In this lesson, students will learn about the Periodic Table, including how to distinguish between periods and groups, and understand the Periodic Law. The efforts of 19th-century chemists to systematically arrange elements led to the development of the Periodic Table, which organized known elements based on their properties and predicted the existence of undiscovered elements. In this table, vertical columns are called groups, while horizontal rows are known as periods.
Students will classify elements into these two categories based on the configuration of their outermost electrons, identify the demarcation of the table into s-block and p-block elements, and explain the overall shape of the Periodic Table. They will also explore the concept of families within the table and recognize the similarities in the physical and chemical properties of elements in the same family. Additionally, the lesson will cover how electronic configuration relates to an element’s position in the table, the influence of the shielding effect on periodic trends, and how electronegativities change within a group and across a period.
Periodic Table
The discovery of the periodic table significantly simplified the study of individual properties of known elements by grouping them in a systematic way. Various attempts were made to classify elements into a tabular form.
Dobereiner’s Triads: German chemist Dobereiner observed that in groups of three elements, known as triads, the atomic mass of the middle element was the average of the other two. An example is the triad of calcium (40), strontium (88), and barium (137). However, only a few elements could be arranged this way, and this classification did not gain wide acceptance.
Newlands Octaves: After Cannizzaro accurately determined atomic masses in 1860, British chemist Newlands proposed the “law of octaves” in 1864. He noted that the chemical properties of every eighth element repeated when arranged by increasing atomic mass, similar to musical notes. However, his work was not widely recognized because it did not accommodate undiscovered elements, such as the noble gases.
Mendeleev’s Periodic Table: Russian chemist Mendeleev arranged 63 known elements by increasing atomic mass in horizontal rows called periods, placing elements with similar properties in vertical columns. This arrangement became known as the Periodic Table, and he formulated the periodic law: “properties of the elements are periodic functions of their atomic masses.” Despite being the first significant attempt to organize the elements, Mendeleev’s table had limitations, such as not explaining the position of isotopes and incorrectly ordering some atomic masses.
Periodic Law: In 1913, H. Moseley discovered the concept of atomic number, which he determined should dictate an element’s position in the periodic table. This led to the revised periodic law: “properties of the elements are periodic functions of their atomic numbers.” The atomic number corresponds to the number of electrons in a neutral atom, providing a basis for understanding electronic configurations.
Modern Periodic Table
The atomic number of an element is a more fundamental property than atomic mass for two main reasons: it increases regularly from one element to the next, and it is unique to each element. The discovery of atomic number in 1913 led to a revision of Mendeleev’s periodic law, which was originally based on atomic mass. The modern periodic table arranges elements according to increasing atomic number. When elements are organized this way, their properties repeat at regular intervals, allowing elements with similar characteristics to be grouped together.
For example, after every eighth element, the ninth element exhibits similar properties to the first. Sodium (Z=11) shares properties with lithium (Z=3), and after atomic number 18, every nineteenth element shows similar behavior. This organization allowed for the creation of a table with vertical and horizontal rows.
The long form of the periodic table emphasizes the significance of atomic number, as it reflects the electronic configuration of elements. The horizontal rows, known as periods, feature elements with continuously increasing atomic numbers and changing electronic configurations, leading to varying properties. Elements are positioned within periods based on the number of valence electrons; for instance, alkali metals with one valence electron occupy the leftmost position, while noble gases with eight valence electrons are found on the right.
The vertical columns, called groups, are numbered from 1 to 18. While the atomic numbers in a group increase with irregular gaps, elements within the same group share similar electronic configurations and have the same number of valence electrons, resulting in similar chemical properties.
Salient Features of the Long Form of the Periodic Table:
- It contains seven horizontal rows (periods).
- The first period has two elements; the second and third periods each have eight elements; the fourth and fifth periods contain 18 elements each; the sixth period has 32 elements, while the seventh is incomplete with 23 elements.
- Elements in a period exhibit different properties.
- There are 18 vertical columns (groups) numbered from left to right.
- Elements in a group exhibit similar chemical properties.
- Elements are classified into four blocks based on their electron configurations.
Periodicity of Properties
Atomic Size and Atomic Radius
Atoms are extremely small and lack defined boundaries, making it challenging to measure their size directly. To estimate atomic size, we assume that atoms are spherical and that they touch when they are close together. The atomic radius is defined as half the distance between the nuclei of two bonded atoms. For example, in elemental carbon, the distance between the nuclei of two carbon atoms is 154 picometers (pm), which means the radius of a carbon atom is 77 pm.
As we move from left to right across a period in the periodic table, the size of atoms decreases even though the atomic number increases. This decrease in size occurs because the effective nuclear charge (the attraction between the nucleus and the outer electrons) increases with the addition of protons, while the electrons are added to the same valence shell. This increasing nuclear charge pulls the outermost shell closer to the nucleus, resulting in a smaller atomic size. For instance, the atomic size decreases from lithium (152 pm) to neon (69 pm) in period 2.
Conversely, atomic size increases as we move down a group because each successive element has an additional electron shell, which decreases the effective nuclear charge felt by the outermost electrons.
Shielding Effect
The shielding effect refers to the phenomenon where inner electrons reduce the nuclear charge experienced by the outermost shell electrons. This reduction occurs because the presence of inner electrons partially shields the outer electrons from the full attractive force of the nucleus. As a result, the valence electrons experience a lower effective nuclear charge than the actual charge from the nucleus.
As atomic number increases, the number of electrons also rises, which enhances the shielding effect. This effect becomes more significant as you move down a group in the periodic table. For example, it is easier to remove an electron from potassium (Z=19) than from sodium (Z=11) due to the increased shielding effect in potassium.
