T1chem

Particles in the Atom and Atomic Radius

Protons, Neutrons and Electrons

• The nucleus of an atom concentrates the mass of the atom.

• Subatomic Particle Characteristics:

  • Protons:

  • Relative mass: 1

  • Relative charge: 1+

  • Neutrons:

  • Relative mass: 1

  • Relative charge: 0

  • Electrons:

  • Relative mass: 1/1836

  • Relative charge: 1-

Charge Interaction between Particles

• When subatomic particles are passed between two oppositely charged plates:

  • Protons are deflected towards the negative plate (due to positive charge).

  • Electrons are deflected towards the positive plate (due to negative charge).

  • Neutrons continue on a straight path (due to no charge).

• If particles have the same energy when passing between charged plates, the amount of deflection of protons and electrons is equal but in opposite directions.

• When particles have the same speed, lighter electrons will be deflected more strongly than protons.

Atomic and Mass Numbers

• Atomic Number (Z): The number of protons in an atom (also called proton number).

  • Atoms of the same element share the same atomic number and number of protons.

  • Atoms are neutral, thus number of electrons equals the number of protons.

• Ions: Charged species where the number of electrons differs from protons. For ions:

  • Positive charge: Number of electrons = Atomic Number - Number of electrons lost.

  • Negative charge: Number of electrons = Atomic Number + Number of electrons gained.

• Mass Number (A): Total number of protons and neutrons in an atom (also called nucleon number).

Atomic Radius

• Across a Period: Atomic radius decreases due to an increase in protons, increasing nuclear charge, which attracts outer electrons closer to the nucleus, thus reducing atomic radius.

• Down a Group: Atomic radius increases due to increased energy levels occupied by outer electrons, making them further from the nucleus. Moreover, higher energy levels are more shielded by inner shell electrons.

Ionic Radius

• Across a Period:

  • For positive ions: Ionic radius decreases because the nuclear attraction increases with increasing proton number; the electron configuration remains unchanged.

  • For negative ions: Ionic radius increases because added electrons result in more electrons than protons, leading to weaker nuclear attraction.

The Nucleus of the Atom

• The nucleus comprises protons and neutrons, termed nucleons.

Isotopes

• Definition: Isotopes are variants of an element that have the same number of protons and electrons but different numbers of neutrons.

• Isotopes share chemical properties (due to identical electron configurations) but exhibit different physical properties (due to differing mass numbers).

Representation of Isotopes

• Isotope notation illustrates the mass number and atomic number related to the element, for example, $^{A}_{Z} ext{X}$

Electrons

Shells

• Electrons occupy distinct shells.

• The innermost shell has the least energy, with energy increasing further from the nucleus.

Orbitals

• Definition: An orbital is a defined area of space where up to 2 electrons can be found.

  • Principal Quantum Number (n): Represents the shell of the electrons; a higher value indicates both higher energy and greater distance from the nucleus.

• Types of Orbitals:

  • s Orbitals: Spherical shape, one s orbital per shell starting from n=1 (total of 2 s electrons per shell, lowest energy).

  • p Orbitals: Dumb-bell shaped, three p orbitals from n=2 onward (total of 6 p electrons per shell, higher energy than s).

  • d Orbitals: Five d orbitals from n=3 onward (total of 10 d electrons per shell, higher energy than p).

• Electrons fill orbitals in order of increasing energy; for example, the filling sequence is: 1s, 2s, 2p, 3s, 4s, 3p, 3d, etc.

• The 4s orbital is filled before the 3d due to lower energy. Electrons must fill each subshell with unpaired electrons before pairing in that subshell occurs.

Electron Configuration Exceptions

• Two exceptions in standard electron configuration include:

  • For Chromium: Configuration is 1s²2s²2p⁶3s²3p⁶3d⁵4s¹ due to stability preference of either half-full or full d sublevel over a partially filled d sublevel.

  • For Copper: Configuration is 1s²2s²2p⁶3s²3p⁶3d¹⁰4s¹.

