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Chemistry 1: Test 2 Study Guide

by: Madison Greer

Chemistry 1: Test 2 Study Guide ch 1213

Marketplace > Mississippi State University > Chemistry > ch 1213 > Chemistry 1 Test 2 Study Guide
Madison Greer

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About this Document

These notes cover everything in chapters 5, 6, and 7.
Chemistry 1
Erin Dornshuld
Study Guide
Chemistry, 5, 6, 7, Test 2, dornshuld, chem one, chem 1, Study Guide
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This 10 page Study Guide was uploaded by Madison Greer on Thursday February 25, 2016. The Study Guide belongs to ch 1213 at Mississippi State University taught by Erin Dornshuld in Summer 2015. Since its upload, it has received 337 views. For similar materials see Chemistry 1 in Chemistry at Mississippi State University.


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Date Created: 02/25/16
• ionization energy (IE)- the energy required to remove an electron from a neutral atom • electron affinity (EA)- the energy change when a neutral atom attracts an electron to become a negative ion • Ionic compounds tend to form between metals and nonmetals when electrons are transferred from an element with a low ionization energy (the metal) to one with a high electron affinity (the nonmetal) • compound- substance composed of two or more elements combined in a specific ratio and held together by chemical bonds • Maximum stability results when a chemical species is isoelectronic with a noble gas. • Example- Na+= isoelectronic with Ne; Cl-= isoelectronic with Ar • Lewis dot symbol- consists of an element’s symbol with dots. Each dot represents a valence electron • cation- lose electron (positive) • anion- gain electron (negative) • ionic bonding- refers to the electrostatic attraction that holds oppositely charged ions together in an ionic compound • chemical formula- denotes the constituent elements and the ratio in which they combine • lattice- a 3D array of oppositely charged ions • lattice energy- amount of energy required to convert a mole of ionic solid to its constituent ions in the gas phase • monatomic cation- named by adding the word ion to the name of the element • monatomic anion- named by changing the ending of the element’s name to -ide • In cases where a metal cation may have more than one possible charge, the charge is indicated in the name of the ion with a Roman numeral in parenthesis. • Lewis theory of bonding- where electrons are shared, not transferred; covalent bonds • molecule- a combination of at least two atoms in a specific arrangement held together by chemical forces • Law of definite proportions- different samples of a given compound always contain the same elements in the same mass ratio • Law of multiple proportions- if two elements can combine with each other to form two or more different compounds, the ratio of masses of one element that combines with the fixed mass of the other element can be expressed in small whole numbers (works for simple compounds) • diatomic molecules- molecules that contain only two atoms • homonuclear- both atoms are the same; H , F , 2i ,2etc2 • heteronuclear- both atoms are different; HF, NaCl, etc… • polyatomic molecules- contain more than two atoms • chemical formula- denotes the composition of the substance • molecular formula- shows the exact number of atoms of each element in a molecule • structural formula- shows not only the elemental composition but also the general arrangements (and sometimes connectivity) • Some elements have two or more distinct forms known as allotropes. Example- oxygen (O 2 and ozone (O ) 3re allotropes of oxygen. empirical formula- written by reducing the subscripts to the smallest possible whole numbers • • binary molecular compounds- compounds with two elements • There are metals that will form covalent bonds rather than ionic. MEMORIZE B2H6 diborane SiH4 silane NH 3 ammonia MEMORIZE PH 3 phosphane H 2 water H 2 hydrogen sulfide • acid- a substance the produces hydrogen ions when dissolved in water • inorganic compounds- generally defined as compounds that do not contain carbon • organic molecules- molecules that contain carbon and hydrogen • hydrocarbons- the simplest organic compounds that consist only of carbon and hydrogen MEMORIZE CH 4 methane C2H 6 ethane C3H 8 propane C 4 10 butane C 5 12 pentane C 6 14 hexane C 7 16 heptane C 8 18 octane C 9 20 nonane C10 22 