The Wonders of the Periodic Table
Upon completion of this lesson, the students should be able to:
1. list the characteristics of the modern periodic table
2. classify the elements in the periodic table
3. explain the periodicity of elements
explain the periodicity of elements
Development of Periodic Table
As early as 1800, chemists began to determine the atomic weights of some elements with fair accuracy. Several attempts were made to classify the elements on this basis.
1. Johann Wolfgang Dobereiner (1829)
He classified the elements in groups of 3 called triads, based on similarities in properties and that the atomic mass of the middle member of the triad was approximately the average of the atomic masses of the lightest elements.
2. John A. New Lands (1863)
He arranged the elements in the order of increasing atomic mass. The eight elements starting from a given one is a kind of repetition of the first like the eight notes of the octave of music and called it the law of octaves.
3. Lothar Meyer
He plotted a graph showing an attempt to group elements according to atomic weight.
4. Dmitri Mendeleyeev (1869)
He worked out a Periodic Table of Elements were the elements were arranged in the order of increasing atomic weights with a regular repetition (periodicity) of physical and chemical properties.
5. Henry Moseley (1887)
He arranged the elements in the order of increasing atomic numbers, which relates that the properties of the elements are periodic functions of their atomic numbers. This is known as the Modern Periodic Law.
What are periods, groups and families?
Periods are the 7 horizontal rows in the periodic table
- Period 1 has 2 elements corresponding to 2 electrons in the s sublevel.
- Periods 2 and 3 have 8 elements corresponding to 8 sublevel electrons in the s and p sublevels.
- Periods 4 and 5 have 18 elements corresponding to 18 electrons in the s,p and d sublevels.
- Periods 6 and 7 also include the 14 f electrons but the seventh period is incomplete.
Groups are the vertical columns in the periodic table, which are divided into A, and B subgroups. The A subgroups are often called families. Some of the A families are designated by the following:
a. Group IA - Alkali Metals
b. Group IIA - Alkaline Earth Metals
c. Group VIIA - Halogens
d. Group VIIIA - Noble Gases
Other A subgroups are classified according to the first element in the column:
a. Group IIIA - Boron Family
b. Group IVA - Carbon Family
c. Group VA - Nitrogen Family
d. Group VIA - Oxygen Family
Classification of Elements in the Periodic Table
1. Representative Elements are the elements in A Group/ Family. The term representative element is related to stepwise addition of electrons to the s and p sub levels of the atoms. Elements belonging to the same group or family have similar properties.
2. Noble Gases or Inert Gases are the elements in the last group with completely filled set of s and p orbitals.
3. Transition Elements are the elements in the columns IB - VIIIB which are called the B Group/Family. Take note that they start with IIB up to VIIB, which have 3 columns and then end with IB and IIB. These sequences, which contain 10 elements each, are related to the stepwise addition of the 10 electrons to the d sub level of the atoms. These elements are metallic-dense, lustrous, good conductor of heat and electricity and in the most cases are hard. They form the many colored compounds and form polyatomic ions like Mn04 and CrO4.
4. Inner Transition Elements are the 2 additional horizontal rows below composed of 2 groups of elements which were discovered to have similar characteristics as Lanthanum in the 6th period called Lathanoids (Rare Earth Metals) and Actinium (Heavy Rare Elements). The Lanthanoids are all metals while the Actinoids are all radioactive. All the elements after Uranium are produced artificially by nuclear reactions.
The Periodic Table and Electronic Configuration
The Concept of Valence
Elements within any group exhibit a characteristic valence. The alkali metals of group IA exhibit a valence of +1, since the atoms easily lose the one electron in the outer level. The halogen of Group VIIA has a valence of -1, since one electron is readily taken up. In general, atoms, which have less than 4 valence electron, tend to give up electron thus having a positive valence corresponding to the number of electrons lost. While atoms with more than 4 valence corresponding to the number of electrons gained.
