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Chem 1030 - Week of February 14

by: Emma Shoupe

Chem 1030 - Week of February 14 Chemistry 1030

Marketplace > Auburn University > Chemistry > Chemistry 1030 > Chem 1030 Week of February 14
Emma Shoupe
GPA 3.7

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About this Document

These notes are from the first two classes following our first exam.
General Chemistry 1
Dr. Livia Streit
Class Notes
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This 4 page Class Notes was uploaded by Emma Shoupe on Thursday February 18, 2016. The Class Notes belongs to Chemistry 1030 at Auburn University taught by Dr. Livia Streit in Spring 2016. Since its upload, it has received 163 views. For similar materials see General Chemistry 1 in Chemistry at Auburn University.


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Date Created: 02/18/16
General Chemistry I – 1030 Dr. Streit February 17, 2016 Ionization Energy  Ionization energy- minimum energy required to remove an electron from an atom in the gas phase  Results in ion – chemical species with a net charge  Cation – positive charge (means loss of electron) st  1 ionization energy – removal of the most loosely held electron  Ionization Energy: the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.A A + +e - (g)  Bigger going right, smaller going down  Exceptions: between Group 2&13, Group 5&6 o z effective= effective nuclear charge increases, ionization energy also increases – from left to right o Within a given shell, electrons with a higher value of quantum number L are higher in energy and easier to remove because they are more stable o Removing a paired electron is easier because of the repulsive forces between 2 electrons in the same orbital o Elements with larger atomic radius have a lower first ionization energy because forces get weaker – the electrons are not as attached to the nucleus o Removing more electrons, means there is a greater attraction to the nucleus which means lnds rdieldith. Therefore, it becomes harder to remove the 2 , 3 , or 4 electron o Principal Energy Level: more energy levels means bigger atoms o Nuclear Charge (# of p in an atom): attraction of electrons to nucleus; increased nuclear charge causes atomic radius to decrease o Shielding Effect: inner electrons increase atomic size by reducing the attractive force on outermost electrons - o Effective Nuclear Charge: force of attraction felt valence e from nucleus; a high ENC means smaller ionic radius (greater attraction to outermost electrons)  Electron affinity- energy released when an atom in the gas phase accepts an electron  Results in an ion  Anion – negative charge (means gain of electron) o Easier to add electron to an s orbital than to a p orbital o Electron affinity increases left to right  Easier to add an electron as the positive charge of a nucleus increases o Within a p subshell, it is easier to add an electron to an empty orbital than to add one to an orbital that already contains an electron o While many first electron affinities are positive, subsequent electron affinities are ALWAYS negative  Considerable energy is required to overcome the repulsive forces between the electron and the negatively charged ion  Electron Affinity: amount of energy required to add an electron to a gaseous atom: − −  Cl(g)+e Cl  Larger going right, smaller going down  Exceptions: between Group 1&2, Group 14&15  Electronegativity: tendency to attract electrons in a covalent bond  Increases going up and right  Metallic character o Metals  Shiny, lustrous, malleable, ductile  Good conductors of heat and electricity  Low ionization energies (form cations) o Nonmetals  Vary in color, not shiny  Brittle  Poor conductors of heat and electricity  High electron affinity (form anions) o Metalloids  Elements with properties intermediate in between metals and nonmetals o Metals: tend to form cations, metal oxides are basic o Nonmetals: tend to form anions, nonmetal oxides are acidic, poor conductors of electricity o Metallic character increases down a group, decreases across a period o Alkali metals (1A)—The most reactive metal family, these must be stored under oil because they react violently with water! They dissolve and create an alkaline, or basic, solution, hence their name. o Alkaline earth metals (2A)—These also are reactive metals, but they don’t explode in water; pastes of these are used in batteries. o Halogens (7A)—Known as the “salt formers,” they are used in modern lighting and always exist as diatomic molecules in their elemental form. o Noble gases (8A)—Known for their extremely slow reactivity, these were once thought to never react; neon, one of the noble gases, is used to make bright signs.  Isoelectric – species with identical electron configurations to the noble gas to the right o Common monatomic ions arranged by their positions in the periodic table  Write the configuration of atom  Add or remove appropriate number of electrons  Na+ = 1s2 2s2 2p6 3s1 = isoelectric to 1s2 2s2 2p6 [Ne] February 18, 2016  High Electron affinity  easy to accept electrons (+)  Low electron affinity  hard to accept electrons (-)  Low ionization energy  easy to form cations  High ionization energy  hard to lose electrons  Ions of d block elements o Formed by removing electrons first from the shell with the highest value of n  For Fe to form Fe 2+, 2 electrons are lost from the 4s subshell not the 3d  Fe can also form Fe 3+, in which case the 3 electron is removed from the 3d subshell  Ionic Radius o The radius of a cation or anion o When an atom loses an electron and becomes a cation, the radius DECREASES due in part to a reduction in electron- electron repulsions in the valence shell and when all of an atoms valence electrons are removed  Comparing Ionic radius with Atomic radius o When an electron gains 1 or more electrons and becomes an anion, its radius INCREASES due to an increase in electron- electron repulsions o Isoelectric series- a series of 2 or more species that have identical electron configurations, but different nuclear charges Chapter 5 – Ionic and Covalent Compounds  Compound- composed of 2 or more elements combined in a specific ratio and held together by covalent bonds o ex- water and salt (sodium chloride) 1:1 ratio  Lewis dot symbols o When atoms form compounds, it is their valence electrons that actually interact o Consists of the element’s symbol with dots o For main group elements, such as Na, the number of dots is the number of electrons that are lost o For non metals in the 2 ndperiod, the number of unpaired dots is the number of bonds the atom can form o


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