CHEM 101 Chapter 7
CHEM 101 Chapter 7 Chem 101
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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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