decane • For alkanes, the names end in -ane. • functional group- groups that determine many chemical properties of a compound because that is typically where a chemical reaction occurs • polyatomic ions- a charged chemical species compound of either two or more atoms that are covalently bonded or a metal complex that can be considered to be acting as a single unit • MEMORIZE *look at notes for list of polyatomic ions* • oxoanions- polyatomic ions that contain one or more oxygen atoms and one atom of another element; ex- chlorate, nitrate, sulfate • oxoacids- produce hydrogen ions and the corresponding oxyanions when dissolved in water; can be monoprotic or polyprotic • Oxyanions whose name ends in -ate, can be named as follows: • An acid based on an -ate ion is called -ic acid; ex- HClO i3 called chloric acid • An acid based on an -ite ion is called -ous acid; ex- HClO is2called chlorous acid • Prefixes in oxyanion names are retained in the names of the corresponding oxoacids; ex- HClO an4 HClO are called perchloric acid and hypochlorous acid • hydrate- a compound that has a specific number of water molecules within its solid structure • anhydrous- the compound no longer has water molecules associated with it; driven off all the water from the hydrate MEMORIZE CO 2 dry ice NaCl salt N2O nitrous oxide, laughing gas CaCO 3 marble, chalk, limestone NaHCO 3 baking soda MgSO 4 *H 2 Epsom salt Mg(OH) 2 milk of magnesia • molecular mass (molecular weight)- the mass in atomic mass units (amu) of an individual molecule. Because the atomic masses on the periodic table are averaged masses, the molecular mass is also and averaged value. • percent composition by mass- a list of percent-by-mass of each element in a compound • molecule= covalent bonds, molecular mass • ionic compound= ionic bonds, formula mass • molar mass- the mass in grams of one mole of a substance • *know how to do problems like Example 5.16 in notes* • octet rule- atoms will lose, gain, or share (valence) electrons to achieve a noble gas configuration • Lewis structure- a representation of covalent bonding in molecules • single bond- 2 electrons, one electron pair • double bond- 4 electrons, 2 electron pairs • triple bond- 6 electrons, 3 electron pairs • bond length- the distance between two nuclei of two covalently bonded atoms in a molecule bond length bond strength • shortest longest weakest strongest triple < double < single single < double < triple • pure covalent bond- neutral atoms held together by equally shared electrons • polar covalent bond- partially held together by unequally shared electrons • ionic bond- oppositely charged ions held together by electrostatic attraction • electronegativity- the ability of an atom to draw shared electrons toward itself • Six Steps for Drawing Lewis Structures for Molecules and Polyatomic Ions 1. Draw the skeletal structure of the compound. 2. Count the total number of valence electrons present; add electrons for negative charges and subtract electrons for positive charges. 3. For each bond in the skeletal structure, subtract two electrons from the total valence electrons. 4. Use the remaining electrons to complete octets of the terminal atoms by placing pairs of electrons on each atom. 5. Place any remaining electrons in pairs on the central atom. 6. If the central atom has fewer than 8 electrons, move one or more pairs from the terminal atoms to form multiple bonds between the central atom and the terminal atoms. • formal charge- used to determine the most plausible Lewis structure • formal charge “equation”- (# of valence electrons for that atom)-(# of bonds to that atom)-(# of dots around that atom)= formal charge • resonance structure- one of two or more equally valid Lewis structures for a single species that cannot be represented accurately with a single Lewis structure • Resonance structures are NOT rapidly changing structure; they are always equal. • expanded octet- the central atom has more than 8 electrons • free radicals- molecules with an odd number of electrons • coordinate covalent or dative bond- a bond created from a pair of electrons donated from one atom • Lewis acid- can accept a pair of electrons • Lewis base- can donate a pair of electrons • VSEPR (valence-shell electron pair repulsion) theory- assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between electrons in the valence shell of that atom • 180 degrees between 2 electron domains; linear • 120 degrees between 