How do we describe the distribution of electrons in an atom How can we understand the periodic properties of elements Wh

1. Chapter 8 How do we describe the distribution of electrons in an atom? How can we understand the periodic properties of elements? What electrons are involved in chemical reactions? What are some important properties of atoms related to electron configurations? 1

2. 2 Orbital Energies of the Hydrogen Atom

3. 3 Electron Spin Last quantum number describes electron spin. The first three describe the energy, shape and position of the atomic orbital. Electron pairs residing in the same orbital are required to have opposing spins(+1/2 & -1/2). This causes electrons to behave like tiny bar magnets. A beam of hydrogen atoms is split in two by a magnetic field due to these magnetic properties of the electrons. Electron acts as a ball of spinning charge, creating electric and magnetic fields.

4. Nuclear Charge, Repulsion and Penetration Sub-shell energies are split Higher nuclear charge increases attraction between nucleus and electron; orbital energy more stable with higher nuclear charge: Li2+ > He+ > H Shielding by other electrons decreases attraction of outer electron to nucleus Penetration increases overall attraction of the electron to the nucleus and causes splitting of sub-shell energies Order of sub-shell energies: s<p<d<f 4

5. 5 Electron Configuration An electron configuration of an atom is a particular distribution of electrons among available sub-shells. The notation for a configuration lists the sub-shell symbols sequentially with a superscript indicating the number of electrons occupying that sub shell. For example, lithium (atomic number 3) has two electrons in the �1s� sub shell and one electron in the �2s� sub shell: 1s2 2s1.

6. 6 Electron Configuration An orbital diagram is notation used to show how the orbitals of a sub shell are occupied by electrons. Each orbital is represented by a circle. Each group of orbitals is labeled by its sub shell notation. Electrons are represented by arrows: up for ms = +1/2 and down for ms = -1/2

7. 7 The Pauli Exclusion Principle The Pauli exclusion principle, which summarizes experimental observations, states that no two electrons can have the same four quantum numbers. In other words, an orbital can hold at most two electrons, and then only if the electrons have opposite spins.

8. 8 The Pauli Exclusion Principle The maximum number of electrons and their orbital diagrams are:

9. Quiz The maximum number of electrons that can be accommodated in a sub-shell where l = 4 is a. 12 b. 14 c. 16 d. 18 e. 20 9

10. 10 Aufbau Principle Every atom has an infinite number of possible electron configurations. The configuration associated with the lowest energy level of the atom is called the �ground state.� Other configurations correspond to �excited states.�

11. 11 Aufbau Principle The Aufbau principle is a scheme used to reproduce the ground state electron configurations of atoms by following the �building up� order. Listed below is the order in which all the possible sub-shells fill with electrons. You need not memorize this order. As you will see, it can be easily obtained.

12. 12 Order for Filling Atomic Subshells

13. 13 Orbital Energy Levels in Multi-electron Systems

14. 14 Aufbau Principle The �building up� order corresponds for the most part to increasing energy of the subshells. By filling orbitals of the lowest energy first, you usually get the lowest total energy (�ground state�) of the atom. The number of electrons in the neutral atom equals the atomic number, Z.

15. 15 Here are a few examples. Using the abbreviation [He] for 1s2, the configurations are Aufbau Principle

16. 16 With boron (Z=5), the electrons begin filling the 2p subshell. Aufbau Principle

17. 17 With sodium (Z = 11), the 3s sub shell begins to fill. Aufbau Principle

18. 18 Note that elements within a given family have similar configurations. For instance, look at the noble gases. Configurations and the Periodic Table

19. 19 Configurations and the Periodic Table Note that elements within a given family have similar configurations. The Group IIA elements are sometimes called the alkaline earth metals.

20. Practice Problem Which of the following electron configurations represents an excited state? a. He: 1s2 b. Ne: 1s2 2s2 2p6 c. Na: 1s2 2s2 2p6 3s1 d. P: 1s2 2s2 2p6 3s2 3p2 4s1 e. N: 1s2 2s2 2p3 20

21. Quiz Which of the following electron configurations is NOT possible? a. 1s2 2s2 2p3 b. 1s2 2s2 2p6 c. 1s2 2s2 2p2 d. 1s2 2s2 2p6 2d2 e. 1s2 2s2 2p6 3s1 21

22. 22 Configurations and the Periodic Table Electrons that reside in the outermost shell of an atom - or in other words, those electrons outside the �noble gas core�- are called valence electrons. These electrons are primarily involved in chemical reactions. Elements within a given group have the same valence shell configuration. This accounts for the similarity of the chemical properties among groups of elements. Noble gas core: an inner shell configuration resembling one of the noble gases. Pseudo-noble gas core: noble gas core + (n-1)d10 electrons.

