Electronic Structure of Atoms and the Periodic table. Electron Spin Quantum # m s. The Pauli Exclusion Principle

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1 Electronic Structure of Atoms and the Periodic table Chapter 6 & 7, Part 3 October 26 th, 2004 Homework session Wednesday 3:00 5:00 Electron Spin Quantum # m s Each electron is assigned a spinning motion along an imaginary axis. The electrons present in each orbital must spin in opposite directions; clockwise (m s = +1/2) and anticlockwise (m s = -1/2). A pair of electrons in an orbital with opposite spins are called paired. 42 The Pauli Exclusion Principle No two electrons in an atom can have the same 4 quantum numbers. If an electron has n =1, l = 0, m l = 0 and m s = +1/2, no other electron can have the same quantum numbers. 43 1

2 Electron Configuration Specific arrangement of electrons among the available orbitals. Some of the rules that we have to follow: Determine the number of electrons in the neutral atom from the atomic number. Start by placing electrons in the first available lowest energy orbital. Each orbital can contain ONLY 2 electrons. Each principal energy level can have a MAXIMUM of 2(n) 2 electrons. 44 Orbital Filling Order 4f 3d 4d 5d 2p 3p 4p 5p 6p 1s 2s 3s 4s 5s 6s 7s ENERGY 45 Orbitals and their energies 46 2

3 The Aufbau Principle The Aufbau (building up) principle is used to describe the electron configuration of the elements. Electron configurations are written by writing the orbitals in an increasing order of energy. The number of electrons in each orbital is written as a superscript. 47 Element # of electrons Electron Configuration Hydrogen (H) 1 1s 1 Helium (He) 2 1s 2 Lithium (Li) 3 1s 2 2s 1 Beryllium (Be) 4 1s 2 2s 2 Boron (B) 5 1s 2 2s 2 2p 1 Carbon (C) 6 1s 2 2s 2 2p 2 Nitrogen (N) 7 1s 2 2s 2 2p 3 Oxygen (O) 8 1s 2 2s 2 2p 4 Fluorine (F) 9 1s 2 2s 2 2p 5 Neon (Ne) 10 1s 2 2s 2 2p 6 48 Element Sodium Na Magnesium Mg Aluminum Al Silicon Si Phosphorus P Sulfur S Chlorine Cl Argon Ar Potassium K Calcium Ca # of electrons Electron Configuration 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 1 1s 2 2s 2 2p 6 3s 2 3p 2 1s 2 2s 2 2p 6 3s 2 3p 3 1s 2 2s 2 2p 6 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s

4 Orbital diagrams Also known as a box diagram. Each orbital is represented by a box and each orbital can only contain 2 electrons with opposite spins. 50 Hund s Rule The lowest energy arrangement of electrons in a subshell is obtained by putting electrons into orbitals in the same subshell with the same spin before pairing electrons. Consider the C atom which has the electron configuration 1s 2 2s 2 2p Importance of Electron Configuration Compounds are formed by the combination of atoms of different elements. These atoms come together by the formation of bonds. Valence electrons are responsible for the bond formation. Valence electrons are present in the outermost principal energy level (shell) of an atom. Examining the electron configuration allows us to determine the # of valence electrons. 52 4

5 Valence Electrons Consider the Sodium atom with the electron configuration Na: 1s 2 2s 2 2p 6 3s 1, the total # of electrons = 11. The # of valence electrons in Na = 1. Neon has the electron configuration Ne: 1s 2 2s 2 2p 6, the # of valence electrons = 8. Calcium has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 and has 2 valence electrons. The periodic table is a very useful tool for determining the # valence electrons. 53 History of The Periodic Table Scientists had been observing chemical properties of different elements and had tried to put the elements in groups so that the elements with similar properties were together. During , Johann Dobereiner published results that grouped elements in triads. All the elements in a triad had similar chemical properties. During , John Newlands proposed his law of octaves. According to this law if elements were arranged in the order of increasing atomic masses, then the 1 st element is similar to the 8 th, the 2 nd similar to the 9 th and so on. 54 Modern Periodic Table Dmitri Mendeleev and Lothar Meyer independently developed and published the periodic table as we know it now. Mendeleev s Periodic Law states that elements arranged in the order of increasing atomic masses show a regular variation in their properties. The current form of the periodic table is based on the ATOMIC NUMBER and not the ATOMIC MASS, where the elements are arranged in the order of increasing atomic numbers. 55 5

6 Valence electrons and the Periodic Table The group number of all A group elements = # of valence electrons in that element. Na is in Group 1A and the # of valence electrons in Na = 1. For Ca (Group 2A) = 2, Al = 3 etc. The period number = n ; for Na (third period), the valence electrons are in the shell with n = Valence Electrons & the Octet Rule Atoms of different elements will react in a manner so that they achieve the electron configuration of the noble gas nearest to them. Noble gases have electron configurations in which the valence shell is completely filled. Atoms can lose or accept electrons to form ions with positive (lose e - ) or negative (accept e - ) charges. The positively charged ions are called cations while the negatively charged ions are called anions. 57 Abbreviated Electron Configuration Uses the fact that noble gases have completely filled shells. Puts the valence electrons on display. The noble gas that is previous to the element of interest is written first followed by the valence electrons. Na: 1s 2 2s 2 2p 6 3s 1, the noble gas before Sodium is Neon, which has the electron configuration 1s 2 2s 2 2p 6. The electron configuration of Sodium can now be written as [Ne] 3s 1, where the [Ne] represents all the core electrons. 58 6

7 Element # of electrons Electron Configuration H 1 1s 1 He 2 1s 2 Li 3 [He] 2s 1 Be 4 [He] 2s 2 B 5 [He] 2s 2 2p 1 C 6 [He] 2s 2 2p 2 N 7 [He] 2s 2 2p 3 O 8 [He] 2s 2 2p 4 F 9 [He] 2s 2 2p 5 Ne 10 [He] 2s 2 2p 6 59 Effective Nuclear Charge Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nuclear charge. The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons. As the average number of screening electrons (S) increases, the effective nuclear charge (Z eff ) decreases. As the distance from the nucleus increases, S increases and Z eff decreases. 60 Sizes of Atoms and Ions Consider a simple diatomic molecule. The distance between the two nuclei is called the bond distance. If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom. 61 7

8 Sizes of Atoms and Ions As the principal quantum number increases, the size of the orbital increases. Consider the s orbitals. All s orbitals are spherical and increase in size as n increases. The spherical symmetry of the orbitals can be seen in the contour plots. Contour plots are connecting points of equal electron density. 62 8