Orbital Diagrams and Electron Configuration

1. Orbital Diagrams and Electron Configuration • Drawing orbital diagrams gives information not only about the orbitals that are/have been filled but also about the number of unpaired electrons. • Orbital diagrams can be cumbersome!!

2. Electron Configuration • A short-hand notation is commonly used in place of orbital diagrams to describe the electron configuration of an atom. • Electron configuration: • a particular arrangement of electrons in the orbitals of an atom

3. Electron Configuration • The electron configuration tells the number of electrons found in each subshell. • If there are three electrons in a 2p subshell, we would write: 2p3 where the superscript (3) indicates the number of electrons in that subshell

4. 1s 2s 2p 3s Electron Configuration • The orbital diagram for an O atom: The electron configuration for an O atom: 1s22s22p4

5. Electron Configuration • To determine the electron configuration of an atom (or ion) without first writing the orbital diagram: • determine the number of electrons present • add electrons to each subshell in the correct filling order until all electrons have been added • use the “diagonal” diagram to help determine the filling order

6. Electron Configuration Example: Write the electron configuration of a Mn atom (Z = 25). 1s2 2s22p6 3s23p6 4s2 3d5

7. Electron Configuration Example: Write the electron configuration of an O2- ion (Z = 8). An O2- ion has 8 protons and 10 electrons 1s22s22p6

8. Electron Configuration Write the electron configuration of a krypton atom (Z = 36). 1s22s22p63s23p64s23d104p6 This is the Kr “core”  [Kr] • The noble gas “core” can be used to write the electron configuration of an elementusing core notation: noble gas “core” + valence electrons

9. Core notation To write the electron configuration using the core notation: • Find the Noble Gas that comes before the atom. • Determine how many additional electrons must be added beyond what that noble gas has. (= Atomic number of atom minus atomic number of noble gas) • Determine the period that element is in. (This determines the value of n of the s subshell to start with when adding extra electrons) • Add electrons starting in that “n” subshell

10. Electron Configuration Write the core electron configuration of Sr (Z = 38). Previous noble gas: Kr (Z = 36) Extra electrons: 38 (e of Sr) - 36 = 2 Period number of Sr: 5 So: Kr core plus 2 extra e- starting in 5s [Kr] 5s2

11. Electron Configuration Write the core electron configuration of Br (Z = 35). Previous noble gas: Ar (Z = 18) Extra electrons: 35 - 18 = 17 Period number: 4 So: Ar core plus 17 extra e- starting with 4s [Ar] 4s23d104p5

12. Isoelectronic Series • When atoms ionize, they form ions with the same number of electrons as the nearest (in atomic number) noble gas. Na = 1s22s22p63s1 = [Ne]3s1 Na+ = 1s22s22p6 = [Ne] Cl = 1s22s22p63s23p5 = [Ne]3s23p5 Cl- = 1s22s22p63s23p6= [Ar]

13. Isoelectronic Series • N (7 e-): 1s22s22p3 • O (8 e-): 1s22s22p4 • F (9 e-): 1s22s22p5 • N3- (10 e-): 1s22s22p6 = [Ne] • O2- (10 e-): 1s22s22p6 = [Ne] • F- (10 e-): 1s22s22p6 = [Ne]

14. Isoelectronic Series • Na (11 e-): 1s22s22p63s1 • Mg (12 e-): 1s22s22p63s2 • Al (13 e-): 1s22s22p63s23p1 • Na+ (10 e-): 1s22s22p6 = [Ne] • Mg2+ (10 e-): 1s22s22p6 = [Ne] • Al3+ (10 e-): 1s22s22p6 = [Ne]

15. 1A Ionsof the highlighted elements are isoelectronicwith Ne. 8A H 2A He 3A 4A 5A 6A 7A Li Be B C N O F Ne Na Mg Al Si P S Cl Ar 7B 8B 8B 8B 1B 2B 3B 4B 5B 6B K Ca Sc Ti V Cr Fe Co Ni Cu Zn Ga Ge As Se Br Kr Mn Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Md No Lr Fm

