CHM 6

CHM 010 Chapter 6: The Periodic Table

Learning Goals

1. Discuss the historical basis and significance of the periodic table.

2. Distinguish between a period and a group in the periodic table.

3. State the difference between an A-group (representative) element and a B-group element (transition metal).

4. Predict trends for properties such as atomic radius, ionization potential, and electron affinity.

The Modern Periodic Table

The modern periodic table consists of seven horizontal rows known as periods and 18 vertical columns called groups (or families). Every atom of elements in a given group has the same number of electrons in their outermost shells, leading to similar chemical behavior. For instance, elements in Group IA react violently with water due to having one electron in their outermost shell.

Historical Development

Early Knowledge of Elements

• From ancient Greek times until 1800, only 14 elements were known.

• From 1800 to 1810, an additional 14 elements were discovered.

• By 1830, this number increased to 45 known elements.

• As new elements emerged, chemists began to ponder how many elements actually existed.

Johann Döbereiner (1829)

• The German chemist Johann Döbereiner noted patterns of similar properties among elements. He observed that bromine possessed characteristics between chlorine and iodine, with its atomic mass positioned almost midway between them.

• To establish order, Döbereiner identified groups of three elements that shared chemical properties, calling these groups triads.

  • Triad Examples:

  • First Triad:

    • Calcium (Ca)

    • Strontium (Sr)

    • Barium (Ba)

  • Second Triad:

    • Sulfur (S)

    • Selenium (Se)

    • Tellurium (Te)

• The following table illustrates the properties of these triad elements:

John Newlands (1864)

• The English chemist John Newlands arranged the elements by increasing atomic mass, discovering a recurring pattern where every eighth element exhibited properties similar to the first. This observation was termed the Law of Octaves. However, Newlands' arrangement faced criticism due to the inconsistencies it presented, especially beyond calcium, leading to its rejection by the scientific community.

Dmitri Mendeleev and Lothar Meyer (1869)

• Independently, they proposed a more systematic tabulation based on the recurring properties of elements, known as periodicity.

• Mendeleev created an extensive periodic table of 56 elements, positioning them primarily by atomic mass and sometimes adjusting for chemical properties.

• He established the concept of periods, starting with hydrogen, followed by periods of seven elements.

Mendeleev's Adjustments

• Mendeleev occasionally prioritized elements based on chemical similarities, even placing heavier elements before lighter ones. He anticipated the existence of certain elements by leaving gaps in his table, predicting properties based on other known elements. For example:

  • Eka-aluminum (Gallium) was predicted by Mendeleev, and upon its discovery six years later, it closely matched Mendeleev's expectations.

Inconsistencies of Mendeleev's Table

• Mendeleev's reliance on atomic mass caused some inaccuracies. For instance, argon and potassium would switch places if ordered by atomic mass, contradicting the periodicity of properties.

Modern Periodic Table

• Today, the periodic table is organized by atomic number rather than atomic mass, leading to a better reflection of periodicity. The groups are categorized as A groups (representative elements) and B groups (transition metals).

Classification of Elements

• Elements are categorized based on the type of subshell containing the outermost electrons:

  • Main group elements

  • Noble gases

  • Transition elements (transition metals)

  • Lanthanides

  • Actinides

Periodicity and Electron Configuration

• A-group elements share the same number of outer shell electrons, reflected by their group number. For example, Group IA has one outer electron, while Group VIA has six.

• The outermost electrons are called valence electrons, crucial for chemical bonding. For instance, halogens (Group 7A) uniformly have outer electron configurations of ns²np⁵.

Group Numbers and Valence Electrons

• Group elements uniformly possess the same number of valence electrons.

• Example for Group 2A (2):

  • Beryllium (Be): 1s² 2s²

  • Magnesium (Mg): 1s² 2s² 2p⁶ 3s²

  • Calcium (Ca): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

Periodic Trends

Atomic Radius

• The atomic radius refers to the distance from the nucleus center to the outermost electron, measured in angstroms (1 Å = 10⁻¹⁰ m).

• Trends Resulting in Atomic Size:

  • Increases down a group due to more energy levels.

  • Decreases across a period due to increasing nuclear charge attracting valence electrons more strongly.

Ionization Energy

• Ionization energy is the energy required to remove an electron from a neutral atom. Generally, there is an increase in ionization energy across a period from left to right and a decrease down a group.

• Metals typically have lower ionization energies than nonmetals, with noble gases exhibiting the highest due to their complete octets.

Electron Affinity

• Electron affinity defines an atom’s ability to accept an additional electron, reflecting in the energy released upon acceptance. Electron affinity tends to increase across a period and decreases down a group, influenced by increasing nuclear charge and energy level distance.

Summary of Periodic Trends

Trends from Top to Bottom in a Group

• Atomic radius: Increases

• Ionization potential: Decreases

• Electron affinity: Decreases

Trends from Left to Right across a Period

• Atomic radius: Decreases

• Ionization potential: Increases

• Electron affinity: Increases

Properties of Elements

• Metals: Shiny, malleable, good thermal and electrical conductors; low ionization energies.

• Nonmetals: Variable in appearance, brittle, poor conductors;

• Metalloids: Intermediate properties between metals and nonmetals.