Atomic and Electronic Structure 2 PDF

MED-102 General Chemistry Atomic and Electronic Structure 2  Learning Objectives (LOBs) Discuss the rules that govern the values allowed for the various quantum numbers Determine electronic configurations for neutral atoms as well as ions. Apply Hund’s Rule in writing electronic config...

MED-102 General Chemistry Atomic and Electronic Structure 2  Learning Objectives (LOBs) Discuss the rules that govern the values allowed for the various quantum numbers Determine electronic configurations for neutral atoms as well as ions. Apply Hund’s Rule in writing electronic configurations. Define the Pauli Exclusion Principle in determining a set of quantum numbers for an electron in an atom. Explain atomic periodic properties on the basis of electronic configurations.  Quantum Numbers Principal Quantum number, n Determines the size of the orbital. Also called the Shell quantum number Azimuthal quantum number, l Determines the shape of the orbital. Also called the Subshell quantum number Magnetic quantum number, ml Determines the 3D orientation of the orbital  Quantum Numbers - Rules What values can the quantum numbers take? There are certain rules that must be followed Principal quantum number, n (shell) n = 1,2,3,4,…, Azimuthal quantum number, l (subshell) l = 0,1,2,3,…,n-1 Magnetic quantum number, ml (orbital) ml = -l,-l+1,…,0,1,2,3,…,+l Examples of Quantum Numbers More Examples of Quantum Numbers  More Examples of Quantum Numbers - Revision Slide WATCH: https://www.youtube.com/watch?v=Aoi4j8es4gQ Let us work out the orbitals for n = 3 (third shell). For n = 3, the possible values of l are = 0, 1, 2 (n-1). So, we have three subshells in the third shell. These subshells are labelled 3s, 3p, and 3d. Now, for each subshell, we need to find out how many orbitals there are. The quantum number ml goes from –l to +l in steps of 1. For 3s, l = 0. Thus ml = 0 only. There is one orbital in the 3s subshell. For 3p, l = 1. Thus ml = -1, 0, +1. There are three orbitals in the 3p subshell. For 3d, l = 2. Thus ml = -2, -1, 0, +1, +2. This gives five orbitals in the 3d subshell. There are n2 orbitals in a shell with principal quantum number n. Each orbital can accommodate two electrons. Thus, there are 2n2 electrons in each shell with quantum number n.  Electron-spin Quantum Number We have seen that we need 3 quantum numbers to describe one orbital. Each orbital can hold a total of 2 electrons. In order to differentiate between the two electrons in a given orbital we need another quantum number  Electron-spin – Revision Slide Electrons have a negative charge, and when they ‘spin’ in an atom they create a small magnetic field. Spin-up electrons align in the same direction as the magnetic field. Spin-down electrons align opposite to the magnetic field, and this means that they will have slightly higher energy than the spin-up electrons. An orbital contains a maximum of two electrons. One must be spin-up and the other spin-down. This is necessary so that we can distinguish between the two electrons in the atom. If placing one electron in an empty orbital, it is preferred to have it spin-up, which will give the atom its lowest possible energy. Exercise on Quantum Numbers Consider the orbital box diagram for the element Oxygen: What is a possible set of the four quantum numbers for the electron labelled with the red asterisk (*)? Find n, l, ml, and ms.  Exercise on Quantum Numbers - Answer Consider the orbital box diagram for the element Oxygen: What is a possible set of the four quantum numbers for the electron labelled with the red asterisk (*)? Find n, l, ml, and ms. The electron is in the 2p subshell. The ‘2’ indicates n, n = 2. The ‘p’ corresponds to l = 1. We don’t know the ml number because the orbitals in the 2p subshell are not labelled. It could be -1, 0, or +1. Any of these three values is correct. The electron is spin-up, therefore ms = +1/2.  Summary These are subshells, each containing specific numbers of orbitals Order of orbitals (Aufbau Principle)  Aufbau Principle – Revision Slide WATCH: https://www.youtube.com/watch?v=SvlDqyV5lXQ Aufbau is German for ‘building up.’ We will build/construct electronic configurations. When we place the electrons in individual orbitals, we begin with the lowest energy. The first electron (H atom) will go spin-up in the 1s orbital. The second electron (He atom) will go spin-down in the 1s orbital. Then we start filling in the 2s orbital in the same way, reaching higher and higher in energy. The details and rules of how we place the electrons in orbitals will be discussed in more detail in the next session. Filling in the orbitals Aufbau Principle  Hund’s Rule - Carbon Maximize unpaired spin-up electrons  Hund’s Rule – Carbon – Revision Slide Hund’s Rule states that we should maximize the number of unpaired spin-up electrons. When we pair up electrons in the same orbital, we are