Chemistry Notes Week 3
Chemistry Notes Week 3 Chem 1010
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This 6 page Class Notes was uploaded by Courtney Burke on Wednesday September 30, 2015. The Class Notes belongs to Chem 1010 at University of Denver taught by Teresa Cowger in Fall 2015. Since its upload, it has received 80 views. For similar materials see General Chemistry 1010 in Science at University of Denver.
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Date Created: 09/30/15
CHEMISTRY 1010 NOTES Week 3 September 28 October 2 2015 Courtney Burke 7072664834 SECTION 27 Compounds Introduction to Bonding F i The overwhelming majority of elements occur in compounds combined with other elements The noble gases helium neon argon krypton xenon and radon occur in the air as separate atoms Oxygen nitrogen and sulfur also occur in their most common elemental form as the molecules 02 N2 and Sg Elements combine in two general ways 0 Transferring electrons from one element to another to form ionic compounds 0 Sharing electrons between atoms of different elements to form covalent bonds Chemical bonds The forces that hold the atoms together in a compound Ions Charged particles that form when an atom gains or loses one or more electrons The simplest type of ionic compound is a binary ionic compound one composed of two elements and typically forms when a metal reacts with a nonmetal 0 Each metal atom loses one or more electrons and becomes a cation a positively charged ion 0 Each nonmetal atom gains one or more electrons lost by the metal atom and becomes an anion a negatively charged ion A cation or anion derived from a single atom is called a monatomic ion All binary ionic compounds are solid arrays of oppositely charged ions Coulomb s Law lii39ill l E I39quot Where FE is the electric feree he is the Coulomb eeneteiit is BEETEKIUQN in ict q ie the eherge en ehjeeti and rie the dieteiiee between the ehergee O Ions with higher charges attract or repel each other more strongly than ions with lower charges 0 Smaller ions attract or repel each other more strongly than larger ions because the charges are closer to each other Ionic compounds are neutral because they contain equal numbers of positive and negative charges Nonmetal atoms gain electrons to form ions with the same number of electrons as in an atom of the nearest noble gas Metals lose electrons elements in Group 1A1 lose one electron elements in Group 2A2 lose two electrons and elements in Group 3A3 lose three electrons Nonmetals gain electrons elements in Group 7A17 gain one electron elements in Group 6A16 gain two electrons and elements in Group 5A15 gain three electrons Atoms of different elements share electrons to form the molecules of a covalents compound Most covalent substances consist of molecules and there are no molecules in an ionic compound The nature of the particles attracting each other in covalent and in ionic substances is fundamentally different Many ionic compounds contain polyatomic ions which consist of two or more atoms bonded covalently and have a net positive or negative charge SECTION 91 Atomic Properties and Chemical Bonds Bonding lowers the potential energy between positive and negative particles There is in general a gradation from more metallic elements to more nonmetallic elements across a period and up a group Metal with nonmetal electron transfer and ionic bonding We observe electron transfer and ionic bonding between atoms with large differences in their tendencies to lose or gain electrons Nonmetal with nonmetal electron sharing and covalent bonding When two atoms differ little or not at all in their tendencies to lose or gain electrons we observe electron sharing and covalent bonding which occurs most commonly between nonmetals The shared electron pair is typically localized between the two atoms linking them in a covalent bond of a particular length and strength Metal with metal electron pooling and metallic bonding In the simplest model of metallic bonding the enormous number of atoms in a sample of a metal pool their valance electrons into a sea of electrons that ow between and around each metalion core thereby attracting and holding them together Electrons in metallic bonding are delocalized moving freely throughout the entire piece of metal Lewis Symbols and the Octet Rule 0 Lewis electrondot symbol the element symbol represents the nucleus and inner electrons and dots around the symbol represent the valence electrons 0 For a metal the total number of dots is the number of electrons an atom loses to for a cation 0 For a nonmetal the number of unpaired dots equals either the number of electrons an atom gains to form an anion or the number it shares to form covalent bonds 0 Octet rule When atoms bond they lose gain or share electrons to attain a filled outer level of eight electrons SECTION 92 The Ionic Bonding Molecule 0 The transfer of electrons from metal atoms to nonmetal atoms to form ions that attract each other and form a solid compound 0 In ionic bonding the total number of electrons lost by the metal atoms equals the total number of electrons gained by the nonmetal atoms Whv Ionic Compounds Form The Importance of Lattice Energv 0 Energy is actually absorbed during electron transfer 0 The electron transfer process 0 The first ionization energy is the energy absorbed when 1 mol of gaseous atoms loses 1 mol of valence electrons O The first electron affinity EAl is the energy released when 1 mol of gaseous atoms gains 1 mol of electrons 0 If the overall reaction releases energy there must be some step that is exothermic enough to outweigh the endothermic steps This step involves the strong attraction between pairs of oppositely charged ions 0 Even more energy is released when the separate gaseous ions coalesce into a crystalline solid because each ion attracts several oppositely charged ions 0 The lattice energy AH39lat ce is the enthalpy change that accompanies the reverse of the previous equation 1 mol of ionic solid separating into gaseous ions 0 In a BornHaber cycle a series of steps from elements to ionic solid for which all the enthalpies are known except the lattice energy 0 The BomHaber cycle shows that the energy