12 Mg 1s2 2s2 2p6 3s2
Magnesium will give up its 2 valence electrons forming a +2 valence
8 O 1s2 2s2 2p4
Oxygen has 6 valence electron thus it will gain 2 electrons -2 valence Group VIIIA has a stable outer configuration of electrons (with 8 valence electrons) and would not be expected to give up or take up electrons. Thus, this group has a zero valence.
In the B series, the incomplete level contributes to valence characteristics. One or two electrons from an incomplete inner level may be lost in chemical change and added to one or two electrons in the outer level, which allows possibilities of valence among the transition elements.
26 Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6
Iron may exhibit valence of +2 by loss of the 2 outer electrons or a valence of +3 when additional electron is lost from the incomplete 3rd level.
Lewis Dot System: Kernel Notation and Electron Dot Notation
Metals, Nonmetals and Metalloids
Metals are at the left and in the center of the Periodic Table. About 80 elements are classified as metals including some form in every group except Groups VIIA and VIIIA. The atoms of metals tend to donate electrons.
Nonmetals are at the far right and toward the top of the Periodic Table. They are composed of about a dozen relatively common and important elements with the exception of Hydrogen. Atoms of non-metals tend to accept electrons.
Metalloids or borderline elements are elements that to some extent exhibit both metallic and nonmetallic properties. They usually act as electron donor with metals and electron acceptor with non-metals. These elements lie in the zigzag line in the Periodic Table.
Positions of metals, nonmetals and metalloids in the Periodic Table
Trends in the Periodic Table
The atomic radius is approximately the distance of the outermost region of electron charge density in an atom drops off with increasing distance from the nucleus and approaches zero at a large distance. Therefore, there is no sharply defined boundary to determine the size of an isolated atom. The electron probability distribution is affected by neighboring atoms, hence, the size of an atom may change from one condition to another as in the formation of compounds, under different conditions. The size of the atomic radius is determined on covalently bonded particles of elements as they exist in nature or are in covalently bonded compounds.
Going across any period in the Periodic Table, there is a decrease in the size of the atomic radius. Going from left to right, the valence electron are all in the same energy level or the same general distance from the nucleus and that their nuclear charge increased by one. Nuclear charge is the force of attraction being offered by the nucleus towards electrons. Therefore, the greater the number of protons, the greater is the nuclear charge and the greater is the over pull of the nucleaus on the electron.
Consider the atoms of Period 3:
Na 2e 8e 1e Mg 2e 8e 2e Al 2e 8e 3e
Consider the electronic configuration of Group IA elements:
Na 2e 8e 1e
K 2e 8e 8e 1e
Rb 2e 8e 18e 8e 1e
Cs 2e 8e 18e 18e 8e 1e
Fr 2e 8e 18e 18e 18e 8e 1e
Although the number of protons from top to bottom within the same group increases, the atomic size still increases due to the additional energy level as seen from the above illustration. Therefore, atomic size increases from top to bottom within the same group.
Atomic size and Periodic Table
When an atom losses or gains electron, it becomes a positively/negatively charge particle called ion.
Magnesium losses 2 electrons and becomes Mg+2 ion.
Oxygen gains 2 electrons and becomes 0-2 ion.
The loss of electrons by a metal atom results in a relatively large decrease in size, the radius of the ion formed is smaller than the radius of the atom from which it was formed. For nonmetals, when electrons are gained to form negative ions, there is a rather large increase in the size due to the repulsion of the electrons for one another.
Ionic size and Periodic Table
Ionization energy is the amount of energy required to remove the most loosely bound electron in a gaseous atom or ion to give a positive (+) particle of cation. The first ionization energy of an atom is the amount of energy required to remove the first valence electron from that atom. The second ionization energy of an atom is the amount of energy required to remove the second valence electron from the ion and so forth. The second ionization energy is always higher than the first, since an electron is removed from a positive ion, and the third is likewise higher than the second.
