1. Chem 1151: Ch. 3 Electronic Structure and Periodic Law

2. All elements in our universe categorized in periodic table. – Based on atomic number (number of protons). Arranged in columns (groups) and rows (periods). – Groups have similar properties. – Periods correspond to filling of quantum shells by electrons. http://analytical.wikia.com/wiki/Periodic_table_of_elements

3. Origin of the Periodic Table Periodic Law (1869, Julius Meyer, Dmitri Mendeleev): When all elements are arranged in order of increasing atomic numbers, elements with similar chemical properties will occur at regular periodic intervals. http://web.sbu.edu/chemistry/wier/atoms/meyer.htmlhttp://web.sbu.edu/chemistry/wier/atoms/meyer.html; http://www.chemistry.co.nz/mendeleev.htmhttp://www.chemistry.co.nz/mendeleev.htm

4. Meyer and Mendeleev today… http://photos.lucywho.com/zz-top-photos-t93044.html

5. Electrons in Atoms – Early Theories 1911 - Rutherford (Ernest) model: Low-mass electrons follow circular orbits around the nucleus, (planets circling the sun). – Problem: according to current theory at that time, the atoms would essentially collapse as flying electrons lost energy. – How can the atom remain stable over time? + - - - - - - Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

6. Bohr’s Postulates Bohr proposed that the electron in a hydrogen atom moved in any one of a series of circular orbits around the nucleus. The electron could change orbits only by absorbing or releasing energy. Based on work by Planck, Einstein and Balmer, Niels Bohr proposed two postulates to explain the energy stability of the atom.

7. Quantum Mechanics 1926-1927 Austrian physicist Erwin Schrödinger devised a theory to find wave properties of electrons in atoms. – Quantum mechanics (wave mechanics) mathematical description of the wave properties of submicroscopic particles. Here is what is important! – Paths of electrons cannot be determined accurately. – Location and energy of electrons can be specified by: Shell – Group level (corresponds with period) Subshell - Divisions of shells (s, p, d, f) Orbital – space where e- moves around nucleus – Location is a probability function

8. 1927 – Werner Heisenberg’s Uncertainty Principle The position and momentum of a particle cannot be simultaneously measured with arbitrarily high precision. Due to wave-particle duality: electron behaves more like a wave than a hard sphere. The less mass an object (e.g., particle) has, the more uncertainty http://www.goiit.com/posts/show/813684/atomic-structure-the-de-broglie-relation-804303.htm

9. Shells, Subshells, Orbitals  The number of subshells in a shell = shell number  The first subshell s has 1 orbital. Each successive subshell adds 2 more orbitals (1, 3, 5, 7, etc).  Each orbital can hold only 2 electrons of opposite spin.  An atom with n = 3 also includes all subshells and orbitals for n < 3:  1s, 2s, 2p, 3s, 3p, 3d

10. Atomic Orbitals http://www.chemcomp.com/journal/molorbs.htm s subshell ( l =0) 1 orbital ( m l = 0) p subshell ( l =1) 3 orbitals ( m l = -1, 0, +1) d subshell ( l =2) 5 orbitals ( m l = -2, -1, 0, +1, +2)

11. Shells, Subshells, Orbitals: So What? H and He only have a single 1s orbital. Li and Be have a 2s orbital and a 1s orbital. Na and Mg have 3s, 2s, 1s and 2p orbitals. Ca has what? 1s 2s 1s2s3s 2p

12. Energy associated with shells, subshells, orbitals 1.Energy of electron is most dependent upon the shell number (4 > 3 > 2 >1). – 3s > 2s – 2s > 1s Why? Orbital size increases as n increases (e- moves farther from nucleus). 2.Energy of electron in subshell increases moving from s to f. 3.3d > 3p > 3s > 2p > 2s … 4.Orbitals in the same subshell have the same energy level (degenerate). 1s2s3s 2p Energy

13. Why do atoms fill orbitals the way they do? Aufbau Principle Electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states (e.g. 1s before 2s). Sometimes a low energy subshell has lower energy than upper subshell of preceding shell (e.g., 4s fills before 3d). Pauli exclusion principle QM principle that no two identical fermions (particles with half-integer spin) may occupy the same quantum state simultaneously (why paired electrons have different spin). Hund's rule Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin. 2p 2s 1s Energy

