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Unit 3 Electrons and the Periodic Table

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Unit 3 Electrons and the Periodic Table UNIT 3 ELECTRONS AND THE PERIODIC TABLE Chapter 5 & 6 CHAPTER 5: ELECTRONS IN ATOMS WHAT ARE THE CHEMICAL PROPERTIES OF ATOMS AND MOLECULES RELATED TO? • ELECTRONS! • Number of valence electrons • How much an atom wants to “gain” or “lose” electrons • How many electrons “gained” or “lost” ELECTRON HISTORY • Dalton’s Model • Didn’t even know of electrons or charge – just ratios of atoms. • Thomson’s Model • ELECTRONS EXIST!!!! • Where are they? No clue. • Plum Pudding? • Rutherford’s Model • Gold Foil Experiment • Electrons are found outside the nucleus. • Now knows where they are, but still doesn’t know how they are organized. Rutherford’s model BOHR’S MODEL • Electron’s are found in energy levels. Quantum: Specific energy values that have exact values in between them. • Pictured the atom like a solar system with electrons “orbiting” around the nucleus. • Each orbit path associated with a particular energy. • Now a better understanding of how electrons are organized, but it still isn’t quite correct. QUANTUM MECHANICAL MODEL ERWIN SCHRODINER (1926) • Electrons are found in regions of empty space around the nucleus. “electron clouds”. • Shape of the region corresponds to different energy levels. • Shape of the region represents the probability of the electron’s location 90% of the time. • PROBLEM? These are still called “atomic orbitals”. SCHRODINGER’S MODEL • Electrons are “wave functions”. • We can’t actually know exactly where an electron is at any moment. • We can know a probability of an electrons location based on its energy. • Ex: a windmill blade ATOMIC ORBITALS • Shape describing the probability of finding an electron at various locations around the nucleus. • 4 main shapes: s, p, d, f • f is too complicated for us to worry about drawing. s: consists of 1 “sub-orbital” p: 3 “sub-orbitals” d: 5 “sub-orbitals” f: 7 “sub-orbitals” Orbitals Overlap! OVERALL ENERGY VALUES Label this information on your mini periodic table. ELECTRON CONFIGURATIONS • How electrons are arranged in various orbitals around the nucleus. • Three rules explain how: Aufbau Principle, Pauli Exclusion Principle, and Hund’s rule. ELECTRON CONFIGURATIONS Aufbau Principle: • Electrons enter orbitals of lowest energy first. • Patterns: Use the Periodic Table. ELECTRON CONFIGURATIONS Pauli Exclusion Principle: • Only 2 electrons per orbital • Each orbital contains electrons of opposite spin. (up arrow, or down arrow) ELECTRON CONFIGURATIONS Hunds Rule: • When electrons enter orbitals of equal energy, all the sublevels are filled with 1 electron of the same spin first. • Example: • P’s (3 total sublevels) • Each of the 3 sublevels will fill with only 1 electron first • Then they will begin to fill with the second electron of opposite spin. ELECTRON CONFIGURATIONS Element 1s 2s 2px 2py 2pz 3s Electron Configuration H He Li C N Na Noble Gas Abbreviation? 1. Sample Problem 5.1: • The atomic number of phosphorus is 15. Write the electron configuration of a phosphorus atom. 1s 2s 2p 3s 3p 4s 3d P = 1s22s22p63s23p3 or [Ne]3s23p3 2. Complete the following Electron Configurations Carbon Sulfur Full Electron Configuration Full Electron Configuration Abbreviated Configuration Abbreviated Configuration 2. Complete the following Electron Configurations Nickel Krypton Full Electron Configuration Full Electron Configuration Abbreviated Configuration Abbreviated Configuration 3. Why do we care about electron configurations? • Chemical properties are related to electrons. • They explain the organization of electrons. 4. What is the expected electron configuration of Mo? • 1s22s22p63s23p64s23d104p65s24d4 Fill this into a detailed orbital diagram. 5. The actual electron configuration for Mo is: • 1s22s22p63s23p64s23d104p65s14d5 Fill this into a detailed orbital diagram 6. Why is Mo’s electron configuration different than expected? • An (s) orbital electron gets “promoted” to the d orbitals • Hund’s Rule: All 5 d orbitals with 1 electron is more “stable” than having 1 of the d orbitals empty. 1s22s22p63s23p64s23d104p65s14d5 7. How many unpaired e-s are in the electron configuration above? • (6) 1s22s22p63s23p64s23d104p65s14d5 8. Which e-s are most energetic? Least energetic? • 5s1 are most energetic (4d5 are close). • All the others are LOWER in Energy. 9. How could you abbreviate the electron configuration above? • [Kr] 5s14d5 10. Why would you abbreviate and leave off the “inner-core” electrons? • They are lower in energy and NOT VALENCE. They are more strongly attracted to the nucleus. 11. What other elements would you predict have electron configurations like Mo? • Cr & W (Same COLUMN) 12.Which electrons are usually closest to the nucleus in the following: 1s22s22p63s23p64s1 • 1s22s22p63s23p6 (the “inner core” energy level electrons) 13.Which electrons are furthest from the nucleus? • 4s1 (the highest principle” energy level electrons) 14.Which electrons would you think would be easiest to remove from the atom above? Why? • 4s1 (The electrons farther from the nucleus are attracted with a smaller force to the nucleus.) 15.What name is given to these “outermost” electrons? • VALENCE electrons IONS AND ELECTRONS • Ions have additional or missing electrons. • These electrons fill or come from the outermost energy level. • Electrons gained or lost are VALENCE ELECTRONS. Anions: Non-metals that gain electrons. Cations: Metals that lose electrons. • Example: O2- Anion • has 10 electrons; gained 2 electrons • 1s22s22p6 • same electron configuration as [Ne] 16. Give the Full & noble gas (abbrev) GROUND STATE electron configuration for the following ions: • O2- • Cu2+ • N3- • Cu1+ (Special Case)***.
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