3 electron domains; trigonal planar • 109.5 degrees between 4 electron domains; tetrahedral • trigonal pyramidal- 5 electron domains • octahedral- 6 electron domains electron domain geometry- the arrangement of electron domains around the central atom • • molecular geometry- the arrangement of bonded atoms (ignore all lone electron pairs) • bond angle- the angle between two adjacent A-B bonds KNOW THIS!!! total number number of type of electron- molecular example of electron lone pairs molecule domain geometry domains geometry 3 1 AB trigonal planar bent SO 2 2 4 1 AB 3 tetrahedral trigonal NH 3 pyramidal 4 2 AB 2 tetrahedral bent H2O 5 1 AB 4 trigonal seesaw- SF4 bipyramidal shaped 5 2 AB 3 trigonal T-shaped ClF 3 bipyramidal 5 3 AB 2 trigonal linear IF2 bipyramidal 6 1 AB 5 octahedral square BrF 5 pyramidal 6 2 AB 4 octahedral square planar XeF 4 • Loan pairs take up more space than bonds. • VSEPR (valence-shell- electron pair repulsion)- assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between electrons in the valence shell of that atom • electron domains • lone electron pairs • single bonds • double bonds • triple bonds • 2 electron domains= linear= 180 degrees • 3 electron domains= trigonal planar= 120 degrees • 4 electron domains= tetrahedral= 109.5 degrees • 5 electron domains= trigonal bipyramidal • 6 electron domains= octahedral • electron domain geometry- the arrangement of electron domains around the central atom • molecular geometry- the arrangement of bonded atoms (ignore all lone electron pairs) • bonds angle- the angle between two adjacent A-B bonds • Electron Domain and Molecular Geometries and Molecules with Lone Pairs on the Central Atom total number number of type of electron- molecular of electron domain example domains loan pairs molecule geometry geometry 3 1 AB trigonal planar bent SO 2 2 4 1 AB 3 tetrahedral trigonal NH 3 pyramidal 4 2 AB 2 tetrahedral bent H2O trigonal seesaw- 5 1 AB 4 bipyramidal shaped SF 4 trigonal 5 2 AB 3 T-shaped ClF3 bipyramidal 5 3 AB 2 trigonal linear IF2- bipyramidal square 6 1 AB 5 octahedral pyramidal BrF5 6 2 AB 4 octahedral square planar XeF 4 • Strong Weak Ion-Ion Ion-Dipole Hydrogen Bond Dipole-Dipole Dipole-Induced Dipole Dispersion • polar bond- atoms bonded together are different • non polar bond- atoms held together are the same • structural isomers- molecules with the same molecular formula but bonded in either a different way or arranged in a different way • trans- different sides • cis- same sides; boiling point is higher • intermolecular forces- attractive forces between neighboring molecules; collectively known as van der Waals forces • dipole-dipole interactions- attractive forces that act between polar molecules • hydrogen bond- an important type of dipole-dipole interaction; occurs between molecules that contain H bonded to a small, highly electronegative atom such as N, O, or F • dispersion forces- also known as London dispersion forces; attractive forces that arise from temporary dipoles • ion-dipole interactions- the Coulombic attraction between ions (positive or negative) and polar molecules • induced dipole- a dipole is induced by a neighboring charge • valence bond theory- atoms share electrons when an atomic orbital in one atom overlaps with an atomic orbital of another • Energy always wants to flow downhill. • p-orbitals=90 degrees • hybridization- the mixing of atomic orbitals Number of Electron Domains and Hybrid Orbitals on a Central Atom number of electron domains on 2 3 4 5 6 central atom hybrid orbitals sp sp 2 sp3 sp d sp d2 • sigma (σ) bond- all single bonds are sigma bonds; double and triple bonds each contain a single sigma bond • pi (π) bond- single bonds contain no pi bonds; double bonds contain a single pi bond; a triple bond contains two pi bonds • molecular orbital theory- another bonding theory which is based on the premise that atomic orbitals combine to form molecular orbitals that are a ‘property’ of the entire molecule • bonding molecular orbital- constructive combination of the two 1 s orbitals gives rise to a molecular orbital that lies along the intermolecular axis • anti bonding molecular orbital- destructive combination of the two 1 s orbitals gives rise to a molecular orbital that lies along the intermolecular axis, but does not lie between the two nuclei • bond order- the higher the bond order, the more stable the molecule is


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