Unit 5: Electron Configuration, Periodic Table Chapter 3 & 4

Key Concepts Orbitals, Energy Levels, Energy Sublevels Electron Configuration Noble Gas Shortcut s, p, d, f Blocks Early Periodic Table, Modern Periodic Table Periodic Law Periods, Groups Valence Electrons, Octet/Duet Rules, Ions Alkali Metals, Alkaline Earth Metals, Halogens, Noble Gases Metals, Nonmetals, Transition Metals Periodic Table Trends

Orbitals Atomic Orbitals: where electrons are likely to be found in an atom Energy Levels are the size of orbitals Energy Sublevels are the shapes of orbitals

Energy Levels (size of orbital) Lower energy levels are smaller and closer to the nucleus Higher energy levels are bigger and farther away from the nucleus IR Visible Light Electrons in higher energy levels have greater potential energy. Dropping to n=1 (UV) releases more energy than dropping to n=2 (visible light) releases more energy than dropping to n=3 (IR), because of the difference in energies/distance. [x-ray and gamma ray are radiation particles that leave the nucleus, not electrons in energy levels] UV

Energy Sublevels (shape of orbitals) p, d, and f orbitals have multiple orientations in space s orbital p orbital d orbital f orbital 1 orientation 3 orientations 5 orientations 7 orientations

Energy Sublevels (shape of orbitals) Orbitals Website http://www.ptable.com/#Orbital Orbitals Website: http://www.ptable.com/#Orbital

Orbitals Energy Levels: 1, 2, 3, 4… Energy Sublevels: s, p, d, f

Orbitals 2 electrons fit in each orbital! Energy Level (size of orbitals) Energy Sublevel (shape of orbitals) Number of Orbitals Number of Electrons 1 s 1 total 2 e- 2 s, p 1, 3 = 4 total 8 e- 3 s, p, d 1, 3, 5 = 9 total 18 e- 4 s, p, d, f 1, 3, 5, 7 = 16 total 32 e- 2 electrons fit in each orbital!

Electron Configuration Electron Configuration: the energy levels and sublevels that electrons exist in Ex. The electron configuration of neon is written Ne: 1s2 2s2 2p6 e- always fill up orbitals in the same order! (lower energy level  higher energy level)

Electron Configuration Write out the electron configuration of neon: How many electrons does the element have? Ne has 10 e- Fill the orbitals with electrons Ne: 1s2 2s2 2p6 2 e- in orbital 1s, 2 e- in orbital 2s, & 6 e- in orbital 2p = 10 e- total

Electron Configuration Write the electron configuration for the following elements O K+ Mg S2- O (8 e-): 1s2 2s2 2p4 Mg (12 e-): 1s2 2s2 2p6 3s2 K+ (18 e-): 1s2 2s2 2p6 3s2 3p6 S2- (18 e-): 1s2 2s2 2p6 3s2 3p6 (Should be 3p4, but because 2 added e- it is 3p6)

Noble Gas Shortcut Elements gain or lose e- until they’re isoelectronic (have the same e- configuration) with a noble gas Noble Gas Shortcut: put the closest noble gas before an element in [brackets] to shorten its electron configuration Ex. What is the noble gas shortcut of phosphorus? P: [Ne] 3s2 3p3

Periodic Table

Periodic Table Rows = Energy Levels Blocks = Energy Sublevels 3 2 1 4 5 6 7 1 3 2 4 5 6 3 4 5 6 7 4 5 Rows = Energy Levels Blocks = Energy Sublevels

Periodic Table 7p Rows = Energy Levels Blocks = Energy Sublevels

Early Periodic Table Mendeleev: invented the first periodic table Grouped the known elements together according to their properties Put elements in order of their atomic mass Mendeleev left gaps for undiscovered elements. For example, he knew all the elements in group 4 except for the one between Silicon and Tin, and he even made predictions about its properties. When Germanium was discovered, his predictions were confirmed.