required for elements to form ions is supplied by the attraction among the ions in the solid And the take home lesson is that ionic solids exist only because the lattice energy far exceeds the total energy needed to form the ions 0 Electrostatic energy decreases between cations and anions and thus lattice energy should decrease as well How the Model Explains the Properties of Ionic Compounds 0 A typical ionic compound is hard rigid and brittle 0 Ionic compounds typically do not conduct electricity in the solid state but do conduct when melted or dissolved We expect ionic compounds to have high melting points and much higher boiling points Ion pairs gaseous ionic molecules rather than individual ions Ionic compounds are solid arrays of ions and no separate molecules exist SECTION 93 The Formation of a Covalent Bond Sharing electrons is the main way that atoms interact A covalent bond arises from the balance between the nuclei attracting the electrons and electrons and nuclei repelling each other Formation of a covalent bond always results in greater electron density between the nuclei The shared bonding pair is represented by a pair of dots or a line An outerlevel electron bonding pair that is not involved in bonding is called a lone apir or unshared pair Prooerties of a Covalent Bond Order Energy and Length The bond order is the number of electron pairs being shared by a given pair of atoms A single bond is the most common bond and consists of one bonding pair of electrons a single bond has a bond order of 1 Multiple bonds usually involve C O andor N atoms A double bond consists of two bonding electron pairs four electrons shared between two atoms so the bond order is 2 A triple bond consists of three shared pairs two atoms share six electrons so the bond order is 3 The strength of a covalent bond depends on the magnitude of the attraction between the nuclei and shared electrons The bond energy called bond enthalpy is the energy needed to overcome this attraction and is defined as the standard enthalpy change for breaking the bond in 1 mol of gaseous molecules The same quantity of energy absorbed to break the bond is released when the bond forms Bond formation is an exothermic process so the sign of its enthalpy change is always negative Ag Bg I A Bg AH39bondforming BEA B always lt 0 Stronger bonds have a larger BE because they are lower in energy have a deeper energy well Weaker bonds have a smaller BE because they are higher in energy have a shallower energy well Bond length is the distance between the nuclei of two bonded atoms The order energy and length of a covalent bond are interrelated 0 For a given pair of atoms a higher bond order results in a smaller bond length and a higher bond energy Thus a shorter bond is a stronger bond How the Model Explains the Properties of Covalent Substances 0 Strong bonding forces hold the atoms together within the molecule and weak intermolecular forces act between separate molecules in the sample 0 Network covalent solids are held together by covalent bonds between atoms throughout the sample and their properties do re ect the strength of covalent bonds 0 Most covalent bonds are poor electrical conductors because their electrons are localized as either shared or unshared pairs and no ions are present SECTION 101 Depicting Molecules and Ions with Lewis Structures 0 Lewis structure or Lewis Formula shows symbols for the atoms the bonding electron pairs as lines and the lone electron pairs that fill each atom s outer level valence shell as pairs of dots Applying the Octet Rule to Write Lewis Structures 0 The four steps for writing Lewis structures for species with only single bonds 0 Place the atoms relative to each other 0 Determine the total number of valence electrons 0 Draw a single bond from each surrounding atom to the central atom and subtract from the total for each bond to find the number of e remaining 0 Distribute the remaining electrons in pairs so that each atom ends up with 8e 0 Hydrogen atoms form one bond carbon atoms form four bonds nitrogen atoms form three bonds oxygen atoms for two bonds and surrounding halogens form one bond uorine is always a surrounding atom 0 Steps for writing Lewis Structures for species with multiple bonds 0 Same steps for single bonds 0 And then if a central atom does not end up with an octet form one or more multiple bonds Resonance Delocalized ElectronPair Bonding 0 Resonance Structures Forms Have the same relative placement of atoms but different locations of bonding and lone electron pairs 0 An average of the resonance structures is called a resonance hybrid 0 Electronpair delocalization In a single double or triple bond each electron pair is localized between the bonded atoms In a resonance hybrid two of the electron pairs one bonding and one lone pair are delocalized their density is spread over a few adjacent atoms Formal Charge Selecting the More Important Resonance Structure Formal charge The charge an atom would have if the bonding electrons were shared equally Formal charges must sum to the actual charge on the species Smaller formal charges positive or negative are preferable to larger ones The same nonzero formal charges on adjacent atoms are not preferred A more negative formal charge should reside on a more electronegative atom For a formal charge bonding electrons are shared equally by the atoms so each atom has half of them Formal charge valence e lone pair e 12 bonding e For an oxidation number bonding electrons are transferred completely to the more electronegative atom as if the bonding were pure ionic Oxidation number valence e lone pair e bonding e Electron deficient They have fewer than eight electrons around the central atom Most molecules have a central atom from an oddnumbered group such as N Group 5A15 or Cl Group 7A17 These are called free radicals species that contain a lone unpaired electron which makes them paramagnetic and extremely reactive Many molecules and ions have more than eight valence electrons around the central atom That atom expands its valence shell to form more bonds which releases energy Expanded valence Shells occur only with nonmetals from period 3 or higher because they have d orbitals available