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Notes on Lessons 27-29

by: Corymarie Notetaker

Notes on Lessons 27-29 CHMY 141N - 00

Corymarie Notetaker
College Chemistry I
Mark Cracolice (P)

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About this Document

These notes include main points from lessons 27-29 and a summary of each lesson. There are also helpful charts and tables for quick studying.
College Chemistry I
Mark Cracolice (P)
Class Notes
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This 10 page Class Notes was uploaded by Corymarie Notetaker on Wednesday November 18, 2015. The Class Notes belongs to CHMY 141N - 00 at University of Montana taught by Mark Cracolice (P) in Fall 2015. Since its upload, it has received 20 views. For similar materials see College Chemistry I in Chemistry at University of Montana.


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Date Created: 11/18/15
CH MY 141 Notes Lessons 2729 Lesson 27 How do Atoms Combine to Form Ionic Compounds and Molecules Chemical bond describes the forces that hold atoms together to form molecules or polyatomic ions that hold atoms together in metals or that hold oppositely charged ions together to form ionic compounds 0 chemical bonds can break and reform in new combinations when atoms molecules or ions collide I in such a collision the first contact is between the outermost electrons of two particles valence electrons Elements in the same chemical family usually form monatomic ions having the same charge by gaining or losing valence electrons until they become isoelectronic with a noblegas atom and the octet of electrons is complete cation an ion with a positive charge because the atom loses one or more electrons anion an ion with a negative charge because the atom gains one or more electrons Not all elements form monatomic ions naturally hydronium ion H30 the ion of hydrogen does not formally exist itself but rather as a hydrated hydrogen ion ionic compounds compounds made up of ions 0 several combinations of changes appear in ionic compounds but they are always in such numbers that the compound is electrically neutral 0 The solids have a definite geometric structure called a crystal Ions in a crystal are arranged so the potential energy resulting from the attractions and repulsions between them is at a minimum I monatomic and polyatomic ions form crystal structures 0 ionic bonds the strong electrostatic forces that hold the ions in fixed positions I The bonds in an ionic crystal are very strong which is why nearly all ionic compounds are solids at room temperature 0 Solid ionic compounds are poor conductors of electricity because the ions are locked in place in the crystal however liquid ionic compounds and water solutions of ionic compounds are good conductors o dissolving ionic compounds is a physical change Lesson 27 How do Atoms Combine to Form Ionic Compounds and Molecules 0 Molecular compounds compound whose ultimate structural unit is an individual particle known as a molecule 0 the physical and chemical properties of a molecule are different from the properties of the atoms that make up the molecule 0 Lewis proposed that two atoms in a molecule are held together by a covalent bond in which they share one or more pairs of electrons I the bonding pair of electrons spend most of their time between two atomic nuclei they attract both positively charged nuclei and couple the atoms to each other I electron cloud or charge density the greatest probability of location the bonding atoms 0 The electron cloud formed by the two electrons is concentrated in the region between the two nuclei The atomic orbitals of the separated atoms are said to overlap o Orbital overlap refers to the covalent bond formed when the atomic orbitals of individual atoms extend over one another so that the two nuclei in the bond share electrons 0 Lewis diagram Lewis formulas or Lewis structures electron dot symbols that show the bonding arrangement between atoms in a molecule lone electrons unshared electron pairs When two atoms share two bonding electrons the electrons effectively belong to both atoms Atoms bond to achieve an octet of electrons I Octet rule when an atom has a full eight electrons or a completed its octet Single bond sharing of one pair of electrons by two bonded atoms Multiple bonds I double bond two atoms bonded by two pairs of electrons stronger than a single bond I triple bond two atoms bond by three pairs of electrons most abundant substance containing a triple bond N2 stronger than double bond 0 bond strength is measured as the energy required to break a bond I All bonded electrons are valence electrons I Many polyatomic molecules contain multiple bonds Lesson 27 How do Atoms Combine to Form Ionic Compounds and Molecules o A bond in which bonding electrons are shared equally is a nonpolar bond 0 the charge density is centered in the region between the bonded atoms a A bond in which bonding electrons are shared unequally is a polar bond 0 the charge density is shifted towards one atom with the higher electronegativity 