Structure of an atom

Chapter 3: Atoms and Molecules

Fundamental Concepts

Atoms and molecules are the building blocks of matter. Different types of matter arise from different combinations of atoms.

Key Questions Addressed

• What distinguishes one element's atom from another?

• Are atoms indeed indivisible?

Subatomic Particles

Exploration of subatomic particles and their arrangement within atoms. The late 19th century focus was on determining the atom's structure through experiments. Early hints that atoms are divisible came from studies of electricity.

Charged Particles in Matter

Activities to Explore Charge:

• Charge comb dry hair and observe the attraction of small paper pieces.

• Rub a glass rod with a silk cloth and bring it near an inflated balloon.Conclusion: Rubbing objects together results in them becoming electrically charged, suggesting that atoms contain charged particles.

Discovery of Charged Particles

J.J. Thomson identified the electron in 1897, establishing the atom contained subatomic particles. Canal rays discovered by E. Goldstein in 1886 led to the identification of protons as another subatomic particle with a positive charge.

Subatomic Particle Characteristics

• Electrons (e–)

  • Charge: -1 (negative)

  • Mass: Negligible relative to protons.

• Protons (p+)

  • Charge: +1 (positive)

  • Mass: Considered one unit; approximately 2000 times the mass of an electron.

Understanding Atomic Structure

4.1 Emergence of Atomic Models

• Thomson's Model: Thomson proposed the atom as a positively charged sphere with electrons embedded within it, a model likened to a Christmas pudding or a watermelon. This "plum pudding model" emphasized charge neutrality and was the first to suggest that atoms contained smaller particles (electrons). However, it failed to explain certain experimental results related to the atom's structure and behavior, such as the distinct spectral lines of elements

• Rutherford's Model:

  • Alpha Particle Scattering Experiment: Conducted by Ernest Rutherford in 1909, this involved bombarding a thin gold foil with alpha particles to study the atomic structure.

  • Observations:

    • Most alpha particles passed through the foil unhindered.

    • A few were slightly deflected at small angles.

    • A minority bounced back at high angles, leading to surprising observations.

  • Conclusions Drawn:

    • Most of the atom is empty space.

    • Positive charge is concentrated in a small, dense nucleus, which contains most of the atom's mass.

  • Features of Rutherford's Model:

    • Proposed a model with a positively charged nucleus and electrons orbiting around it, similar to how planets orbit the sun.

    • Size of the nucleus is very small compared to the overall size of the atom, making it a concentrated area of mass.

  • Drawbacks: This model couldn't explain the stability of atoms since charged particles in motion would radiate energy, leading to a loss of energy and the spiral of electrons into the nucleus.

• Bohr's Model of the Atom: Developed by Niels Bohr in 1913, this model directly addressed the limitations identified in Rutherford's model. It introduced the idea of quantized energy levels (orbits) for electrons surrounding the nucleus.

  • Key Elements of Bohr's Model:

    • Electrons occupy distinct orbits with fixed energies without radiating energy.

    • Electrons can jump between energy levels by absorbing or emitting specific amounts of energy (photons).

    • This was significant in explaining the hydrogen emission spectrum and line spectra of other elements, which could not be accounted for by previous models.

4.4 Electron Shells

• Bohr-Bury Rules for Electron Distribution:

  • Maximum electrons in a shell = 2n², where n = shell number.

  • K-shell: 2 electrons

  • L-shell: 8 electrons

  • M-shell: 18 electrons

  • N-shell: 32 electrons.

• Outer Shell: Can contain a maximum of 8 electrons before filling inner shells.

• Valence electrons are those in the outermost shell, which dictate chemical reactivity.

4.5 Valency and Chemical Activity

• Valency: Determined by the number of electrons in the outer shell, it is crucial for understanding the reactive properties of elements. Atoms attain stability by reaching a full outer shell, commonly referred to as the octet rule.

  • Example: Sodium can lose one electron (valency = 1) to achieve a stable electronic configuration by resembling the nearest noble gas configuration.

Atomic Number and Mass Number

4.6 Definitions

• Atomic Number (Z): The number of protons within an atom's nucleus, which defines the element and determines its position on the periodic table.

• Mass Number (A): The total number of protons and neutrons in the nucleus, influencing the atomic mass and stability of the isotope forms of an element.

Isotopes and Isobars

4.7 Isotopes

Atoms of the same element that have the same atomic number but differ in mass numbers due to varying numbers of neutrons (e.g., Hydrogen has isotopes: Protium, Deuterium, Tritium).

4.8 Isobars

Different elements having the same mass number but distinct atomic numbers (e.g., Calcium and Argon).

Key Takeaways

• Discovery of subatomic particles revolutionized atomic theory.

• Thomson's Model introduced the concept of charged particles in atoms.

• Rutherford's nuclear model solidified the understanding of atomic structure, despite its limitations.

• Bohr's modifications incorporated quantized electron shells and stability while explaining spectral lines.

• Understanding valency reveals the combining capacity of elements, which is crucial for chemical reactions.

• Isotopes and isobars expand the understanding of atomic variety and the properties of elements across different scenarios.