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A periodic table of partial ground-state electron configurations

A periodic table of partial ground-state electron configurations. ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY. Arrangement of Electrons in Atoms. Each orbital can be assigned no more than 2 electrons!.

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A periodic table of partial ground-state electron configurations

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  1. A periodic table of partial ground-state electron configurations

  2. ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY Arrangement of Electrons in Atoms Each orbital can be assigned no more than 2 electrons! This is tied to the existence of a 4th quantum number, the electron spin quantum number, ms. Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (ml)

  3. Electron Spin Quantum Number, ms Can be proved experimentally that electron has a spin. Two spin directions are given by ms where ms = +1/2 and -1/2. Diamagnetic: NOT attracted to a magnetic field. Paramagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.

  4. Categories of Electrons Inner (core) electrons: fill all the lower energy levels of an atom Outer electrons: those electrons in the highest energy level (highest n value) of an atom Valence electrons: those involved in forming compounds; the bonding electrons; among the main-group elements, the valence electrons are the outer electrons

  5. General Observations about the Periodic Table A. The group number equals the number of outer electrons (those with the highest value of n) (main-group elements only) B. The period number is the n value of the highest energy level. C. The n value squared (n2) gives the total number of orbitals in that energy level; 2n2 gives the maximum number of electrons in the energy level. KEY PRINCIPLE All physical and chemical properties of the elements are based on the electronic configurations of their atoms.

  6. QUANTUM NUMBERS n ---> shell 1, 2, 3, 4, ... l ---> subshell 0, 1, 2, ... n - 1 ml ---> orbital -l ... 0 ... +l ms ---> electron spin +1/2 and -1/2 When n = 1, then l = 0 this shell has a single orbital (1s) to which 2e- can be assigned. When n = 2, then l = 0, 1 2s orbital 2e- three 2p orbitals 6e- TOTAL = 8e-

  7. principal n positive integers (1, 2, 3,…) orbital energy (size) angular momentum l integers from 0 to n-1 orbital shape (l values of 0, 1, 2 and 3 correspond to s, p, d and f orbitals, respectively.) magnetic ml integers from -l to 0 to +l orbital orientation spin ms +1/2 or -1/2 direction of e- spin Table 8.2 Summary of Quantum Numbers of Electrons in Atoms Name Symbol Allowed Values Property Each electron in an atom has its own unique set of four (4) quantum numbers.

  8. PROBLEM: Write a set of quantum numbers for the third electron and a set for the eighth electron of the fluorine (F) atom. 9F 1s 2s 2p n = n = l = l = ml = ml = ms= ms= Determining Quantum Numbers from Orbital Diagrams Sample Problem 1 PLAN: Use the orbital diagram to find the third and eighth electrons. Up arrow = +1/2 Down arrow = -1/2 SOLUTION: The third electron is in the 2s orbital. Its quantum numbers are: 2 0 0 +1/2 The eighth electron is in a 2p orbital. Its quantum numbers are: 2 1 -1 -1/2

  9. Pauli Exclusion Principle No two electrons in the same atom can have the same set of 4 quantum numbers. That is, each electron in an atom has a unique address of quantum numbers. Hund’s Rule When orbitals of equal energy are available, the electron configuration of lowest energy has the maximum number of unpaired electrons with parallel spins. For a given n value, the lower the l value, the lower the sublevel energy; thus…. s < p < d < f

  10. Assigning Electrons to Atoms • Electrons generally assigned to orbitals of successively higher energy. • For many-electron atoms, energy depends on both n and l. • In H atom all subshells of same n have same energy. • In many-electron atom: • subshells increase in energy as value of (n + l) increases. • b) for subshells of same • (n + l), the subshell with • lower n is lower in energy.

  11. Electron Filling Order n values are constant horizontally l values are constant vertically combined values of n+1 are constant diagonally

  12. Orbital occupancy for the first 10 elements, H through Ne Figure 1 He and Ne have filled outer shells: confers chemical inertness

  13. Cr and Cu: Half-filled and filled sublevels are unexpectedly stable!

  14. Effective Nuclear Charge, Z* • Z* is the nuclear charge experienced by the outermost electrons. • Z* increases across a period owing to incomplete shielding by inner electrons. Atom Z* Experienced by Electrons in Valence Orbitals • Li +1.28 • Be ------- • B +2.58 • C +3.22 • N +3.85 • O +4.49 • F +5.13 Increase in Z* across a period

  15. Smaller orbitals. Electrons held more tightly. General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity Higher effective nuclear charge. Electrons held more tightly

  16. Atomic Size • Size goes UP on going down a group. • Because electrons are added farther from the nucleus, there is less attraction. • Size goes DOWN on going across a period.

  17. Atomic Radii Figure 8.9

  18. Ion Sizes Does the size go up or down when losing an electron to form a cation?

  19. + + Li , 78 pm 2e and 3 p Ion Sizes Forming a cation. • CATIONS are SMALLER than the atoms from which they come. • The electron/proton attraction has gone UP and so size DECREASES. Li,152 pm 3e and 3p

  20. Ion Sizes Does the size go up or down when gaining an electron to form an anion?

  21. - - F, 71 pm F , 133 pm 9e and 9p 10 e and 9 p Ion Sizes Forming an anion. • ANIONS are LARGER than the atoms from which they come. • The electron/proton attraction has gone DOWN and so size INCREASES. • Trends in ion sizes are the same as atom sizes.

  22. Trends in Ion Sizes Figure 8.13

  23. Ionization EnergySee Screen 8.12 IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg+ (g) + e-

  24. Ionization EnergySee Screen 8.12 Mg (g) + 735 kJ ---> Mg+ (g) + e- Mg+ (g) + 1451 kJ ---> Mg2+ (g) + e- Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + e- Energy cost is very high to dip into a shell of lower n. This is why ox. no. = Group no.

  25. Trends in Ionization Energy • IE increases across a period because Z* increases. • Metals lose electrons more easily than nonmetals. • Metals are good reducing agents. • Nonmetals lose electrons with difficulty.

  26. Trends in Ionization Energy • IE decreases down a group • Because size increases. • Reducing ability generally increases down the periodic table. • See reactions of Li, Na, K

  27. Electron Affinity A few elements GAIN electrons to form anions. Electron affinity is the energy change when an electron is added: A(g) + e- ---> A-(g) E.A. = ∆E

  28.     -  [He] O ion + electron       [He] O atom Electron Affinity of Oxygen ∆E is EXOthermic because O has an affinity for an e-. EA = - 141 kJ

  29.     [He] N atom + electron   N- ion    [He] Electron Affinity of Nitrogen ∆E is zero for N- due to electron-electron repulsions. EA = 0 kJ

  30. Trends in Electron Affinity • Affinity for electron increases across a period (EA becomes more negative). • Affinity decreases down a group (EA becomes less negative). Atom EA F -328 kJ Cl -349 kJ Br -325 kJ I -295 kJ

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