0 when the charge density shift is extreme the bonding electrons are effectively transferred to one atom and an ionic bond results a The electronegativity of an element is the ability of an atom of that element in a molecule to attract bonding electron pairs to itself 0 electronegativities tend to be greater at the top of a group in a smaller atom the bonding electrons are closer to the nucleus and therefore are attracted by it more strongly electronegativity increases from left to right across any period matched the increase in nuclear charge among atoms whose bonding electrons are in the same principal energy level electronegativities are highest in the upper right region of the periodic table and lowest in the lower left region can estimate the polarity of a bond by calculating the difference between the electronegativity values for the bonded elements the greater the difference the more polar the bond 0 The more electronegative element toward which the bonding electrons are displaced acts as the negative pole in a polar covalent bond I indicated by an arrow pointing to the negative pole ex HHF I or 5HF539 o memorize F40gtO34gtCl32gtN30Br30gt27gtC26Hgt22 OOOOO Lesson 28 How can the TwoDimensional Arrangement of Atoms in Molecules be Predicted 0 Step 1 Determine the total number of valence electrons I Sum the number of valence electrons contributed by each atom Add or subtract electrons if the species is an ion add electrons for an anion and subtract for a cation I Each main group element atom contributes a number of valence electrons equal to its column number in the periodic table 1A has 1 valence electron 2A has 2 0 Ex CF 4 1x4ve C 4x7ve F 32 valence electrons 0 Step 2 Determine the Central Atoms I the least electronegative atom is usually central I Chemical formulas are often written in order of atoms of increasing electronegativity so the first atom is usually the center of the molecule Or they are written to suggest the sequence of atom connectivity through experience one will be able to recognize the difference When there is more than one carbon atom they tend to bond to one another When a molecule contains a nonmetal atom plus oxygen and hydrogen the oxygen will be bonded to a central nonmetal atom and one hydrogen atom will often bond to each oxygen atom If there is more than one oxygen atom they are typically bonded to the central atom in a symmetrical pattern The formula of the compound will often give guidance to the sequence of atoms in the molecule 0 C25 is less electronegative than F40 and is therefore central 0 Step 3 Draw the first Draft of the Lewis Diagram I Connect the central atom to terminal atoms outside or single bonded atoms with single bonds I Give each atom an octet 4 pairs of electrons Except Hydrogen which will always be joined to a central atom with a single bond 1 pair 0 Step 4 If the first draft of the Lewis diagram has too many electrons delete two unshared pairs on adjacent atoms and add one bonding pair between them Repeat if necessary I This process reduces the electron count in the rough draft by two while preserving the octet of electrons on each affected atom I An ion is surrounded by square brackets and the magnitude and sign of the charge is indicated to the right of the top of the right bracket Lesson 28 How can the TwoDimensional Arrangement of Atoms in Molecules be Predicted 0 Step 5 Perform a formal charge analysis Optimize the formal charges by minimizing the number of atoms with a formal charge and if necessary place negative formal chargess on the outer most electronegative atoms I In a significant number of cases more than one Lewis Diagram can be constructed for a molecule or polyatomic ion In such cases we need a method to determine which diagram is most likely to match the experimentallydetermined actual structure 0 a relatively accurate methodology for doing so involves analysis of the formal charge of each atom in a Lewis diagram ls based on two quantities 0 First for an atom to be neutral it needs to have the same number of valence electrons as it would if it were not part of a molecule 0 Second to determine the formal charge of an atom is the number of electrons allocated to it by the formal charge methodology This is accomplished by assigning all unshared electrons to the atom plus exactly half of the bonding electrons I formal charge the charge that an atom would experience if bonding electrons were shared equally I formal charge Electrons on separate atomelectrons on atom in molecule I Guidelines to choose the best Lewis diagram 0 1 the sum of the formal charges of the atoms must equal the charge on the species 0 2 the best Lewis diagram has the least number of atoms with formal charges 0 3 the best Lewis diagram places negative formal charges on the most electronegative atoms o 4 if more than one atom has a formal charge opposite charges tend to occur on neighboring atoms Lesson 28 How can the TwoDimensional Arrangement of Atoms in Molecules be Predicted Summary of how to draw a Lewis Diagram Determine the total number of valence