16. Isoelectronic Series • Isoelectronic:having the same number of electrons • N3-, O2-, F-, Ne, Na+, Mg2+, and Al3+ form an isoelectronic series. • A group of atoms or ions that all contain the same number of electrons

17. Isoelectronic Series • Examples of isoelectronic series: • N3-, O2-, F-, Ne, Na+, Mg2+, Al3+ • Se2-, Br-, Kr, Rb+, Sr2+, Y3+ • Cr, Fe2+, and Co3+

18. Periodic Properties of Elements • Chemical and physical properties of the elements vary with their position in the periodic table. • Atomic size • Size of Atom vs. Ion • Size of Ions in Isoelectronic series • Ionization energy • Electron affinity • Metallic character

19. What Is the Size of an Atom? The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

20. Periodic Properties--Atomic Size • The relative size (radius) of an atom of an element can be predicted by its position in the periodic table. • Trends • Within a group (column), the atomic radius tends to increase from top to bottom • Within a period (row), the atomic radius tends to decrease as we move from left to right

21. Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row …increase from top to bottom of a column

22. Periodic Properties--Atomic Size • Example: Which element would have the larger atomic radius, Ar or Br? • Br should have the larger radius • more towards the bottom • more towards the left side

23. Periodic Properties – Atom vs. Ion Size • Trends to know: • Cations (+) are smaller than their parent atoms. • Electrons are removed from the outer shell. • Anions (-) are larger than their parent atoms. • Electron-electron repulsion causes the electrons to spread out more in space.

24. Sizes of Ions - Trends • Ions increase in size as you go down a column.

25. Sizes of Ions - Trends • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge.

26. Ionization Energy • The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion to form a cation or more positively charged cation. • The first ionization energy is that energy required to remove first electron. • The second ionization energy is that energy required to remove second electron, etc.

27. Ionization of Gaseous Sodium: Na (g)  Na+ (g) + e- • As the ionization energy increases, it becomes harder to remove an electron.

28. Ionization Energy • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap.

29. Trends in First Ionization Energies • As one goes down a column, less energy is required to remove the first electron. • valence electrons are farther from the nucleus. Within each row, the ionization energy increases from left to right

30. Ionization Energy Which element has the higher ionization energy, Br or Ca? Which one will lose an electron easier? • Br has the higher ionization energy • further to the right • Ca will lose an electron easier because its ionization energy is lower.

31. Periodic PropertiesElectron Affinity • The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity. Cl (g) + e- Cl- (g) • The electron affinity becomes increasingly negative as the attraction between an atom and an electron increases • more negative electron affinity = more likely to gain an electron and form an anion

32. Electron Affinity • Trends: • Halogens have the most negative electron affinities. • Electron affinities become increasing negative moving from the left toward the halogens. • Electron affinities do not change significantly within a group. • Noble gases will not accept another electron. • To do so would require adding an electron to a new electron shell (significantly higher in energy)

33. Metallic Character • Metals: • shiny luster • malleable and ductile • good conductors of heat and electricity • form cations • Metallic character • increases from top to bottom • Increases from right to left

34. Properties of Metal, Nonmetals,and Metalloids

35. Metals versus Nonmetals Differences between metals and nonmetals tend to revolve around these properties.

36. Metals versus Nonmetals • Metals tend to form cations. • Nonmetals tend to form anions.

37. Metals They tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

38. Metals • Compounds formed between metals and nonmetals tend to be ionic. • Metal oxides tend to be basic.

39. Nonmetals • These are dull, brittle substances that are poor conductors of heat and electricity. • They tend to gain electrons in reactions with metals to acquire a noble gas configuration.

40. Nonmetals • Substances containing only nonmetals are molecular compounds. • Most nonmetal oxides are acidic.

41. Metalloids • These have some characteristics of metals and some of nonmetals. • For instance, silicon looks shiny, but is brittle and fairly poor conductor.