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CHEM 101 Chapter 7

by: Lyna Nguyen

CHEM 101 Chapter 7 Chem 101

Marketplace > Texas A&M University > Chemistry > Chem 101 > CHEM 101 Chapter 7
Lyna Nguyen
Texas A&M
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Combines lecture + textbook
General Chemistry 1
Dr. Daniel Collins
Class Notes
Chemistry, Chem, CHEM 101, tamu




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This 4 page Class Notes was uploaded by Lyna Nguyen on Wednesday February 3, 2016. The Class Notes belongs to Chem 101 at Texas A&M University taught by Dr. Daniel Collins in Fall 2015. Since its upload, it has received 40 views. For similar materials see General Chemistry 1 in Chemistry at Texas A&M University.


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Date Created: 02/03/16
10/29/15 ­ 11/03/15 Chemistry 101  Chapter 7  The Pauli Exclusion Principle o Wolfgang Pauli in 1925 o Exclusion Principle: No more than 2 electrons can be assigned to the same orbital, and if there are 2 electrons in the same orbits, they must have opposite signs  Unique wave function o General Statement: no 2 electrons in an atom can have the same set of quantum numbers o Orbital Box Diagram: electron is in orbital 1s  Box with label and arrow representing electron o Helium with Pauli Exclusion  2 electrons in 1s orbital o Max number of electrons in a shell = 2n 2  Atomic Subshell Energies and Electron Assignments o Aufbau principle: procedure by which electrons are assigned to orbitals  Lowest energy first  “Ground up” o Order of Subshell Energies and Assignments  Subshell energies in multi-electron atoms depend on both n and l  2 rules that help predict arrangements  Electrons are assigned to subshells in order of increasing “n + l” value  For 2 subshells with the same value of “n + l” electrons are assigned first to the subshell of lower “n”  No electron is equal o Quantum numbers: high probability of finding an electron  Shell: electrons with the same “n” value  Subshell: electrons with the same “n” and “l” values  Orbital: electrons with the same “n”, “l”, and “m” value  Same energy in subshell o Effective Nuclear Charge (Z)  Effective Nuclear Charge: net charge experienced by a particular electron in a multi-electron atom resulting rom a balance of the attractive force of the nucleus and the repulsive forces of other electrons  “Positive charge” felt by every electron  On period table:  Top -> bottom: increase  Left -> right: increase 1 10/29/15 ­ 11/03/15  Electron Configurations of Atoms o Electron configuration: arrangements of electrons  How electrons are distributed among the various atomic orbitals o Manipulate outside electrons o Ground state electron configurations: where electrons are found in the shells, subshells, and orbital that result in lowest energy for an isolated atom of the element in the gaseous state  In general, assigned to orbitals in order of increasing “n + l” value o Isoelectronic: have the same number of electrons and hence the same ground state electron configuration  By adding or subtracting electrons o Stability of Orbitals  Hierarchy of stability:  Empty > full > half filled (hund compliant)  Partially filled = most unstable  Anomalous Series  Ex: Copper 2 9 1 10 o [Ar] 2s 3d -> [Ar] 2s 3d o Filled and partially filled -> half full and full  For stability o Electron Configurations of the main group elements  Lithium (Li) and other elements of group 1A  Noble gas notation: abbreviated form by writing in brackets the symbol preceding the element with spdf notation  Core electrons: electrons included in the noble gas notation  Valence electrons: beyond the core o Determine chemical properties  Beryllium and other elements of group 2A  S-block elements: because group 1A and 2A have electron configurations of ns and ns 2  Boron and other elements of group 3A  P-block elements: group 3A-7A  Carbon and other elements of group A  When assigning to orbitals, each successive electron is assigned to a different orbital with the same spin until subshell is half full  Nitrogen (N) and Oxygen and elements of group 5A and 6A  5A = 5 valence electrons  6A = 6 valence electrons  Elements of Period 3  P after s  Electron configuration of transition elements  Transition elements: partially filled d shells  Inner transition elements: fill f subshell 2 10/29/15 ­ 11/03/15 o Lanthanides and actinides  Lanthanides and actinides o Assigned to f orbital first o 14 electrons in f orbital o Hund’s Rule: the most stable arrangement of electrons in a subshell is that which the maximum number of unpaired electrons, all with the same direction spin  Prefer unpaired electrons before pairing up  Rule of subshell o P-block is the representative elements  Electron Configuration of ions o Anions and cations  Anion: one or more valence electrons are added; next noble gas  Cation: one or more valence electron removed; previous noble gas  Always removed first from highest “n” o Max “l” if n = n o Diamagnetism and Paramagnetism  Paramagnetic: elements and compounds with unpaired electrons are attracted to a magnet  Diamagnetic: substances in which all electrons are paired experience a slight repulsion when subjected to a magnetic field o Rule of thumb: add electrons using the atomic number/trend of periodic table  Atomic Properties and Periodic Trends o Realized similarities in properties of the elements are the result of similar valence shell electron configurations o Atomic size (radii):  Left -> right: decrease  Top -> bottom: increase  For a main group element, radii generally increases down a group and decrease across a period  Exceptions: Ga, 6A (Hund’s Rule) o Ionization Energy  Ionization energy: minimum energy (kJ/mol) required to remove an electron from an atom in the gas phase in its ground state  3 qualifiers:  I + X -> X + e - o First ionization energy o Always endothermic + -  Atom in ground state -> Atom + e  ΔU = IE  1 < I2< I3  Overcome the attraction  Farther away from nucleus = smaller IE  IE always positive 3 10/29/15 ­ 11/03/15  Increase across a period, decrease down a group  Owning to an increase in effective nuclear charge across a period, radius decrease, IE increase o Electron Attachment Enthalpy and Electron Affinity  Electron Attachment enthalpy (Δ H): eEAhalpy change occurring when a gaseous atom adds an electron, forming anion  A (g) + e -> A (g) Δ HEA o Exothermic  Electron Affinity (EH): equal in magnitude, opposite sign of ΔU  Negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion  Becomes more negative across a period (increase)  Enthalpy decreases down a group  Noble gases <0  Example: o F (g) + e -> F (g) ΔH = -328 kJ/mol EA = +328 kJ/mol o O (g) + e -> O (g) ΔH = -141 kJ/mol EA = +141 kJ/mol  Trends in Ion Sizes o Positive and negative ions increase in size when descending down the group (still follows trends down a column) o Radius of a cation is always smaller than that of the atom from which it is derived o Anions are always larger than the atoms form which they are derived o Isoelectronic ions have the same number of electrons but a different number of protons  Periodic Trends and Chemical Properties o Properties: atomic and ionic radii, ionization energies, and electron attachment enthalpies o Grouping elements with similar properties o Main group metals generally form cations with an electron configuration equivalent o hat of the preceding noble gas o General statement: nonmetals generally acquire enough electrons to form an anion with the electron configuration of the next noble gas o Ionization energy increases across a period o Carbon: unfavorable as cation and anion o Diagonal relationship:  Li -> Mg: similar charge density  Group 1: fire, more reactive down the column  Group 2: less reactive, filled orbital  3A: redox elements  7A: halogens, hx acids, more reactive up the column  8A: octet rule compliant, least reactive 4


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