Shells and Subshells

1. 3s 3s 3s 3p 3p 3p 3d 3d 3d n=3 n=3 n=3 Shells and Subshells • The orbitals in an atom are arranged in shells and subshells. • Shell: all orbitals with the same value of n • Subshell: all orbitals with the same value of both n and l orbital

2. n=3 3s 3p 3d n=2 Energy 2s 2p n=1 1s Subshell Energy • For a hydrogen atom (or an ion containing only 1 electron) all orbitals within the same shell are degenerate.

3. 4p 3d 4s 3p 3s 2p Energy 2s 1s Subshell Energy • For atoms with more than one electron, electron-electron repulsion causes different subshells within the same shell to have different energies. • Within the same shell: • s < p < d < f

4. Subshell Energy • The relative energies of the various subshells can be predicted using the diagonal diagram: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f

5. Arrangement of Electrons • The arrangement of the electrons in an atom can be depicted using three different but related methods: • Orbital diagram • Electron configuration 1s22s22p63s1 • Electron configuration using core notation [Ne]3s1

6. Arrangement of Electrons • Rules for populating orbitals with electrons: • Pauli Exclusion Principle: Each electron in an atom must have a unique set of four quantum numbers n, l, ml, and ms. • In order to put more than one electron in an orbital, electrons must have different values of ms. • i.e. they must have different spins • Maximum of 2 electrons per orbital

7. Arrangement of Electrons • Rules for populating orbitals with electrons: • Aufbau Principle: Electrons are placed in the lowest energy orbital available. • Hund’s Rule: If more than one orbital in a subshell is available, electrons will fill empty orbitals in that subshell first. • Keep electrons unpaired (in an orbital by itself) as long as an empty orbital with the same energy is available.

8. Orbital Diagrams Example: Draw an orbital diagram for each of the following atoms. Hydrogen: Helium: Lithium: Beryllium:

9. Orbital Diagrams Example: Draw an orbital diagram for each of the following atoms. Boron: Carbon: Nitrogen: Neon:

10. 1s 2s 2p Orbital Diagrams • The orbital diagram for Ne: • The 2p subshell is completely filled. • The outermost shell (n=2) contains an octet (8) of electrons.

11. “n”s “n”p Orbital Diagrams • All noble gases except helium have a similar octet of electrons in their outermost shell. • This configuration is exceptionally stable. • Responsible for the unreactive nature of the noble gases. • Main group elements that ionize easily generally do so in a way that gives them the same octet of electrons. where n = period number

12. Orbital Diagrams Example: Draw an orbital diagram for each of the following atoms or ions. Iron: Bromine: Sodium ion:

13. p block s block d block f block Useful Information from the Periodic Table • The period number of the element indicates the highest shell (value of n) that contains electrons for that atom. • An element in the fourth period will have one or more electrons in the n=4 shell. • The location of the atom in the periodic table indicates the subshell where the last e-’s are found.

14. Electron Configuration • A short-hand notation (electron configuration) is commonly used instead of an orbital diagram. • The electron configuration designates: • Each subshell that contains electrons in order of increasing energy • The number of electrons found in the subshell

15. 1s 2s 2p Electron Configuration • The orbital diagram for an oxygen atom: • The electron configuration for an oxygen atom: 1s22s22p4 Notice: no commas between subshells!

16. Electron Configuration • Process for writing an electron configuration: • Determine the number of electrons present • Add electrons to each subshell in order of increasing energy until all electrons have been designated • Use diagonal diagram • Remember the maximum # of e- per subshell: • _s2 • _p6 • _d10 • _f14

17. Electron Configuration Example: Write the electron configuration for each of the following atoms: Titanium Lead

18. Electron Configuration Example: Write the electron configuration for each of the following ions: Oxide ion: Potassium ion:

19. Electron Configuration Using Core Notation • Calcium atoms contain 20 electrons: • The first 18 electrons are arranged exactly as the electrons present in argon. • The argon core: [Ar] • The last two electrons are referred to as valence electrons. • Electrons located in the outermost shell that can be transferred to or shared with another atom during the formation of ions or covalent bonds • Electrons over and above those found in the previous noble gas

20. Electron Configuration Using Core Notation • Argon: • 1s22s22p63s23p6 • Calcium: • 1s22s22p63s23p64s2 • Electron configuration using core notation: • [Ar]4s2 Valence e- [Ar]

21. Electron Configuration Using Core Notation • The electron configuration using core notation contains two components: • the noble gas core • valence electrons • Fe: [Ar]4s23d6 • C: [He]2s22p2

22. Electron Configuration Using Core Notation • To write the electron configuration using core notation: • Find the noble gas that comes before the atom and place its elemental symbol in [ ] • Calculate the number of additional electrons • Atomic # of atom – atomic # noble gas • Determine the period number “n” of the atom and begin placing valence electrons in the “n”s subshell. • Use diagonal diagram to determine order in which subsequent subshells are filled

23. Electron Configuration Using Core Notation Example: Write the electron configuration using core notation for each of the following atoms. Ni: Bi:

24. Electron Configuration Using Core Notation Example: Write the electron configuration using core notation for each of the following ions. Iodide ion: Magnesium ion:

25. Transition Metal Ions • Transition metal ions form when electrons are lost from the parent atom in the following order: • s electrons from outermost shell first • d electrons from previous shell next • Example: • Ti: [Ar]4s23d2 • Ti3+: [Ar]3d1

26. Anomalies Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

27. Anomalies For instance, the core electron configuration for chromium is [Ar] 4s1 3d5rather than the expected [Ar] 4s2 3d4. The core electron configuration for copper is [Ar]4s13d10instead of [Ar]4s23d9.

28. Isoelectronic Series • The following ions contain the same number of electrons (10) as Ne. • These ions are isoelectronic with each other and neon. • Having the same number of electrons • These ions and neon form an isoelectronic series. • A group of atoms and ions with the same number of electrons • Nitride ion • Oxide ion • Fluoride ion • Sodium ion • Magnesium ion • Aluminum ion

29. Isoelectronic Series Example: Which of the following atoms or ions in each group are isoelectronic? Fe2+, Co3+, Mn, Cr Se2-, Br, Kr, Sr2+

30. 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

31. 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

32. Periodic Properties--Atomic Size Lower “lefter” larger Periodic Table Increasing size Increasing size

33. 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.

34. Periodic Properties – Ion Size • Trends to know: • For ions in the same group (same charge), size increases from top to bottom. • Same trend as for the size of parent atoms • I- is larger than F- • For an isoelectronic series of ions, the size decreases with increasing atomic number. • Na+ is smaller than O2-

35. Periodic Properties - Ionization Energy • The ease with which an electron can be removed from an atom to form an ion is an important indicator of its chemical behavior. • Ionization energy:the minimum energy required to remove an electron from the ground state of an isolated gaseous atom or ion. • Formation of cation (+) or more positively charged cation Na (g)  Na+ (g) + e-

36. Periodic Properties - Ionization Energy • As ionization energy increases it becomes harder to remove an electron/form a cation. • Within each row, the ionization energy increases from left to right. • Metals form cations more easily than nonmetals. • Within each group, the ionization energy generally decreases from top to bottom. • It’s easier to form K+ than Li+.

37. Periodic Properties – Electron 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

38. Periodic Properties – 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.

39. Periodic Properties – 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