Modern Periodic Table Elements are now organized by their atomic number (# of protons), not their atomic mass

Periodic Law Periodic Law: when elements are arranged by their atomic numbers, elements with similar properties appear at regular intervals Elements in the same columns have similar properties! Periodic Law: elements with similar properties appear “periodically”

Period vs. Group Period: horizontal row on the periodic table Group: vertical column on the periodic table

Groups s Block p Block d Block f Block Alkali Metals: Group 1 Alkaline Earth Metals: Group 2 p Block Other metals, nonmetals, & metalloids: Groups 3-6 Halogens: Group 7 Noble Gases: Group 8 d Block Transition Metals f Block Rare Earth Elements

Groups

Valence Electrons Valence Electrons: electrons in the outermost shell/orbital of an atom Valence electrons of an atom participate in chemical reactions with other atoms Elements with the same number of valence electrons tend to react in the same way C H

Octet Rule Octet Rule: atoms gain or lose e- to attain an e- configuration of the nearest noble gas Noble gases have complete outer e- shells, which make them stable! Duet Rule is for H & He only Octet Rule is for all other elements

Ions +1 +2 +3 ±4 -3 -2 -1 TM form ions with various charges When an atom/molecule gains a negative electron, it becomes a negative ion or an anion. When an atom/molecule loses a negative electron, it becomes a positive ion or cation. When an atom or molecule gains a negative electron, it becomes a negative ion When an atom or molecule loses a negative electron, it becomes a positive ion

Alkali Metals Alkali Metals: Highly reactive, reacts with water Loses 1 valence e- to get to stable e- configuration Stored in oil so they don’t react with the moisture in the air

Alkaline Earth Metals Alkaline Earth Metals: Highly reactive, reacts with oxygen Loses 2 valence e- to get to stable e- configuration Slightly less reactive than alkali metals because it takes more energy to lose 2 e- than 1 e-

Halogens Halogens: Highly reactive, reacts with metals making salts Gains 1 valence e- to get to stable e- configuration

Noble Gases Noble Gases: stable, not reactive Full set of e- in their outermost energy level

Metals vs. Nonmetals

Metals vs. Nonmetals Metals Nonmetals Left side of periodic table Shiny Malleable/ductile Conductive Solid at room temp Forms ions with positive charge Right side of periodic table Dull Brittle Not conductive Gas at room temp Forms ions with negative charges Metalloids: along staircase between metals & nonmetals Both metallic and non-metallic properties

Transition Metals Same properties as the other metals with a few additional properties Forms ions with various positive charges Forms colorful compounds Good catalysts Catalysts accelerate a chemical reaction without being affected by it. They aren’t actually part of the chemical reaction, they just speed it up. Catalysts: -Manganese dioxide -Iron oxide -Copper oxide Catalyst

Transition Metals Fe(SO4) Fe(NO3)3 Work out the charges of iron in both compounds. Fe2- is green and Fe3- is orange. Fe(SO4) Fe(NO3)3

Transition Metals + Potassium dichromate Sodium sulfite The color change indicates a change in charge. Potassium dichromate Sodium sulfite +

Periodic Table Trends Radius: size of atom Electronegativity: attraction of e- in a bond Ionization Energy: energy required to remove e- When atoms are smaller, their e- are closer

Periodic Table Trends Radius: size of atom Electronegativity: attraction of e- in a bond Ionization Energy: energy required to remove e- Radius Decreases Electronegativity Increases Ionization Energy Increases When atoms are smaller, their e- are closer

Periodic Table Trends When atoms are smaller, their e- are closer to the nucleus so it takes more energy to remove them. When atoms are bigger, their e- are farther from the nucleus so it takes less energy to remove them. When atoms are smaller, their e- are closer

Periodic Table Trends Ex. Compare barium and sulfur’s radius, electronegativity, and ionization energy. Ex. What is the largest element on the periodic table? What is the smallest element on the periodic table? Radius Decreases Electronegativity Increases Ionization Energy Increases