Atomic Structure and Electronic Configuration PDF

## Atomic Structure and Electronic Configuration ### Introduction - All matter, whether in solid, liquid, or gaseous form, is composed of **atoms**, the smallest indivisible particles. - **Atoms** of the same element are identical in weight, size, and properties. - **Atoms** of different elements...

## Atomic Structure and Electronic Configuration ### Introduction - All matter, whether in solid, liquid, or gaseous form, is composed of **atoms**, the smallest indivisible particles. - **Atoms** of the same element are identical in weight, size, and properties. - **Atoms** of different elements differ in weight, size, and other characteristics. - The size of atoms is on the order of 1 Å (10⁻¹⁰ m). - A material consisting of only one type of atom is called an **element**. Examples include nitrogen, carbon, hydrogen, aluminum, copper, gold, and iron. - A group of atoms that tend to exist together in a stable form are called **molecules**. Examples include H₂, O₂, and N₂. - Many molecules in nature are combinations of atoms from different elements, such as water (H₂O). - **Molecules** can be monoatomic, diatomic, triatomic, or polyatomic. ### The Electron - Scientists like Sir J.J. Thomson, Rutherford, Niels Bohr, and others discovered that atoms, though indivisible, are composed of smaller particles called **electrons**, **protons**, and **neutrons**. - **Electrons** and **protons** carry electrical charges, while **neutrons** are neutral. - Rutherford and his co-workers showed that the mass of an atom is concentrated in its center, called the **nucleus**. - The **nucleus** is made up of **protons** and **neutrons**, and it's surrounded by a highly structured configuration of **electrons** revolving around it in shells or orbits. - The size of the **nucleus** is on the order of 10⁻¹⁴ m, while its density is about 2 × 10¹⁷ kg/m³. - The mass of an **electron** is negligible compared to the mass of **protons** and **neutrons**, so the mass of an atom mostly depends on its **protons** and **neutrons**. ### The Electron - Michael Faraday, in his experiments on electrolysis, was the first to provide evidence that electrical charge exists in discrete units. - J.J. Thomson observed cathode rays when studying electrical discharges in low-pressure gases. - W. Crookes investigated the properties of cathode rays and confirmed that they: - travel in a straight line and cast shadows, - carry a negative charge and have momentum, - possess high kinetic energy and can induce chemical reactions, excite fluorescence on certain substances. - Thomson's hypothesis stated that cathode rays are composed of a stream of particles, each with a mass (m) and charge (e) of -1.602 × 10⁻¹⁹ C. He called these particles **electrons**. - Thomson determined the specific charge (e/m) of the **electron** to be -1.76 × 10¹¹ C/kg, signifying its universality. ### Protons - The **nucleus** of a hydrogen atom is called a **proton**. - A **proton** carries a unit positive charge of +1.602 × 10⁻¹⁹ C. - The mass of a **proton** is 1.672 × 10⁻²⁷ kg. - The **proton** and **neutron** are considered to be two different charge states of the same particle, called a **nucleon**. - The number of **protons** in a nucleus is the **charge number (Z)**, which determines the element. ### Neutrons - **Neutrons** are electrically neutral particles, about 1.008 times heavier than **protons**. - The mass of each **neutron** is 1.675 × 10⁻²⁷ kg. - A **neutron** can be considered as composed of one **proton** and one **electron**. - The number of **neutrons** in a nucleus is denoted by **N**. - For most nuclei, N ≥ Z. - For light nuclei, the N/Z ratio is approximately 1. - For heavier nuclei, the N/Z ratio is approximately 1.6. - The **mass number (A)** of a nucleus is the total number of nucleons, A = N + Z. ### Other Subatomic Particles - **Positrons (e⁺)** are the antimatter counterparts of electrons. They are positively charged and annihilate with electrons, producing gamma rays. - **Neutrinos (v)** and **antineutrinos (v)** have small mass and no charge but may be associated with energy changes during radioactive emissions. - **Mesons** are particles with mass between that of an electron and a proton. - **Deutrons (D)** are heavy isotopes of hydrogen with mass twice that of ordinary hydrogen. ### Atomic Models - Atomic models aim to explain the structure of atoms and the arrangement of their electrons. #### Thomson's Atomic Model - Thomson proposed the **plum pudding model** in 1911. - It suggested that atoms are uniform spheres with electrons distributed like plums in a pudding, with a uniform positive charge throughout the sphere. - This model failed to explain the phenomena of alpha particle scattering and the emission of spectral series. #### Rutherford's Atomic Model - Rutherford conducted alpha particle scattering experiments in 1911. - He observed that most alpha particles passed through a thin gold foil without deflection but some were deflected at large angles, some even rebounding back. - His observations contradicted Thomson's model, leading to the **nuclear model**. - The **nuclear model** proposed: - The positive charge and most of the mass of an atom are concentrated in a small core called the **nucleus** at the center. - **Electrons** revolve around the **nucleus** in orbits at relatively larger distances. - The **nucleus** occupies a small volume compared to the overall size of the atom. - The **electrons** carry a negative charge and are balanced by the positive charge of the **nucleus**, making the atom electrically neutral. - The Rutherford model explained the scattering of alpha particles but failed to explain the stability of the atom as it predicted that electrons should continuously lose energy by emitting radiation and eventually collapse into the nucleus. #### Bohr's Atomic Model - Bohr proposed a model for the hydrogen atom in 1913. - **Bohr’s Model** incorporated quantized energy levels based on Planck's quantum theory. - **Bohr's postulates:** 1. **Stationary orbits:** The electron can only exist in specific circular orbits around the nucleus, called stationary orbits, with fixed angular momentum. The electrons do not emit radiation while moving in these orbits. 2. **Energy transitions:** Radiation is only emitted when an electron jumps from a higher energy state to a lower energy state, with the difference in energy being emitted as a photon. 