Atomic Structure and Periodicity Part 2 PDF

Chapter 7 Atomic structure and periodicity Part - 2 1 Chapter 7 –Atomic structure and periodicity 7.1 Electromagnetic Radiation 7.2 The Nature of Matter 7.3 The Atomic Spectrum of Hydrogen 7.4 The Bohr Model 7.5 The Quantum Mechanical Description of the Atom...

Chapter 7 Atomic structure and periodicity Part - 2 1 Chapter 7 –Atomic structure and periodicity 7.1 Electromagnetic Radiation 7.2 The Nature of Matter 7.3 The Atomic Spectrum of Hydrogen 7.4 The Bohr Model 7.5 The Quantum Mechanical Description of the Atom 7.9 The characteristics of Hydrogen Orbitals 7.10 Electron Spin and the Pauli Principle 7.13 The Aufbau Principle and the Periodic Table 7.15 Periodic Trends in Atomic Properties 2  The Quantum Mechanical Description - Since the Bohr model was not a valid model. - Heisenberg, De Broglie, and Schrödinger developed the wave mechanics or quantum mechanics. - De Broglie originated the idea that the e- shows wave properties. - Schrödinger decided to explain the atomic structure by emphasizing on the wave properties. - Using a complicated mathematical technique, he formulated an equation that describes the behavior and energies of submicroscopic particles. - The Schrödinger equation specifies the possible energy states the e- can occupy in a hydrogen atom and identifies the corresponding wave functions (Ψ). - A specific wave function for a given e- is often called an orbital. - However, an orbital is not a Bohr orbit. 3 The Quantum Mechanical Description Orbit 4  The Quantum Mechanical Description - The wave function gives us no information about the movements of the e-. - Heisenberg uncertainty principle: There is a fundamental limitation to just how precisely we can know both the position and the momentum (mass x velocity) of a particle at a given time. - Quantum mechanics cannot pinpoint the electron in an atom it defines a region where the electron might be at a given time. - Electron density is the measure of the probability of an electron being present at a specific location. - Regions of high electron density represent a high probability of locating the electrons. - An atomic orbital is a region around the nucleus where there is a high probability of finding an electron. 5  The characteristics of Hydrogen Orbitals After solving the Schrödinger equation for the hydrogen atom, it was found that many wave functions (orbitals) satisfy it. Each of these orbitals is characterized by a set of quantum numbers. Quantum numbers are required to describe the distribution of electrons in hydrogen and other atoms. 1. The principle quantum number (n) 2. The angular momentum quantum number (l) 3. The magnetic quantum number (ml) 4. The spin quantum number (ms) 6  The principal quantum number (n) It can have integral values (1,2,3,…) It is related to the size and energy of the orbital As n increases - the orbital becomes larger and e- are farther from the nucleus. - higher energy because e- are less tightly bound to the nucleus, and the energy is less negative. 7  The angular momentum quantum number (l) It can have integral values from 0 to n-1 for each value of n. It is related to the shape of the orbital The values of l are designated by letters s, p, d, f, … Each set of orbitals with a given value of l is called subshell. 8  The magnetic quantum number (ml) It can have integral values between l and -l, including zero. It is related to the orientation in space of the angular momentum associated with the orbital and to the number of orbitals. 9  The magnetic quantum number (ml) Example 7.7 For principal quantum level n=5, determine the number of subshells (different values of l) and give the designation of each. 10  The magnetic quantum number (ml) SOLUTION For n=5, l=0 l=1 l=2 l=3 l=4 5s 5p 5d 5f 5g 11  Orbital Shapes and Energies Orbitals are presented by the surface that surrounds 90% of the total electron probability. s orbitals (l=0) are spherical Figure 12.18 - Three representation of the hydrogen 1s, 2s and 3s orbitals 12  Orbital Shapes and Energies p orbitals (l=1) have two lobes separated by a node at the nucleus. They are labeled according to the axis along which the lobes lie. Figure 12.19 - Representation of the 2p orbitals. 13  Orbital Shapes and Energies d orbitals (l=2) first occur in level n =3 Figure 12.21 - Representation of the 3d orbitals. 14  The spin quantum number (ms) and the Pauli Principle It can only have one of two values, +1/2 or -1/2. Pauli principle: In a given atom no two electrons can have the same set of four quantum numbers (n, l, ml, and ms). An orbital can hold only two electrons, and they must have opposite spins. 15  Orbital energy level for the hydrogen atom For the hydrogen atom, the energy of a particular orbital is determined by its value of n. Thus all orbitals with the same value of n have the same energy – they are said to be degenerate. In the lowest energy state, the ground state, the e- resides in the 1s orbital. If energy is put into the atom, the e- can be transferred to a higher-energy orbital, producing an excited state. Figure 12.23 - Orbital energy levels for the hydrogen atom. 16 End of part - 2 17