General Chemistry I Lecture Notes PDF

# General Chemistry I Dr. Mahmoud A. A. Ibrahim Chemistry Department, Faculty of Science, Minia University Email: [email protected] ## Lecture No. 2 ### Quantum numbers & Electron configuration **Quantum Numbers** * Quantum numbers * Pauli exclusion principle **Electron configuration**...

# General Chemistry I Dr. Mahmoud A. A. Ibrahim Chemistry Department, Faculty of Science, Minia University Email: [email protected] ## Lecture No. 2 ### Quantum numbers & Electron configuration **Quantum Numbers** * Quantum numbers * Pauli exclusion principle **Electron configuration** * Aufbau principle * Orbital diagram * Electron configuration and periodic table * Abbreviated electron configuration * Valence shell electron configuration ## **2. Quantum numbers** **Quantum numbers** * Each electron in an atom is described by four different quantum numbers. **Principle quantum number (n):** this is any positive integral number (excluding zero), specifies the energy of an electron and the size of the orbital. **Angular quantum number (l):** Specifies the shape of an orbital with a particular principal quantum number. The secondary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels). The value of *l* also has a slight effect on the energy of the subshell; the energy of the subshell increases with *l* (s < p < d < f). **Magnetic quantum number (m_l):** m_l = -l, ..., 0, ..., +l. Specifies the orientation in space of an orbital of a given energy (n) and shape (l). This number divides the subshell into individual orbitals which hold the electrons; there are 2l+1 orbitals in each subshell. Thus the s subshell has only one orbital, the p subshell has three orbitals, and so on. **Spin quantum number (ms):** ms = +1/2 or -1/2. Specifies the orientation of the spin axis of an electron. An electron can spin in only one of two directions (sometimes called up and down). **Pauli exclusion principle** >The Pauli exclusion principle states that no two electrons in the same atom can have identical values for all four of their quantum numbers. What this means is that no more than two electrons can occupy the same orbital, and that two electrons in the same orbital have opposite spins. | n | l | m_l | Number of Orbital | Number of electrons | Name | |---|---|---|---|---|---| | 1 | 0 | 0 | 1 | 2 | 1s | | 2 | 0 | 0 | 1 | 2 | 2s | | 2 | 1 | -1, 0, +1 | 3 | 6 | 2p | | 3 | 0 | 0 | 1 | 2 | 3s | | 3 | 1 | -1, 0, +1 | 3 | 6 | 3p | | 3 | 2 | -2, -1, 0, +1, +2 | 5 | 10 | 3d | | 4 | 0 | 0 | 1 | 2 | 4s | | 4 | 1 | -1, 0, +1 | 3 | 6 | 4p | | 4 | 2 | -2, -1, 0, +1,+2 | 5 | 10 | 4d | | 4 | 3 | -3,-2,-1,0,+1,+2,+3 | 7 | 14 | 4f | ## **3. Electron configuration** **The distribution of electrons among the orbitals of an atom is called the electron configuration** **Aufbau principle** >The electrons are filled in according to a scheme known as the Aufbau principle ("building-up"), which corresponds (for the most part) to increasing energy of the subshells: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, >5f write in the orbitals that are occupied by electrons, followed by a superscript to indicate how many electrons are in the set of orbitals (e.g., H 1s¹). **Orbital diagram** >Another way to indicate the placement of electrons is an orbital diagram, in which each orbital is represented by a square (or circle), and the electrons as arrows pointing up or down (indicating the electron spin). When electrons are placed in a set of orbitals of equal energy, they are spread out as much as possible to give as few paired electrons as possible. **Electron configuration and periodic table** >The periodic table is a useful tool for predicting electron configurations of elements. **Abbreviated electron configuration** Electron configurations are often abbreviated by naming the last element with a filled shell (e.g. He and Ne) in brackets and listing only the orbitals after the filled shell. **Valence shell electron configuration** >In this configuration type, only electrons in the outer shell (valence shell) are mentioned out. >Example, write the valence shell electron configuration for Na and Fe atoms: Na: 3S1 Fe: 3d6 4S2 **Unexpected electron configuration** >Some unexpected configurations occur (Chromium, Cr, & Copper, Cu for example) because a special stability occurs when sets of d-orbitals are exactly half filled OR completely filled (more stable) AND s-orbitals are half filled OR completely EMPTY - as a result of the d-orbitals becoming half filled or completely filled with electrons. >Example, Cr: [Ar] 3d5 4s¹ >Example, Cu: [Ar] 3d10 4S1 ## End of lecture 2