3. **Energy quantization:** The angular momentum of the electron in the nth stationary orbit is equal to nh/2π, where n is a positive integer called the principal quantum number. - This quantization of energy levels results in distinct spectral lines at specific frequencies when electrons transition between orbits. #### Sommerfeld's Atomic Model - Sommerfeld extended Bohr's model by incorporating elliptical orbits. - He introduced the **azimuthal (or orbital) quantum number (l)**, with values 0, 1, 2, ... (n-1). - The value l = 0 corresponds to a circular orbit. - The value l = 1 corresponds to an elliptical orbit with a single loop. - The value l = 2 corresponds to an elliptical orbit with two loops. - The values of l are designated by s, p, d, f, ... respectively for 0, 1, 2, 3, and so on. - Sommerfeld's model also introduced the concept of **relativistic correction**, which further explained the fine structure of spectral lines observed. - Sommerfeld's model gave an atomic model with a greater degree of precision than Bohr's model, but it still faced limitations: - It failed to explain the observed energy levels of multielectron atoms. - It did not address electron spin. - It provided nothing about the spatial orientation of electrons. #### The Vector Atomic Model - The **vector atomic model** is the current widely accepted model of the atom. - It is based on the principles of wave mechanics, which emphasizes the wave nature of electrons and the associated principles of quantum mechanics. - The vector atomic model incorporates the **four quantum numbers:** - **Principal quantum number (n)**: Determines the electron's energy level or shell (1, 2, 3, ...). - **Angular momentum quantum number (l)**: Describes the shape of the electron's orbital (0, 1, 2, ... n- 1) and is designated by s,p, d, f for 0, 1, 2, 3, and so on. - **Magnetic quantum number (ml)**: Specifes the spatial orientation of the orbital in space (-l, -l+1, ... 0, ... l-1, l). - **Spin quantum number (ms)**: Describes the intrinsic angular momentum of the electron, which is also quantized, with possible values of +1/2 or -1/2. - The combination of these four quantum numbers uniquely defines the state of an electron, making it inherently unique within an atom. - The vector model has been instrumental in explaining the Zeeman effect (splitting of spectral lines in a magnetic field) and the Stark effect (splitting of spectral lines in an electric field). ### Pauli Exclusion Principle - Wolfgang Pauli's exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers. Each electron must have a unique combination of n, l, ml, and ms. - This principle has significant implications for atomic structure and electronic configurations, as it dictates how electrons occupy shells and subshells. - Every shell has n subshells. Each subshell can accommodate 2(2l + 1) electrons. - The maximum number of electrons in the nth shell is 2n². ### Electronic Configuration - The electronic configuration refers to the arrangement of electrons in an atom's shells and subshells. - The electronic configuration is determined by the filling of the subshells in order of increasing energy. - This order can be obtained by the Aufbau principle (building up principle) or by the diagonal rule. - The electronic configuration of an element indicates its chemical behavior and its position in the periodic table. ### Wave Mechanical Picture of the Atom - Wave mechanics, developed by de Broglie, considers the wave nature of particles, including electrons. - Every particle in motion is associated with a wave, described by the de Broglie wavelength (λ = h/mv). - The uncertainty principle, formulated by Heisenberg, states that it's impossible to simultaneously determine both the position and momentum of a particle with absolute certainty. - This means that the classical model of definite orbits for electrons becomes inadequate. Instead, wave mechanics utilizes probability distributions to describe the location of electrons within a given region of space. - Erwin Schrödinger's time-independent wave equation (ψ²) provides the probability of finding an electron at a specific point in space and time. - The solutions to Schrödinger's equation lead to the wave functions (ψ), which provide a complete mathematical description of the electron's state in the atom. - The wave mechanical model successfully accounts for the quantization of energy levels and the properties of atoms, particularly their spectral behavior. ### Periodic Table - The periodic table is a systematic organization of elements based on their properties. - Mendeleev’s periodic table arranged elements in order of increasing atomic weight, revealing a repeating pattern of chemical properties. - Moseley discovered that the properties of elements are a periodic function of their atomic number. - The modern periodic table is based on the atomic number and is divided into four blocks: s, p, d, and f blocks, depending on the subshells electrons occupy. - The table consists of seven horizontal rows called periods and 18 vertical columns called groups. - Elements in the same group have similar electronic configurations and similar chemical properties, while elements in the same period have different properties but share the same highest principal quantum number (n). ### Summary of Atomic Structure - Atoms are composed of a central positively charged nucleus surrounded by negatively charged electrons. - The nucleus contains protons and neutrons, which account for the majority of the atom's mass. - Electrons occupy specific energy levels or shells, each with a definite quantum number (n). - Every shell has n subshells, each with a unique angular momentum quantum number (l). - Each subshell contains orbitals with well-defined spatial orientations (ml) and spin directions (ms). - The Pauli exclusion principle states that no two electrons in an atom share the same set of four quantum numbers. - Electronic configurations describe the arrangements of electrons in atoms, which in turn influence their chemical behavior. - The periodic table is a systematic compilation of elements in order of increasing atomic number, highlighting the recurring patterns of their chemical properties. - Wave mechanics provides a more realistic model of the atom by considering the wave nature of electrons and introducing probability distributions for their location. This comprehensive description captures the fundamental principles of atomic structure and electronic configuration, highlighting their importance in explaining the properties and periodic trends observed in the chemical world. The progression of atomic models, culminating in the current vector model based on wave mechanics and quantum numbers, illustrates the evolving understanding of atomic structure, leading to a more sophisticated and nuanced comprehension of the fundamental building blocks of matter.