CHEM 101 1st Edition Lecture 20Outline of Last Lecture I. Subshell Energy LevelsII. Aufbau PrincipleIII. Electron ConfigurationOutline of Current Lecture I. Ionization EnergyII. Electron Attachment EnthalpyIII. Trends in Ion SizesIV. Chemical Bonds: Ionic BondsV. Chemical Bonds: Covalent BondsVI. Lewis StructureVII. Atom Formal ChargesCurrent Lecture- Ionization Energyo The ionization energy (IE) is the energy required to remove an electron from an atom in the gas phase Ex: Magnesium, first ionization energy Mg(g) + 738 kJ Mg^(+)(g) + e^(-)1s^(2)2s^(2)2p^(6)3s^(2) 1s^(2)2s^(2)2p^(6)3s^(1) Ex: Magnesium, second ionization energy Mg^(+)(g) + 1451kJ Mg^(2+)(g) + e^(-)1s^(2)2s^(2)2p^(6)3s^(1) 1s^(2)2s^(2)2p^(6) Ex: Magnesium, third ionization energy Mg^(2+)(g) + 7732 Mg^(3+)(g) + e^(-)1s^(2)2s^(2)2p^(6) 1s^(2)2s^(2)2p^(5)o Ionization energy increases from right to left and as you go up the periodic table.(Opposite of Atomic radii) - Electron Attachment Enthalpyo The electron attachment enthalpy (ΔEAH) is the enthalpy change observed when a gaseous atom adds an electron to form a gaseous anion. Ex: Atomic ChlorineCl(g) + e^(-) Cl^(-)(g)[Ne} 3s^(2)3p^(5) [Ne]3s^(2)3p^(6)ΔEAH= -349 kJ/mol Ex: Atomic NeonNe(g) + e^(-) Ne^(-)(g)[He]2s^(2)2p^(6) [He]2^(s)2p^(6)3s^(1)ΔEAH= +116 kJ/molo The closely related electron affinity (EA) is equal in magnitude but opposite in sign to the internal energy change of A(g) + e^(-) A(g)^(-)(EA=- ΔU)- Trends in Ion Sizeso The radii of cations are smaller than radii of the corresponding neutral atoms (same Z, smaller number of electrons)o The radii of anions are larger than radii of the corresponding neutral atoms (same Z, larger number of electrons)o Ex: Li^(+) < Li < Li^(-) in sizeChapter 8: Bonding and Molecular Structure- Chemical Bonds: Ionic Bondso A chemical bond between two atoms is an arrangement of valence electrons resulting in a net attractive force between these two atomso Ionic bonds result from the transfer of an electron (or electrons) from one atom to another. o The transfer results in each attaining an octet of noble gas electron configuration. Ex: Na + Cl 2s^(2)2p^(6)=[Ne] 3s^(2)3p^(6)= [Ar]Both become noble gas electrons- Chemical Bonds: Covalent Bondso If one or more electron pairs are shared between two atoms, then the chemical bond is covalent Ex: the electron pair bond between the two atoms of an H2 molecule is represented by a pair of dots or a line (Lewis structure)H-atom H2 molecule:H H: or H-H Lone pairs- the dots/lines not shared by two elements Single bond- the dots/line shared b two elements1A 2A 3A 4A 5A 6A 7A 8Ans^(1) ns^(2) ns^(2)np^(1) ns^(2)np^(2) ns^(2)np^(3) ns^(2)np^(4) ns^(2)np^(5) ns^(2)np^(6) Li Be B C N: :O: :F: :Ne: - Covalent Bonds and Lewis Structureso Electrons are distinguished between core and valence electrons Ex: B 1s^(2)2s^(2)2p^(1)Core = [He], valence = 2s^(2)2p^(1)o How to determine a Lewis structure of molecules and polyatomic ions in up to 5 steps:1. Determination of the central atom: this is usually the atom with the lowest electron affinity2. Determine the total number of valence electrons in the molecule or ion3. We place one pair of electrons between each pair of bonded atoms to form single bonds4. We use the remaining pairs to form more multiple bonds or lone pairs to complete the electron octet of each atom5. If the central atom has less than eight electrons at this point, we change one or more of the lone pairs on the terminal atoms into a bonding pair between the central and terminal atom to form a multiple bond **Hydrogen atoms will never have lone pairs or multiple bonds!o Ex: determine the Lewis structure of ammoniaStep 1: the central atom is usually the atom of lowest affinity for electrons and never hydrogen. Therefore, N is the central atom.Step 2: H=1 and N=5Total = (3x1) + 5 = 8 electrons = 4 pairs of electronsStep 3: H-N-H l H-Step 4: _ H-N-H or H-N-H l l H HStep 5: unnecessary o Ex: N2 triple bondlN(triple bond) Nlo Ex: CO2 double bondO (double bond) C (double bond) Oo Molecules or ions with the same number of valence electrons and similar Lewis structures are isoelectronic:[N (triple bond) O]^(+) N (triple bond)C (triple bond) O [C (triple bond) O]^(-)- Atom Formal Chargeso The formal charge is the charge that would reside on an atom in a molecule or polyatomic ion if we assume that all bonding electrons are shared equally.o It’s calculated based on the Lewis structure of the molecule or ion:Formal Charge = NVE – [Lone Pair Electrons + ½ Bonding
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