# Brief History of Atomic Models This section explores the evolution of the atomic model, starting from the limitations of early theories and progressing to the discovery of sub-atomic particles and the principles governing atomic structure. --- ## Evolution of Atomic Theory Based on discharge tube experiments in the late 19th century, the Daltonian concept of an indivisible atom was challenged. **Niels Bohr's atomic model** was a significant step forward, introducing the idea of quantized energy levels or electronic shells with fixed energy. Bohr's model successfully calculated the radius of the hydrogen atom and the energy of its electron. It also explained the emission spectrum of hydrogen. However, with the advent of high-resolution spectrophotometers, several limitations became apparent: * **Fine Spectrum:** Spectral lines were found to be split into finer lines, which Bohr's model could not explain. * **Stark Effect:** The splitting of spectral lines in an electric field. * **Zeeman Effect:** The splitting of spectral lines in a magnetic field. These observations, along with **Heisenberg's Uncertainty Principle**, led to a more complex and refined atomic structure involving concepts like subshells and orbitals. For details on these quantum mechanical concepts, refer to . --- ## 2.1.1 Sub-Atomic Particles Atoms are composed of smaller, fundamental particles. The three primary sub-atomic particles are electrons, protons, and neutrons. * **Electron:** * A negatively charged particle discovered by J.J. Thomson. * Deflected towards the positive pole in an electric field. * **Charge:** $-1.6022 \times 10^{-19} \, \text{C}$ * **Relative Charge:** -1 * **Mass:** $9.1095 \times 10^{-31} \, \text{kg}$ * **Proton:** * A positively charged particle. * Deflected towards the negative pole in an electric field. * It is 1836 times heavier than an electron. * **Charge:** $+1.6022 \times 10^{-19} \, \text{C}$ * **Relative Charge:** +1 * **Mass:** $1.6727 \times 10^{-27} \, \text{kg}$ * **Neutron:** * An electrically neutral particle, meaning it is undeflected in electric and magnetic fields. * Slightly more massive than a proton. * **Charge:** 0 C * **Relative Charge:** 0 * **Mass:** $1.6750 \times 10^{-27} \, \text{kg}$ **Table of Sub-Atomic Particle Properties** | Particle | Mass (kg) | Charge (C) | Relative Mass (amu) | Relative Charge | | :--- | :--- | :--- | :--- | :--- | | **Proton** | $1.6727 \times 10^{-27}$ | $+1.6022 \times 10^{-19}$ | ~1 | +1 | | **Neutron** | $1.6750 \times 10^{-27}$ | 0 | ~1 | 0 | | **Electron** | $9.1095 \times 10^{-31}$ | $-1.6022 \times 10^{-19}$ | 1/1836 | -1 | --- ## 2.1.2 Behaviour of Sub-Atomic Particles in an Electric Field When passed through a uniform electric field, the paths of electrons, protons, and neutrons diverge based on their charge. * **Electrons** (negative charge) are deflected towards the positive plate. * **Protons** (positive charge) are deflected towards the negative plate. * **Neutrons** (no charge) pass through the field undeflected. The degree of curvature depends on the particle's mass, charge, and velocity. --- ## 2.1.3 Atomic Number and Mass Number * **Atomic Number (Z):** The number of protons in the nucleus of an atom. In a neutral atom, the number of protons equals the number of electrons. * *Number of Protons = Number of electrons (in a neutral atom)* * **Mass Number (A):** The total number of protons and neutrons in the nucleus. It is also called the nucleon number and is approximately equal to the atomic mass. $$ \text{Mass Number (A)} = \text{Number of Protons (Z)} + \text{Number of Neutrons (N)} $$ $$ A = Z + N $$ * **Calculating Electrons in Ions:** * **Cation (Positive Ion):** Number of electrons = Atomic number (Z) - magnitude of charge. * **Anion (Negative Ion):** Number of electrons = Atomic number (Z) + magnitude of charge. --- ## 2.1.4 Atomic and Ionic Radius * **Atomic Radius:** The average distance from the nucleus to the outermost electron shell. Since the electron cloud boundary is not sharp, this is an average measure of an atom's size. ### Periodic Trends in Atomic Radius 1. **Across a Period (Left to Right):** * The atomic radius *decreases*. * **Reason:** The number of protons (nuclear charge) increases, pulling the electrons closer to the nucleus. The shielding effect from inner electrons remains relatively constant. * *Note:* In transition metals, the trend is less regular as electrons are added to inner d-orbitals, which increases shielding. 