25. 25 Configurations and the Periodic Table The following slide illustrates how the periodic table provides a sound way to remember the Aufbau sequence. In many cases you need only the configuration of the outer electrons. You can determine this from their position on the periodic table. The total number of valence electrons for an atom equals its group number.

26. 26 Configurations and the Periodic Table

27. 27 Orbital Diagrams Consider carbon (Z = 6) with the ground state configuration 1s22s22p2. Three possible arrangements are given in the following orbital diagrams.

28. 28 Orbital Diagrams Hund�s rule states that the lowest energy arrangement (the �ground state�) of electrons in a sub-shell is obtained by putting electrons into separate orbitals of the sub shell with the same spin before pairing electrons. Looking at carbon again, we see that the ground state configuration corresponds to diagram 1 when following Hund�s rule.

29. 29 Orbital Diagrams To apply Hund�s rule to oxygen, whose ground state configuration is 1s22s22p4, we place the first seven electrons as follows.

30. 30 Summary Pauli Exclusion principle: no 2 e�s in an atom can have the same four quantum numbers. Aufbau Principle: obtain electron configurations of the ground state of atoms by successively filling subshells with electrons in a specific order. Hunds Rule: the lowest energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before paring them.

32. 32 Magnetic Properties Although an electron behaves like a tiny magnet, two electrons that are opposite in spin cancel each other. Only atoms with unpaired electrons exhibit magnetic susceptibility. A paramagnetic substance is one that is weakly attracted by a magnetic field, usually the result of unpaired electrons. A diamagnetic substance is not attracted by a magnetic field generally because it has only paired electrons

33. 33 Periodic Properties The periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. Atomic radius Within each period (horizontal row), the atomic radius tends to decrease with increasing atomic number (nuclear charge). Within each group (vertical column), the atomic radius tends to increase with the period number.

34. 34 Periodic Properties Two factors determine the size of an atom. One factor is the principal quantum number, n. The larger is �n�, the larger the size of the orbital. The other factor is the effective nuclear charge, which is the positive charge an electron experiences from the nucleus minus any �shielding effects� from intervening electrons.

35. 35 Representation of atomic radii (covalent radii) of the main-group elements.

36. 36 Periodic Properties Ionization energy The first ionization energy of an atom is the minimal energy needed to remove the highest energy (outermost) electron from the neutral atom. For a lithium atom, the first ionization energy is illustrated by: Li(1s22s1) ? Li+(1s2) + e-; IE = 520 kJ/mol

37. 37 Periodic Properties Ionization energy There is a general trend that ionization energies increase with atomic number within a given period. This follows the trend in size, as it is more difficult to remove an electron that is closer to the nucleus. For the same reason, we find that ionization energies, again following the trend in size, decrease as we descend a column of elements.

38. 38 Ionization energy versus atomic number

39. 39

40. 40 Periodic Properties Ionization energy The electrons of an atom can be removed successively. The energies required at each step are known as the first ionization energy, the second ionization energy, and so forth. IE�s increase for successive electrons in at atom because it is harder to remove electrons from + ions

41. 41 Exceptions to Ionization E Trends A IIIA element (ns2np1) has smaller ionization E than the preceding IIA element (one np electron removed more easily than second ns electron. A VIA element (ns2np4) has smaller ionization E than preceding VA element. As a result of repulsion it is easier to remove e from the doubly occupied np orbital of the VI element that from a singly occupied orbital of the preceding VA element.

42. 42 Periodic Properties Electron Affinity: the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion. For a chlorine atom, the first electron affinity is illustrated by:

43. 43 Periodic Properties Electron Affinity The more negative the electron affinity, the more stable the negative ion that is formed. Broadly speaking, the general trend goes from lower left to upper right as electron affinities become more negative. Highest electron affinities occur for halogens, F and Cl

44. 44 The Main-Group Elements The physical and chemical properties of the main-group elements clearly display periodic behavior. Variations of metallic-nonmetallic character. Basic-acidic behavior of the oxides.

45. 45 Group IA, Alkali Metals Largest atomic radii React violently with water to form H2 Readily ionized to 1+ Metallic character, oxidized in air R2O in most cases

46. 46 Group IIA, Alkali Earth Metals Readily ionized to 2+ React with water to form H2 Closed s shell configuration Metallic

47. 47 Group III A Metals (except for boron) Several oxidation states (commonly 3+)

48. 48 Group V A Form anions generally(1-, 2-, 3-), though positive oxidation states are possible Form metals, metalloids, and nonmetals

49. 49 Halogens Form monoanions High electronegativity (electron affinity) Diatomic gases Most reactive nonmetals (F)

50. 50 Periodic Table

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