Periodic Properties of the Elements
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| Periodic Property | A property that is predictable based on an element's position within the periodic table
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| Electron Configuration for an Atom | Show the particular orbital that electrons occupy for that atom
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| Ground State | Lowest energy state
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| Orbital Diagram | Gives information that symbolizes the electron as an arrow and the orbital as a box
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| Pauli Exclusion Principle | No two electrons in an atom have the same four quantum numbers. Implies that each orbital can have a maximum of only two electrons, with opposing spins.
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| Degenerate | A term describing two or more electron orbitals with the same value of n that have the same energy
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| Aufbau Principle | Electrons fill lowest energy orbitals first
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| Hund's Rule | Electrons occupy orbitals singly before pairing. Filling degenerate orbitals, electrons fill them singly first, with parallel spins.
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| s sublevel(l=0) | The s sublevel has only one orbital and can therefore hold only 2 electrons
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| p sublevel(l=1) | The p sublevel has three orbitals and can hold 6 electrons
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| d sublevel(l=2) | The d subelvel has five orbitals and can hold 10 electrons
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| f sublevel(l=3) | The f sublevel has seven orbitals and can hold 14 electrons
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| Valence Electrons | The electrons that are important in chemical bonding. [main group] the electrons in the outermost principal energy level. [Trans elements] the outermost d electrons even though they are not in an outermost principal energy level
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| Core Electrons | Those electrons in a complete principal energy level and those in complete d and f subelevels
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| 1. Summarizing Periodic Table Organization | 1. The periodic table is divisible into four blocks corresponding to the filling of the four quantum sub-levels(s, p, d, and f)..
2. The group number of main-group element is equal to the number of valence electrons for that element..
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| 2. Summarizing Periodic Table Organization | 3. The row number of a main-group element is equal to the highest principal quantum number of that element
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| Periodic Table Columns | The number of columns in a block corresponds to the maximum number of electrons that can occupy the particular sublevel of that block.
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| Lettered Group Number of Main Group Elements | Except for helium, the number of valence electrons for any main group element is equal to its lettered group number
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| Row Number For Main Group Elements | The row number in the periodic table is equal to the number(or n value) of the highest principal-level. For d block, n value is the row number-1. For f block, n value is the row number-2
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| 1. Summarizing Periodic Table Organization | 1. The periodic table is divisible into four blocks corresponding to the filling of the four quantum sublevels(s, p, d, f)..
2. The group number of main-group element is equal to the number of valence electrons for that element..
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| 2. Summarizing Periodic Table Organization | 3. The row number of a main group element is equal to the highest principal quantum number of that element
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| Inner Electron Configuration | The electron configuration of the noble gas that precedes that element on the periodic table
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| Outer Electron Configuration | The configuration of the electrons beyond the previous noble gas. This is obtained by tracing the elements between the previous noble gas and the element under consideration and assigning electrons to the appropriate orbitals
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| The Following are Exceptions for Electron Configurations |
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| Chromium(Cr) | Chromium(Cr): [Ar]4s^1 3d^5
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| Molybdenum(Mo) | Molybdenum(Mo): [Kr]5s^1 4d^10
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| Copper(Cu) | Copper(Cu): [Ar]4s^1 3d^10
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| Silver(Ag) | Silver(Ag): [Kr]5s^1 4d^10
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| Gold(Au) | Gold(Au): [Xe]6s^1 5d^10
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| An Elements Valence Electrons | The chemical properties of elements are largely determined by the number of valence electrons they contain
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| van der Waals Radius(Nonbonding Atomic Radius) | One-half the distance between the center of adjacent, nonbonding atoms in a crystal
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| Bonding Atomic Radius(Covalent Radius) for Nonmetals | One-half the distance between two of the atoms bonded together
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| Bonding Atomic Radius(Covalent Radius) for Metals | One-half the distance between two of the atoms next to each other in a crystal of the metal
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| Atomic Radius | A set of average bonding radii determined from measurements on a large number of element and compounds. Represents the radius of an atom when it is bonded to another atom and is always smaller than the van der Waals radius
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| Periodic Trends in Atomic Radius | 1. As we move down a column(family) in the per. table, atomic radius increases..
2. As we move right across a period(row)in the per. table, atomic radius decreases..
3. As we move from top right to bottom left of the per. table, atomic radius increases
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| Effective Nuclear Charge | The average or net charge experienced by an electron
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| Effective Nuclear Charge Experience by an Electron Equation | The actual nuclear charge(Z) minus the charge shielded b other electrons(S)...
