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Chapter 8 General Chemistry

by: Cassandra Danhof

Chapter 8 General Chemistry CHEM 1411

Marketplace > Lone Star College-CyFair > Chemistry > CHEM 1411 > Chapter 8 General Chemistry
Cassandra Danhof
Lone Star College-CyFair
GPA 3.21

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These notes go over electronic spin and the Pauli Exclusion Principle, the building up principle and the periodic table, how to write electron configurations using the periodic table, orbital diagr...
General Chemistry I
Prof. Chakranarayan
Class Notes
General Chemistry
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This 9 page Class Notes was uploaded by Cassandra Danhof on Saturday September 24, 2016. The Class Notes belongs to CHEM 1411 at Lone Star College-CyFair taught by Prof. Chakranarayan in Fall 2016. Since its upload, it has received 5 views. For similar materials see General Chemistry I in Chemistry at Lone Star College-CyFair.

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Date Created: 09/24/16
General Chemistry Chapter 8: Electronic Configuration and Periodicity Learning Objectives Important Terms 8.1 Electronic Spin and the Pauli Exclusion Principle Define electron configuration and orbital diagram Electron Configuration:  ● Electron configuration lists the subshells of the atoms A particular distribution  along with the list of electrons consisted in that particularof electrons among the  shell (a statement or shorthand notation) available subshells  ○ Ex. Lithium (atomic number 3) has two Orbital Diagram: A  electrons in its first subshell and one electron in itdiagram to show how  second, or  Li: 1s 2s 1 the orbitals of a  ● Orbital diagrams are a visual representation of the  subshell are occupied  electrons configurations  by electrons Pauli Exclusion  Principle: No two  electrons in an atoms  can have the same four quantum numbers  ○ Ground State: The  State the Pauli exclusion principle lowest­energy  configuration of an  ● No electrons can have the same four quantum  atom numbers (in layman's terms, you cannot place a electron  Excited State: Any  with the same spin in the same orbital)  other allowed  configuration ○ ○ (This can never be a configuration) Apply the Pauli exclusion principle (example 8.1) ● Which of the following orbital diagrams or electron  configurations are possible and which are impossible,  according to Pauli exclusion principle? Explain.  ● A.  ○ Possible orbital diagram ● B. ○ Impossible orbital diagram; there are  three electrons in the 2s orbital, you can only have 2  electrons at a time in an orbital ● C. ○ Impossible orbital diagram; there are  two electrons in the 2p orbital with the same spin, all  electrons in an orbital need to have an opposite spin  from one another  3 1 ● D.  1s 2s ○ Impossible electron configuration;  there are three electrons in the 1s subshell (one  orbital)  ● E.  1s 2s 2 p 7❑ ○ Impossible electron configuration;  there are seven electrons in the 2p subshell (which  can only hold 6 electrons)  2 2 6 2 6 8 2❑ ● F.  1s 2s 2 p 3s 3p 3d 4s ○ Possible. Note that the 3d subshell  can hold as many as 10 electrons  8.2 Building­Up Principle and the Periodic Table Define building up principle  Building­Up (Aufbau)  ● You have to go from the lowest configuration to the  Principle: A scheme  highest configuration when notating an element used to reproduce the  electron configurations  of the ground states of  atoms by successively  filling subshells with  electrons in a specific  order (the building up  order) Noble­Gas Core: An  inner­shell  configuration  corresponding to one of the noble gases  Pseudo­Noble­Gas  Core: The noble gas  core together with 10 (n−1)d electrons  Define noble­gas core, pseudo­noble gas core, and valence  Valence Electron: An  electron electron in an atom  outside the noble­gas  or pseudo gas core  ● The noble gas core is a core corresponding to a  noble gas on the periodic table (the last group to the right on the periodic table) ○ Ex. [He] <­­­ is a noble gas core  2 1 ○ Boron: [He] 2s 2 p ● A pseudo noble gas core is one that is part of the  noble gas core that includes  10 electrons because (n−1)d they are not involved in chemical reactions  ○ Ex. [Ar] 3d 4s 4 p 1 Define main­group element and (d­block and f­block) transition  element ● A main group elements are the elements in the A  groups on the periodic table, or the ones that tower above  the transition elements.  ● The transition elements (the d­block and the f­block)  elements are the elements in the middle of the periodic table 8.3 Writing Electron Configurations Using the Periodic Table Determine the configuration of an atom using the building­up  principle (example 8.2) ● Use the building­up principle to obtain the  configuration for the ground state of the gallium atom (Z=  31)  ● Give the configuration in complete form (don’t  abbreviate for the core). What is the valence­shell  configuration?  ○ Period: ❑ ■ 1st:  1s ■ 2nd:  2s2 p ❑ ■ 3rd:  3s3 p ■ 4th:  4 s3d 4 p ○ Fill the subshell with Gallium’s 31  electrons  ■ 2 2 6 2 6 2 10 1 1s 2s 2 p 3s 3p 4s 3d 4 p ○ Rearrange the subshells by shells to  get the valence shell configuration ■ 1s 2s 2 p 3s 3p 3d6 10 2 1 4s 4 p ● The  valence shell configuration is in purple 8.4 Orbital Diagrams of Atoms; Hund’s Rule State Hund’s rule Hund’s Rule: The  ● You need to put same spin electrons in different  lowest­energy  orbitals to get the lowest­energy of electrons arrangement of  Apply Hund’s rule electrons in a subshell  ● Ex. Florine; Z=9 is obtained by putting  2 2 5 ○ 1s 2s 2 p electrons into separate  orbitals of the subshell  with the same spin  before pairing electrons Paramagnetic  Substance: A  Define paramagnetic substance and diamagnetic substance  substance that is  ● A paramagnetic substance is a substance not paired  weakly attracted by a  with an electron  magnetic field, and this  ○ Ex. Lithium ( 2 1❑ ) 1s 2s attraction is generally  that result of unpaired  electrons  Diamagnetic  Substance: A  substance that is not  ○ attracted by magnetic  ○  (the 2s is unpaired) field or is very slightly  repelled by such a field. ● A diamagnetic substance is a substance where all of  This property generally  their electrons are paired ○ Ex. Beryllium ( 1s 2s 2❑ ) means that the  substance has only  paired electrons  ○ 8.6 Some Periodic Properties  State the periodic law Periodic Law: When  ● The periodic law is about how all of the elements  the elements are  have some sort of pattern to them based on how they  arranged by atomic  arranged horizontally or vertically (period or group) number, their physical  and chemical  properties vary  State the general periodic trends in size of atomic radii  periodically Atomic Radius: The  maximum in the radial  distribution function of  the outer shell of an  atom Atomic (Covalent)  Radius: The value for  that atom in a set of  covalent radii assigned  to atoms in such as  way that the sum of the covalent radii of atoms  A and B predicts the  Define effective nuclear charge  approximate A­B bond  length ● The average nuclear charge felt by an individual  electron in the atom Effective Nuclear  ○ Z =Z−S   Charge: The positive  eff charge that an electron  ○ Z = Number of proton in nucleus experiences from the  ○ S = inner­shell electrons  nucleus, equal to the  ● To determine the effective nuclear charge of an atom for its outer shell, you use the following method:  nuclear charge but  reduced by any  ○ Ex. What is the effective nuclear  shielding or screening  charge felt by an electron in the n =3 shell of sulfur?from any intervening  ○ First, get the electronic configuration  electron distribution of sulfur First Ionization Energy  ■ (First Ionization  1 2 6 2 4 1s 2s 2 p 3s 3 p (refer to period ec)  Potential): Minimum  ● The n=1  energy needed to  2 remove the highest­ shell can hold 2 electrons       ( 1s ) ● The n=2  energy (that is, the  shell can hold 8 electron       ( outmost) electron from  2 6 the neutral atom in the  2s 2 p ) ● The n=3  gaseous state  Electron Affinity: The  shell can hold 18 electron     ( energy required to  3s 3d 3 p 6 ) but since this is  remove an electron  sulfer2 we o4ly fill up 6 of those spots   from the atom’s  ( 3s +3p ) negative ion (in its  ○ Since S in the equation is looking for  ground state)  the inner shell electrons, we subtract the 6 outer shell electrons (the 3 shell) from the inner shell (shell 1  and 2) to get 10 inner shell electrons  ❑ ■ Z eff=16−¿ 10  = Z =6 eff ○ To put shortly, the effective nuclear  charge for an outer shell is always going to be the  number of electrons in that shell ○ Here is a visualization: ● To determine the effective nuclear charge of an atom in an inner shell, you do such ○ Ex. What is the effective nuclear  charge felt by an an electron in the n = 2 shell of  Chlorine? ■ Chlorine has 17 p+ and 17 e­ ■ (remember that the n  =1 shell can hold 2 electrons and the n = 2  shell can hold 8 electrons) ■ 17 ­ 10 = 7 e­ in the 3rd shell (even though it can hold up to 18  electrons, Chlorine only fills up 7 of those  spots) ○ Because we are looking for the  effective nuclear charge in the n =2 shell, we subtract the n = 1 shell because that is the only inner shell we are looking at   ■ Z =Z−S=17−2 eff ❑ ○ Here is the visualization: Determine relative atomic sizes from periodic trends (example 8.5) ● Refer to the periodic table and use the trends noted  for size of atomic radii to arrange the following in order of  increasing atomic radius: Al, C, Si 6 C 13 14 Al Si ● Since the largest atomic radii starts at the bottom left  side, Al would be the largest atomic radii. Since Si and C are in the same group, you would choose the lowest element as  the  ○ 2nd largest radii ○ The atomic radii (in increasing order)  is C < Si < Al State the general periodic trends in ionization energy  Define first ionization energy ● The minimum amount of energy needed to remove  an electron from its outermost (most powerful) shell 2 1 ○ Ex.  Li(1s 2s ) ■ For this, you would  take out the  2s 1 , since it is in the  outermost shell (n=2) −¿ ○ Making:  +¿(1s )+e ¿   ¿ Li ■ (There is a + sign next  to the Li element because we have taken  away one of the electrons from it, making the  element more positive. This is known as a  cation. An anion is the opposite effect (if it  −¿ was  Li¿ )  Determine relative ionization energies from periodic trends  (example 8.6) ● Using the periodic table only, arrange the following  elements in order of increasing ionization energy: Ar, Se, S 16 18 S Ar 34 Se ● Since the largest element is the element from the top right, Ar would be the closest one in this example. Since S  and Se are in the same group, S would be the larger one  because it is at the highest point ○ Increasing ionization energy: Se < S < Ar  Define electron affinity ● Electron affinity is the energy needed to add an  electron from an already negative ion (an anion)  ○ Ex: ❑ ¿ −¿→Cl ● 2 5 ¿ Cl([Ne]3s 3p )+e ¿ ● Note that the regular (ground state) configuration for  2 7 Chloride is  [Ne]3s 3 p . This means that, because of  electron affinity, we are taking an electron out of an already  negative ion. Therefore, first configuration                               ( −¿ 2 5 ¿ ) and the second configuration          Cl([Ne]3s 3p )+e 2 6 [Ne]3 s 3 p ( −¿¿ ) are both the same, just written in different  Cl ¿ ways.  State the broad general trend in electron affinity across any period (All the way up to the 7A group elements, since the noble gases, or  8A group elements, have no electronic affinities because their  valence shell is filled) 


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