• For both cases, upon ion formation, electrons are removed from the 4s orbital first.

Examples of Electron Configurations

• Full electron configuration of Iron (Fe): 1s²2s²2p⁶3s²3p⁶3d⁶4s²

• Shorthand electron configuration of Iron: [Ar] 3d⁶4s²

• Box notation for electrons in Iron: [Ar] ↑↓ ↑ ↑ ↑ ↑ ↑↓

Electron Configuration

Definition

• Electron configurations indicate the distribution of electrons among the orbitals of an atom.

Examples and Explanation

• B: 1s²2s²2p¹ (2 energy levels, 3 outer shell electrons)

• Ne: 1s²2s²2p⁶ (2 energy levels, 8 outer shell electrons)

• Cl: 1s²2s²2p⁶3s²3p⁵ (3 energy levels, 7 outer shell electrons)

• Cl⁻: 1s²2s²2p⁶3s²3p⁶ (1 electron gained)

• Na: 1s²2s²2p⁶3s¹ (3 energy levels, 1 outer shell electron)

• Na⁺: 1s²2s²2p⁶ (1 electron lost)

Ground State

• Ground state indicates that all electrons are in the lowest energy orbitals available, making the configuration stable.

First Ionisation Energy

Definition

• First ionisation energy is defined as the energy necessary to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions, expressed in kJ mol⁻¹.

Example

• $Na(g) <br>ightarrow Na^+(g) + e^-$

• Note: The state symbol (g) indicates that the process occurs in the gas phase.

Factors Affecting Ionisation Energy

• Nuclear Charge: Greater number of protons lead to increased nuclear charge and stronger attraction for outer electrons, increasing first ionisation energy.

• Atomic Radius: A larger atomic radius results in weaker attractive forces between nucleus and outer electrons, decreasing first ionisation energy.

• Electron Shielding: More electron shells increase the electron shielding effect, leading to weaker attraction and lower first ionisation energy.

Ionisation Energy Trends

• Across a Period: First ionisation energy increases due to raising nuclear charge, which reduces atomic radius and increases attraction for outer shell electrons.

• Exceptions:

  • Between Groups 2 and 3: An electron from a 2p orbital is more easily removed due to higher energy than from a 2s orbital.

  • Between Groups 5 and 6: Electron pairs in 2p orbitals repel each other, making removal less energetically demanding.

• Down a Group: First ionisation energy decreases, since both atomic radius and electron shielding increase despite the increase in nuclear charge reducing attraction.

Successive Ionisation Energies

Definition

• Successive ionisation energies refer to the energy needed to remove successive electrons from gaseous ions.

Example of Aluminium

• First Ionization: $Al(g) <br>ightarrow Al^+(g) + e^-$ (1st ionization energy = 577 kJ mol⁻¹)

• Second Ionization: $Al^+(g) <br>ightarrow Al^{2+}(g) + e^-$ (2nd ionization energy = 1820 kJ mol⁻¹)

• Third Ionization: $Al^{2+}(g) <br>ightarrow Al^{3+}(g) + e^-$ (3rd ionization energy = 2740 kJ mol⁻¹)

Trends in Successive Ionisation Energies

• Successive ionisation energies generally increase due to decreasing atomic radius and increased attraction to the nucleus.

• A notable jump in ionisation energy indicates a change in the number of outer electrons, revealing the group of the element (for instance, demonstrating three relatively easy-to-remove electrons compared to a significantly harder fourth).

Example Data for Element

• 1st Ionisation Energy: 801 kJ mol⁻¹

• 2nd Ionisation Energy: 2427 kJ mol⁻¹

• 3rd Ionisation Energy: 3660 kJ mol⁻¹

• 4th Ionisation Energy: 25026 kJ mol⁻¹

• 5th Ionisation Energy: 32827 kJ mol⁻¹

Conclusion

• Identification of element family can be made via ionisation trends, in this case suggesting that the element with 5 outer electrons and 3 removed easily would belong to Group 3, identifying it as boron.