2. **Down a Group (Top to Bottom):** * The atomic radius *increases*. * **Reason:** A new electron shell is added for each period, placing the outermost electrons further from the nucleus. The increased shielding from inner electrons also weakens the nucleus's pull. ### Ionic Radius * **Cations (Positive Ions):** Cations are *smaller* than their parent atoms. * **Reason:** Removing electrons reduces electron-electron repulsion and increases the effective nuclear charge, pulling the remaining electrons closer. * **Anions (Negative Ions):** Anions are *larger* than their parent atoms. * **Reason:** Adding electrons increases electron-electron repulsion, causing the electron cloud to expand. **Trends in Ionic Radius:** * **Down a Group:** Ionic radius increases, similar to atomic radius, due to the addition of electron shells. * **Across a Period:** The trend is complex because elements form ions with different charges. However, for a series of *isoelectronic ions* (ions with the same number of electrons), the radius decreases as the atomic number (number of protons) increases. The trends in atomic and ionic radii are crucial as they influence other properties like ionization energy, electronegativity, bond length, and chemical reactivity. --- ## Worked Examples *From Concept Assessment Exercise 2.1* **1. How many electrons are present in the following species: $Mg^{+2}, Al^{+3}, S^{-2}$?** * **For $Mg^{+2}$:** 1. Atomic number of Mg (Z) = 12. A neutral Mg atom has 12 protons and 12 electrons. 2. The +2 charge means it has lost 2 electrons. 3. Number of electrons = $12 - 2 = 10$ electrons. * **For $Al^{+3}$:** 1. Atomic number of Al (Z) = 13. A neutral Al atom has 13 protons and 13 electrons. 2. The +3 charge means it has lost 3 electrons. 3. Number of electrons = $13 - 3 = 10$ electrons. * **For $S^{-2}$:** 1. Atomic number of S (Z) = 16. A neutral S atom has 16 protons and 16 electrons. 2. The -2 charge means it has gained 2 electrons. 3. Number of electrons = $16 + 2 = 18$ electrons. **2. Find the number of electrons, protons, and neutrons for the following species: $Mg^{+2}, Cl^{-1}, Al^{+3}, S^{-2}$.** *(Assume the most common isotopes: $^{24}Mg, ^{35}Cl, ^{27}Al, ^{32}S$)* | Species | Protons (Z) | Electrons | Neutrons (A-Z) | | :--- | :--- | :--- | :--- | | **$Mg^{+2}$** | 12 | $12 - 2 = 10$ | $24 - 12 = 12$ | | **$Cl^{-1}$** | 17 | $17 + 1 = 18$ | $35 - 17 = 18$ | | **$Al^{+3}$** | 13 | $13 - 3 = 10$ | $27 - 13 = 14$ | | **$S^{-2}$** | 16 | $16 + 2 = 18$ | $32 - 16 = 16$ | --- ## Possible Questions/Answers *From Concept Assessment Exercise 2.2* 1. **Q:** As you move from left to right across a period in the periodic table, the atomic radius generally \_\_\_\_\_. **A:** *decreases*. 2. **Q:** Down a group in the periodic table, the atomic radius tends to \_\_\_\_\_. **A:** *increase*. 3. **Q:** When an atom loses electrons to become a cation, its ionic radius \_\_\_\_\_ compared to its atomic radius. **A:** *decreases*. 4. **Q:** Among elements, as you move from top to bottom within a group, the ionic radius of cations tends to \_\_\_\_\_, while the ionic radius of anions tends to \_\_\_\_\_. **A:** *increase*, *increase*. 5. **Q:** In general, nonmetals tend to form \_\_\_\_\_ ions with \_\_\_\_\_ ionic radii compared to their atomic radii. **A:** *negative (anions)*, *larger*. --- ## References (Derived from FBISE textbook)