electrons Determine the central atoms Draw the first rough draft of the Lewis diagram If the first draft has too many electrons delete two unshared pairs on adjacent atoms and add one bonding pair between them Perform a formal charge analysis Optimize the formal charges by minimizing the number of atoms with a formal charge and if necessary place a negative formal charges on the most electronegative atoms WEAPON Lesson How are Covalent Bonds in Molecules Optimally Arranged and What are Their Lengths and Strengths 0 when there is one single or double bond in a species it is distributed through its resonance structures and is mathematically supported by an electrostatic potential map 0 Electrostatic potential map shows the results of calculations of charge distributions within molecules and polyatomic ions 0 resonance structures the series of Lewis diagrams that illustrate the different possible distributions of valence electrons I the actual structure is an average of the resonance structures 0 resonance hybrid a single unchanging average of what is illustrated by the Lewis diagrams I the general trend is that the larger the number of resonance forms of a species the greater the stability A higher number of resonance forms indicates that the electrons are more spread out and this leads to greater stability 0 Violations of the octet rule 0 we consider three important categories of exceptions to the octet rule I oddelectron molecules odd electron is not localized on the atoms but rather spread over the entire molecule Oddelectron species must have a single electron in place of the typical electron pair at someplace within the Lewis diagram 0 the odd electron will be on the least electronegative atom o For molecules with an odd number of valence electrons 0 species with an odd number of electrons are called free radicals The unpaired electron makes these species highly reactive I molecules with atoms that have more than an octet atoms with n3 and higher have 35 3p and 3d orbitals that can be half filled and bond to other atoms making up to nine bonds potentially possible I molecules with atoms that have less than an octet some compounds such as boron compounds have experimental evidence that indicates they are more stable with just three pairs of electrons surrounding the central atom 0 an analysis of formal charge indicates that molecules are likely to be stable because all atoms have zero formal charge Lesson How are Covalent Bonds in Molecules Optimally Arranged and What are Their Lengths and Strengths Procedure Drawing Lewis Diagrams Step 1 Determine the total number of valence electrons Step 2 Determine the central atoms Step 3 Do the first draft of the Lewis Diagram Step 4 If the first draft has too many electrons delete two unshared pairs on adjacent atoms and add one bonding pair between them Repeat as necessary Step 5 If the first draft has too few electrons add an additional unshared pair of electrons to the central atom Repeat if necessary Step 6 Perform a formal charge analysis Optimize the formal charges by minimizing the number of atoms with a formal charge and if necessary place negative formal charges on the most electronegative atoms Step 7 Check to be sure that C N O and F atoms have an octet H atoms should share two electrons Be sure that the valence electron count in the Lewis diagram matches the total available from Step 1 Step 8 Draw resonance structures if necessary Lesson How are Covalent Bonds in Molecules Optimally Arranged and What are Their Lengths and Strengths 0 Bond order refers to the number of electron pairs shared in a covalent bond between two atoms 0 for species without resonance structures the bond order between any two specified atoms is simply the number of bonding electron pairs between those atoms 0 when a species has resonance structures the bond order is the average number of bonds shared between two atoms 0 bond order Number of shared electron pairs between atoms X and Y Number of links between atoms X and Y 0 Bond length the average distance between the nuclei of two atoms in a molecule 0 the larger an atom the greater the length of the bond from it to a constant second atom I The hydrogen halides show the trend of increasing bond length with increasing Group 7A atom size I The second period nonmetals illustrate the trend of increasing bond length with increasing second period atom size 0 bond length decreases with increasing bond order I single bond gtdouble bond gt triple bond 0 Bond enthalpy is the energy required to break a specified bond in one mole of a compound in the gas phase 0 tables of bond enthalpies report the average value from a range of experiments Therefore bond enthalpies can only be used to estimate AH for a reaction 0 instead of using the enthalpy of formation of the compounds in the reaction we can use the enthalpies of the bonds themselves I AH ZEnthalpies of bonds brokenpositive values ZEnthalpies of bonds formednegative values 0 to break bonds requires energy forming bonds releases energy


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