Going across a period, there is an increase in the ionization energy due to the removal of electron in each case is at the same level and there is a greater nuclear charge holding the electron.
Factors affecting the magnitude of the ionization potential:
- The charge of the atomic nucleus for atoms of similar electronic arrangement. The greater the nuclear charge, the greater the ionization potential.
- The shielding effect of inner electrons. The greater the shielding effect, the smaller the ionization potential.
- The atomic radius. As the atomic size decreases in atoms with the same number of energy levels, the ionization potential increases.
- The extent which the most loosely bound electron penetrates the cloud of inner electrons. The degree of penetration of electrons in a given main energy level decreases in the order of s>p>d>f. All other factors being equal, as in the given atom, it is harder to remove an (s) electron than a (p) electron, a p electron is harder than a (d) electron, and d electron is harder than an (f) electron.
Attractive force between the outer level electrons and the nucleus increases in proportion to the positive charge on the nucleus and decreases with respect to the distance separating the oppositely charged bodies. Outer electrons are not only attracted by the positive nucleus but are also repelled by electrons in the lower energy levels and their own level. This repulsion, which has the net result of reducing the affective nuclear charge, is called the shielding effect or screening effect. Since from top to bottom, ionization energy decreases in A family, the screening effect and distance factors must outweigh the importance of the increased charge of the nucleus.
Ionization Energy and Periodic Table
Electron affinity is the energy given off when a neutral gaseous atom or ion takes in an electron. Negative ions or anions are formed. Determining electron affinities is a hard task; only those for the most nonmetallic elements have been evaluated. A second electron affinity values would involve gain and not loss of energy. An electron added to a negative ion would result in Coulombic repulsion.
0 + e- -------------- 0-1 -33 Kcal/mole
0 + e- -------------- 0-2 +189 Kcal/mole
This periodic trends of electron affinity, of the strongest nonmetals, the halogens, are due to their electron configuration, ns2 np5 that lack a p orbital to have stable gas configuration. Nonmetals tend to gain electrons to form negative ions than metals. Group VIIA has the highest electron affinity since only one electron is needed to complete a stable outer configuration of 8 electrons.
Electron Affinity and Periodic Table
Electronegativity is the tendency of an atom to attract shared electrons to itself when it forms a chemical bond with another atom. Ionization potential and electron affinities are regarded as more or less expressions of electronegativities. Atoms with small size, high ionization potential and high electron affinities would be expected to have high electronegativities Atoms with orbitals nearly filled with electrons will have higher expected electronegativities than atoms with orbitals having few electrons.Non metals have higher electronegativities than metals. Metals are more of electron donors and non metals are electron acceptors. Electronegativity increases from left to right within a period and decreases from top to bottom within a group.
Electronegativity and Periodic Table
Summary of the Trends in the Periodic Table
Readings on Periodic Table
- Periodic Properties of the Elements
Learn about the periodic properties or trends in the periodic table of the elements.
Video on Periodic Table
Self - Progress Test
A. I. Based on the given IUPAC Periodic Table and hypothetical elements as positioned, answer the following:
1. The most metallic element.
2. The most nonmetallic element.
3. The element with the biggest atomic size.
4. The element/s classified as alkali metal/s.
5. The element/s classified as metalloids.
6. The element/s classified alkali-earth metals.
7. The transition element/s.
8. The element/s classified as halogens.
9. The lightest of the noble gas.
10. Element/s with electronic configuration/s ending in d.
11. Element/s with electronic configuration ending in f.
12. Element/s with two (2) valence electrons.
13. Element/s with six (6) valence electrons.
14. Element/s with eight (8) valence electrons.
15. Element/s with one main energy level.
II. Answer fully the following questions:
1. State the Periodic Law.
2. Explain clearly what is meant by the statement that the maximum possible number of electrons in the outermost energy level is eight.
3. What are transition elements? How do you account for the marked differences in their properties?
B. Copy and fill in the table below:
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