14. Pauli Exclusion Principle Electrons behave as if they spin on an axis. According to the Pauli exclusion principle, only electrons spinning in opposite directions (indicated by ↑ and ↓) can occupy the same orbital within a subshell. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

15. How to Get Aufbau from Periodic Table 1 2 3 4 5 6 7 3 4 5 6 2 3 4 5 6 7 4 5 1 For n=2, can only have 2 subshells (s, p) For n=3, can only have 3 subshells (s, p, d)

16. Chemistry of Valence Shell Electrons Valence Shell: Outermost shell (1, 2, 3, etc.) occupied by e-. These are the e- that interact with e- of other elements to form bonds through overlapping orbitals  this is the chemistry we observe. Group number (IA-VIIA) indicates number of valence shell e- (except for He). Elements with similar number of valence shell e- may share similar properties (ex. Halides, noble gases). Octet Rule: Many elements (e.g., noble gases, carbon, oxygen, etc.) are stabilized when surrounded by 8 electrons (octet) in valence shell.

17. Electron pairs form bonds in valence shell 1s2s 3s 2p Na 1s2s3s 3p Cl 2p Atoms stabilized when outer shell filled with e-. In the formation of NaCl, Na “gives” an e- to Cl to form ionic bond between elements. This results in ions Na + having a filled second shell and Cl - having a filled third shell (2 e- in 3s and 6 e- in 3p.

18. NaCl + Na 1+ 1- EN: 0.9EN: 3.0 Ionic Electron pairs form bonds in valence shell

19. 4HC + C H EN: 2.1EN: 2.5 H H H Covalent Electron pairs form bonds in valence shell

20. How the Periodic Table is useful Different ways to classify elements 1.Electronic Configuration 2.Metals, non-metals, metalloids 3.Property trends http://analytical.wikia.com/wiki/Periodic_table_of_elements

21. Electronic Configuration Describes where all of the electrons reside in the atom by shell and subshell, or shell, subshell and orbital. The last e- added to an atom, which goes to the highest energy level subshell, is called the distinguishing e-. Elements can be classified based on where the distinguishing e- resides (in either s, p, d, or f). – Representative elements Distinguishing e- goes to s or p subshell Includes Noble gases – Mostly unreactive elements with 8 valence shell e- (octet) – Transition elements Distinguishing e- goes to d subshell Includes transition metals – Inner transition elements Distinguishing e- goes to f subshell Lanthanides and actinides

22. Distinguishing Electron The distinguishing electron is the last electron listed in the electronic configuration of the element. http://analytical.wikia.com/wiki/Periodic_table_of_elements

23. Element Classification Representative elements have an s or p distinguishing electron. Transition elements have an d distinguishing electron. Inner-transition elements have an f distinguishing electron. http://analytical.wikia.com/wiki/Periodic_table_of_elements

24. Electronic Configuration Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011 Noble gas configurations can be used to write electronic configurations in an abbreviated form in which the noble gas symbol enclosed in brackets is used to represent all electrons found in the noble gas configuration.

25. Metals, Metalloids & Nonmetals Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

26. Properties of elements change in a systematic way within the periodic table. METALLIC AND NONMETALLIC PROPERTIES Most metals have the following properties: high thermal conductivity, high electrical conductivity, ductility, malleability and metallic luster. Most nonmetals have properties opposite those of metals and generally occur as brittle, powdery solids or as gases. Property Trends The Elements of Group VA(15) nitrogenphosphorous arsenicantimony bismuth Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

27. Metallic Property Trends Elements in the same period of the periodic table become less metallic and more nonmetallic from left to right across the period. Elements in the same group of the periodic table become more metallic and less nonmetallic from top to bottom down the group. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

28. Metalloids Metalloids are elements that form a diagonal separation zone between metals and nonmetals in the periodic table. Metalloids have properties between those of metals and nonmetals, and often exhibit some characteristic properties of each type. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

29. Property trends Atomic radius decreases with addition of protons to nucleus which pulls e- closer to center (Electronegativity). Decreasing radius makes it more difficult to remove e- (ionization energy is energy required to remove e-) Atomic radius increases Ionization energy decreases Electronegativity decreases Small radius e- held tightly Large radius e- held loosely