and

nonmetals

Some of the elements classified as metalloids are Boron, Silicon, Germanium, Arsenic, Antimony, and Tellurium. In some books, Polonium and Astatine are considered as metalloids but are still under debate. Metalloids have valuable importance in the computer industry. Some metalloids do not conduct electricity at normal conditions, but when certain conditions are met like temperature or amount of electricity, they allow the flow of electricity, making them valuable components of the circuitry in computers.

5.2 HISTORICAL DEVELOPMENT OF THE PERIODIC TABLE

Since the age of antiquity, many elements were already known to man, some of which are gold, lead, silver, copper, tin, iron, and mercury. The first scientific discovery of an element was done by a German merchant Hennig Brand who discovered phosphorus by distilling human urine. This was an accidental discovery because he was actually looking for the Philosopher’s Stone, the mythical substance that could turn anything into gold, to save him from his bankruptcy. After that, many more scientists followed in discovering new elements.

The earliest attempt to arrange the elements was made by Johann Wolfgang Döbereiner, a German chemist who proposed the Law of Triads.

o   Law of Triads – a group of elements having the same properties where thesecond element of the triad has almost exactly the average atomic mass of the first and third.

                Below are examples of triads given by Döbereiner:

Element	Atomic Mass	Density
Chlorine Bromine Iodine	35.45 g/mol 79.90 g/mol 126.90 g/mol	0.0032 g/ml 3.1028 g/ml 4.933 g/ml
Calcium Strontium Barium	40.09 g/mol 87.62 g/mol 137.33 g/mol	1.55 g/ml 2.54 g/ml 3.59 g/ml

                After Döbereiner’s attempt, another scientist in the name of John Newlands proposed the Law of Octaves.

o   Law of Octaves – when elements are arranged in increasing atomic mass, the physical and chemical properties of elements tend to repeat after every eighth element or after an interval of 7 elements.

John Newlands likened his arrangement of the elements to the octaves in music as shown below:

1	2	3	4	5	6	7	8
do	re	mi	fa	sol	la	ti	do
Li	Be	B	C	N	O	F	Na
Na	Mg	Al	Si	P	S	Cl	K

Lithium in the “lower do” and Sodium in the “higher do” have the same physical and chemical properties. The same rule applies in the “lower notes” in Newlands’ arrangement. The problem with this arrangement was that the pattern was no longer applicable to elements beyond Calcium.

 

 As you can observe from his arrangement, after Fluorine and Chlorine, the elements Sodium and Potassium followed immediately. This is because during those times, scientists haven’t included the noble gases yet to the periodic table as a new class of elements. Newlands made his Law of Octaves in 1865 while Lord Rayleigh and William Ramsay proved the existence of the noble gases and got it included in Mendeleev’s periodic table in 1902.

Newlands’ Law of Octaves received negative feedbacks from his colleagues for likening his arrangement of the elements to the musical scale but later came into importance when Gilbert Lewis proposed his Valence Bond Theory and Irving Langmuir’s  with his Octet Rule.

The most successful attempt to arrange the elements was done by Dmitri IvanovichMendeleev. He was a Russian chemist who, like others before him, arranged the elements according to increasing atomic mass. He grouped the elements based on their physical and chemical properties and made some predictions on the properties of some elements that were not yet discovered during those times. A contemporary of Mendeleev who had the same idea as his was the German chemistJulius Lothar Meyer. He arranged the elements the same way Mendeleev did. But what made Mendeleev’s arrangement different was that he left gaps in his table, leaving spaces for future elements to be discovered. One example is his prediction of an unknown element that has almost the same properties as aluminum. He called this element as eka-aluminum (Ea)(eka- from Sanskrit which means “first”, meaning that eka-aluminum is the first element beneath Aluminum).

	Eka-Aluminum	Gallium
Atomic Mass	68 amu	69.9 amu
Melting Point	Low	30.15oC
Formula of Oxide	Ea2O3	Ga2O3
Density	5.9 g/cm3	121

                Though they had virtually the same arrangement of elements, the larger credit was also given to Mendeleev because of him taking the risk of predicting new elements and publishing his work a few months before Meyer did.

                Further refinements were made in Mendeleev’s periodic table. He noticed that there were anomalies in the properties of elements if the basis of the arrangement is the atomic mass. An example would be Nickel and Cobalt:

                 In the modern periodic table, Cobalt comes first followed by Nickel. But if you will look at the atomic masses of the two elements, Cobalt is heavier than Nickel, therefore, in Mendeleev’s periodic table, he placed Nickel before Cobalt. This arrangement posed a problem because the elements do not fit the properties of the other elements within the group. Mendeleev made a bold decision to switch the position of the two elements, thus, its current position. This problem can also be seen between ₁₈Ar (At. Mass = 39.948) and ₁₉K     (At. Mass = 39.098). This made scientists realize that the atomic mass of elements is not the basis of periodicity.

                This problem was later on resolved by Henry Moseley, a 25-year old English scientist who discovered the Atomic Number. This atomic number corresponds to the nuclear charge of the atom or the number of protons inside the nucleus. He made his discovery by bombarding metals with high energy electrons. Through his discovery of the atomic number, they were able to prove that Cobalt (Atomic number of 27) should come first before Nickel (Atomic number of 28) despite their atomic masses. This led to the formulation of the Periodic Law.

o   Periodic Law – states that the physical and chemical properties of elements are periodic functions of increasing atomic number.