Zeff = Z - S..
(Effective Nuclear Charge) = (actual nuclear charge) - (Charge screened by other electrons)
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| Trend in Atomic Radius with Two Types of Shielding | Core electrons efficiently shield electrons in the outermost principal energy level from nuclear charge, but outermost electrons do not efficiently shield one another from nuclear charge
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| 1. Summarizing Atomic Radii for Main-Group Elements | 1. As we move down a column in the periodic table, the principal quantum number(n) of the electrons in the outermost principal energy level increases, resulting in larger orbitals and therefore larger atomic radii
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| 2. Summarizing Atomic Radii for Main-Group Elements | 2. As we move --> across row in the per. table, the effective nuclear charge(Zeff) experience by the electrons in the outermost princ. energy level increases, making stronger attraction between outermost electrons and the nucleus, and smaller atomic radii
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| Electron Configurations for Anions | We add the number of electrons indicated by the magnitude of the charge of the anions
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| Electron Configurations for Cations | We subtract the number of electrons indicated by the magnitude of the charge
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| Electron Configurations for Transitional Metal Cation | We remove the electrons in the highest n-value orbitals first, even if this does not correspond to the reverse order of filling
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| Paramagnetic | The state of an atom or ion that contains unpaired electrons(not all spin pairred) and is, therefore, attracted to external magnetic fields
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| Diamagnetic | The state of an atom or ion that contains only parried electrons(all spin parried) and is, therefore, slightly repelled by an external magnetic field
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| Ionic Radii of Cations | Cations are much smaller than their corresponding atoms
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| Ionic Radii of Anions | Anions are much larger than their corresponding atoms
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| Isolelectronic | Ions with the same number of electrons
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| Ionic Radii Trend | Cations(less e^-)< Neutral Atom < Anion(more e^1)
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| For Determining Largest Ionic Radii | The largest ion will be the one with the highest energy level which means it will have the greatest number of electrons. If the number of electrons is tied due to ions, the one with the greatest net electron count is the largest
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| Ionization Energy(IE) | The energy required to remove an electron from the atom or ion in the gaseous state. Is always positive because removing an electron always takes energy...
(Lowest energy)IE1 < IE2 < IE3 < ...(Highest Energy)
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| First Ionization Energy(IE1) | The energy required to remove the first electron from an atom or ion...
Na(g) --> Na^+(g) + 1e^-
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| Second Ionization Energy(IE2) | The energy required to remove the second(not two electrons)electron from the atom of ion...
Na^+(g) --> Na^2+(g) + 1e^-
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| Third Ionization Energy(IE3) and so on... | The energy required to remove the third(not three electrons) electron from an atom or ion
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| Periodic Trends in Ionization Energy | Ionization energy increases as we move to the right across a period(row) and increases as we move up a column(family) in the periodic table. Increases from bottom left to top right of the periodic table
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| 1. Summarizing Ionization Energy(IE) for Main-Group Elements | 1. Ionization energy generally decreases as we move down a column(family) in the periodic table because electrons in the outermost principal level are increasingly farther away from the positively charge nucleus and are therefore held less tightly
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| 2. Summarizing Ionization Energy(IE) for Main-Group Elements | Ionization energy generally increases as we mover to the right across a row(period) in the periodic table because electrons in the outermost principal energy level generally experience a greater effective nuclear charge(Zeff)
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| Electron Affinity(EA) | The measure of how easily an atom or ion will accept an additional electron. The energy change associated with the gaining of an electron by the atom in the gaseous state. Can be positive of negative
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| Periodic Trends with Electron Affinity(EA) | Electron affinity becomes more negative as we move to the right across a row. Electron affinity becomes more positive as we mover down the 1A column
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| 1. Summarizing Electron Affinity for Main-Group Elements | 1. Most groups(column) of the periodic table do not exhibit any definite trend in electron affinity. Among the group 1A metals, however, electron affinity becomes more positive as we move down the column(adding an electron becomes less exothermic)
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| 2. Summarizing Electron Affinity for Main-Group Elements | 2. Electron affinity generally becomes more negative(adding an electron becomes more exothermic) as we move to the right across a period(row) in the periodic table
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| 1. Periodic Trends for Metallic Character | 1. As we move to the right across a row(period) in the periodic table, metallic character decreases..
2. As we move down a column(family) in the periodic table, metallic character increases..
3. Metallic character increases from top right to bottom left
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| 2. Periodic Trends for Metallic Character | As we move down group 5A in the periodic table, metallic character increases. As we move to the right in row 3, metallic character decreases
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