Atomic Orbitals & Electron Configurations

1. Atomic Orbitals & Electron Configurations Chemistry

2. Atomic Orbitals & Electron Configurations 1. Identify the relationships among a hydrogen atom’s energy levels, sublevels, and atomic orbitals. 2. Apply the Pauli exclusion principle, the aufbau principle, and Hund’s rule to write electron configurations using orbital diagrams and electron configuration notation. 3. Define valence electrons and draw electron-dot structures representing an atom’s valence electrons. For the Test: -be able to answer the objectives -know all vocabulary -be able to answer all review/practice questions -study your notes (anything can be asked from them)

3. Atomic Orbitals Orbitals do not have an exactly defined size, but are just unoccupied spaces available for electrons should the atom’s energy increase (or decrease).. As the energy level increases, the orbital becomes larger -n = 1, 2, … 7 The shape of the orbital is represented by the sublevels in that orbital -s, p, d, & f

4. Atomic Orbitals Each atomic orbital is designated by energy level and sublevel: 1s represents ‘s’ orbital in the 1st energy level each suborbital holds 2 electrons, each with opposite spin ~ s = 2 ~ p = 6, 3 suborbitals (2 per suborbital) ~ d = 10, 5 suborbitals (2 per suborbital) ~ f = 14, 7 suborbitals (2 per suborbital)

7. Electron Configurations electron configuration: arrangement of electrons in an atom. -electrons fill in so they have the lowest possible energy -there are three rules we follow 1. Aufbau principle: each electron occupies the lowest energy orbital available -follows the aufbau diagram ♦all orbitals in a sublevel are equal energy ♦in multi-electron level, energy sublevels within a principle energy level have different energies

8. ♦in order of increasing energy, the sequence of energy sublevels within a principle energy level is s, p, d, and f ♦orbitals related to energy sublevels within one principle level can overlap orbitals related to energy sublevels within another principle level; begins in n = 3 2. Pauli-exclusion principle: a maximum of 2 electrons may occupy a single orbital, but only if they have opposite spins -arrows pointing up (↑)and down (↓) represent electron spins -paired electrons in same orbital: represented by ↑↓.

9. 3. Hund’s rule: single electrons with the same spin must occupy each equal energy orbital before additional electrons with opposite spins can occupy the same orbital -s sublevel: max 1 unpaired e- p sublevel: max 3 unpaired e- d sublevel: max 5 unpaired e- f sublevel: max 7 unpaired e-

10. Orbital Notation/Electron Configurations There are three ways to indicate electron configurations. 1. orbital notation: an unoccupied orbital is represented by a line, ____, with the orbital’s principle quantum number and sublevel letter written underneath the line. ___ 1s -an unpaired electron is shown as ____ and 1s paired electrons as shown as ____ . 1s

11. Orbital Notation Practice • Mg • Al • Si • P • S • Cl • Ar • Cr • As • Kr

12. 2. electron-configuration notation: eliminates the lines and arrows of orbital notation -the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation; exs. H configuration is 1s1 (1 electron on the 1s orbital) He is 1s2 (2 electrons in the 1s orbital) -use the periodic table to write electron configuration notation: s, p, d, and f block. -does not show orbital distributions of electrons

13. Electron Configuration Notation Practice • Mg • Al • Si • P • S • Cl • Ar • Cr • As • Kr

14. Noble Gas Notation 3. noble-gas notation: shorthand notation using the noble gas elements -noble-gas configuration: outer main energy level fully occupied by eight electrons, except in He -ex: Na 1s22s22p63s1 is the electron-configuration notation (long form) [Ne]3s1 is the noble gas notation (short form) because all orbitals are completely filled up to neon, who’s orbital notation is 1s22s22p6 -this notation is only used for elements above neon! (those in n = 3 and above)

15. Noble Gas Notation Practice • Mg • Al • Si • P • S • Cl • Ar • Cr • As • Kr

16. Exceptions Some transition metals do not follow this trend. -includes groups 6 and 11 (Cr and Cu groups) Cr [Ar]4s23d4 Cu [Ar]4s23d9 Instead, we write it so that all s and d orbitals after the noble gas are half filled-increases stablity. Cr [Ar]4s13d5 Cu [Ar]4s13d10 Practice with Exceptions: 1. Mo 2. Au

17. Valence Electrons Only certain electrons determine an element’s chemical properties -valence electrons: electron’s in the atom’s highest most energy level •sulfur has 16 electrons, but only 6 valence electrons Valence electrons can be determined by writing electrons configurations: -S [Ne]3s23p4 6 -Ga [Ar]4s23d104p1 3 (4th level highest) [Ar] 3d104s24p1

18. Electron Dot Structures Chemists use valence electrons to show how atom’s are involved in bonding. -electron dot structure: consists of the element’s symbol (which represents the nucleus and inner electrons) surrounded by dots (that represent the valence electrons) •unpair the electrons first along the four sides of the symbol S

19. Oxidation Numbers When an atom bonds, it either gains or loses electrons, forming a charge on the atom -oxidation number: charge on an atom after bonding •metals lose to form a positive charge Ca+2 loses two electrons •nonmetals gain to form a negative charge N-3 gains three electrons Look at the electron dot structure to determine how many it will lose or gain.