forcing them to share the same space. Since electrons are negative, they repel each other. Forcing them to pair up, increases the repulsions and increases the energy of the atom, making it less stable. Thus, we keep electrons unpaired, by placing them in adjacent orbitals. We place the unpaired electrons spin-up, because as we stated earlier on, spin-up electrons have lower energy and lend more stability to the atom. The most stable form of Carbon is then the fourth one, the last one above. WATCH: https://www.youtube.com/watch?v=IMbGqcb8aN4 Electronic states  Pauli Exclusion Principle No two electrons in a given atom can have the same set of 4 quantum numbers – Nitrogen atom example below: 5-Minute Break Full Electronic configurations Noble Gas Electron Configurations Condensed Electronic Configurations Different Types of Electronic Configurations WATCH: https://www.youtube.com/watch?v=uZ_dCxKKv58 Anomalous Electronic configurations One electron is promoted from the 4s to the 3d subshell to achieve a d5 or d10 configuration d5 and d10 have special degree of stability WATCH: https://www.youtube.com/watch?v=7R_Z17dOePI Transition Metal Ions 4s electrons leave first!  Alkali metals – Group 1A Electronic configurations – common s1 outermost electron configuration Outermost shell electrons – Valence Electrons  Halogens – Group 7A Electronic configurations – common s2p5 Atomic Radius  Atomic Radius Atomic radius increases from top to bottom This is entirely due to the addition of extra shells as we move down the Group The extra shells have increasing principal quantum number, n, and therefore become larger in size  Atomic Radius Atomic radius decreases from left to right Electrons are being placed in the same shell (same distance from nucleus), but nuclear charge increases by one unit each time we move to the right – the increasing nuclear charge pulls the shells closer to the nucleus, decreasing the size of the atom  Ionization Energy Energy required to remove one electron from a neutral atom e.g. Na(g) → Na+(g) + e- Recall that electrons are attracted to the nucleus – Coulomb’s Law  Ionization Energy – Revision Slide WATCH: https://www.youtube.com/watch?v=hePb00CqvP0 Q1 is the effective nuclear charge, or the number of protons minus the intermediate electrons. Intermediate electrons shield the nucleus and the valence electron does not feel the full positive charge. Q2 is equal to -1 (one electron). As the atomic number increases, so does Q1 and so does the energy of attraction to the electron. As the distance R of the electron from the nucleus increases, the energy of attraction decreases. As the number of shells increases, R increases too. This serves to decrease the attractive energy. Remember: the ionization energy is the energy we have to expend to remove the electron from the attractive force of the nucleus. Trends in Ionization Energy Ionization Energy – Noble Gases Alkali Metals - Trends  Summary for Revision - 1 There are rules that govern the values taken by the three quantum numbers. Following these rules gives rise to shells, subshells in each shell, and orbitals in each subshell. Each orbital can contain two electrons. The two electrons must have different spin. This gives rise to the electron spin quantum number. s subshells have one orbital, p subshells have three orbitals, d subshells have five orbitals, etc. We use the Aufbau Principle to place electrons in orbitals, in terms of increasing energy. This gives rise to electronic configurations of different atoms. The Aufbau Principle is used to construct electronic configurations. We place electrons in orbitals from low energy to higher energy, first spin-up electrons. Two electrons in the same orbital must have opposite spin. Hund’s Rule states that if electrons are being placed in orbitals having the same energy (in a given subshell), they must be placed in separate orbitals and spin-up. Violations of Hund’s Rule lead to excited electronic states, having higher energy than the lowest ground state. The Pauli Exclusion Principle prevents two electrons in a given atom from having the same set of four quantum numbers. There are three types of electronic configurations: full electronic configurations, condensed electronic configurations, and box orbital diagrams.  Summary for Revision - 2 Some metallic elements have unexpected anomalous electronic configurations. They arise from the need to maintain d5 or d10 electronic arrangements, as they are more stable. They usually come about from the promotion of an ns electron to a (n-1)d orbital. Transition metal cations form by loss of ns electrons first, not (n-1)d electrons. Valence electrons are the electrons in the outermost shell. They determine chemical reactivity. Atomic radius (size of the atom) increases down a group and decreases from left to right in a period. Ionization energy has exactly the opposite trend than atomic radius.