Dr. Streit Week 5 Notes
Dr. Streit Week 5 Notes CHEM 1030 - 003
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This 3 page Class Notes was uploaded by Rachel Ferrell on Friday February 19, 2016. The Class Notes belongs to CHEM 1030 - 003 at Auburn University taught by John D Gorden in Fall 2015. Since its upload, it has received 29 views. For similar materials see Fundamentals Chemistry I in Chemistry at Auburn University.
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Date Created: 02/19/16
Rachel Ferrell CHEM 1030 2/15/16 Chapter 4 cont.: Ionization Energy: • =minimum energy required to remove an electron from an atom in the gas phase • result is an ion= a chemical species with a net charge o cation= positive charge o anion= negative charge • Na(g)→Na (g) + e ‐ o Ionization energy of Na= 495.8 KJ/mol o 1 IE of Na→corresponds to the removal of the most loosely held electron in a valence shell • in general→ As Zeff increases, IE also increases o therefore, ionization energy increases from left to right across a period o Exceptions § IE decreases from Group 2A-‐3A • Because within a given shell, higher l value are higher in energy, thus easier to remove • Group 2A= 2s • Group 3A= 2p →easier to remove; lower IE § IE decreases from Group 5A-‐6A • Because removing a paired electron is easier because of repulsive forces between two electrons in the same orbital • Group 5A= np 4 • Group 6A= np →easier to remove; less stable; lower IE • Can also remove additional electrons →IE 2, , 3 etc. o Takes more energy to remove core electron than a valence electron § Because they are closer to nucleus Electron Affinity: • Energy released when an atom in a gas phase accepts an electron o Cl(g)+ e → Cl (g) o Result is always a anion • EA increases from left→right across a period as Zeff increases o Because its easier to add an electron as the (+) charge of the nucleus increases • Exceptions: o EA decreases from Group 1A to Group 2A § Because its easier to add an electron to an s orbital than to add to a p orbital with the same n o EA decreases from Group 4A to Group 5A § Because with a p shell, its easier to add an electron to an empty orbital than to add one to an orbital already with an electron • More than one electron may be added to an atom o EA 2oes down→always negative o Because more energy is required to overcome repulsive force between an electron and a (-‐) charged atom Metallic Character: • Metals tend to o Be shiny, malleable o good conductors of electricity o have low IE→commonly form cations (hard to accept electrons, easy to lose electrons) • Nonmetals o Vary in color, not shiny o Brittle o High EA→commonly form anions (easy to accept electrons, hard to lose electrons) • Metalloids o Properties intermediate between metals and nonmetals • Periodic trends→explained by Columbs law ???????? ???? ???????? o F???? ????^???? o Q1/2= charged objects o D=distance o Explains repulsion and attraction forces o Explains atomic radius, ionization energy, and electron affinity Ions of Main Group Elements: • Isoelectrons= species with the same electron configuration to the noble gases to the right o due to adding or removing electrons o exception: Mercury=polyatomic ion Electron Configuration of Ions: • cation→remove 1 electron • anion→add 1 electron • Steps: o 1) Write the electron config for the atom o 2) Add of remove the appropriate number of electrons Ions of D-‐Block Elements: • ions formed by removing electron first from the shell with the highest value of n o ex. Fe→Fe 2+ § Fe: [Ar] 4s 3d 2+ 6 § Fe : [Ar] 3d § Fe : [Ar] 3d § Therefore, remove from 4s before 3d because it has a higher n value Ionic Radius: • =radius of a cation/anion • when an atom loses an electron to become a cation, radius decreases because o 1) reduction in electron-‐electron replulsions (shielding) in the valence shell o 2) significant decrease in radius occurs when all of atoms valence electrons are removed o cation radius<atomic radius • when an atom gains an electron, radius increases due to increased electron-‐electron repulsions(electrons trying to spread out) o anion radius>atomic radius Isoelectronic Series: • a series of 2 or more species that have identical electron configurations but different nuclear charges Chapter 5: Compounds: • =substance composed of 2 or more elements combined in a specific ratio and held together by chemical bonds o ex. Water and salt (NaCl) Lewis Dot Structures: • atoms combine→more stable electron configuration • maximum stability→isoelectronic with Noble Gases • when compounds form it is the valence electrons that interact with each other • Lewis Dot Symbols= element symbols with dots o each dot=valence electron • for main group metals such as Na→dots are the number of electrons lost in Na • for nonmetals→unpaired dots= number of bonds an atom can form • Ion dot structures o Written in the same way except with brackets and charges 2-‐ o Ex. [O]
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