Chemical Bonds, Nomenclature, Lewis Structure and Molecular Shapes Chapter 5 Chemical Bonds, Nomenclature, Lewis Structure and Molecular Shapes

Homework & Quizzes – Chapter 5 Text Homework (not turned in): pages 147 – 151. Problems: 1, 6 – 8, 17, 27, 35, 39, 41 – 66, 68, 69, 72, 86, 88(b&c), 107, 109, 110, 112, 113. Quiz: Do the graded quiz in Blackboard.

I. Chemical Bonds A. Introduction (summary of chapter) Atoms can combine to produce new larger units called molecules or compounds. Each molecule has a unique name (two rules to learn). Molecules held together by chemical bonds (two types). Bonds result from either transfer of valence electrons (Ionic Bonds) or from sharing of valence electrons – Covalent Bonds. Valence electrons rearrange to mimic closest Group VIIIA (18) structure. Molecules resulting from covalent bonding will have predictable shapes.

I. Chemical Bonds B. Ionic Bonds - Metals (except H) loose electrons, form cations; Nonmetals gain electrons to form anions. Both strive for e- configuration of nearest inert gas. - The resulting opposite ions attract in a ratio which produces a neutral unit. Reduce formula to simplest ratio. Ionic Bond Definition: bond formed by electrostatic attraction between anions (-) and cations (+). Write formula with + element first; do not show charges; Final compound is neutral. Generality: any compound formed from metallic and nonmetallic elements is ionic.

I. Chemical Bonds B. Ionic Bonds Know: Metals combine with nonmetals & form ionic bonds by losing or gaining electrons to mimic closest Inert Gas (VIIIA). IA - Na, K, Li, etc become +1 ions: Na+ IIA - Ca, Mg, etc become +2 ions: Ca+2 IIIA - Al, Ga become +3 ions: Al+3 VA - N, P become -3 ions: N-3 VIA - O, S become -2 ions: O-2 VIIA - F, Cl, Br, I become -1 ions: F-1 Opposite ions attract in a ratio so that the product is neutral. Inert Gas e- Configurations

I. Chemical Bonds B. Ionic Bonds Example Do not show charges in final formula. NaCl NOT Na+Cl-

I. Chemical Bonds B. Ionic Bonds Example

I. Chemical Bonds B. Ionic Bonds - Examples Na+ + Na+ + O-2 Na2O Ca+2 + F- + F- CaF2 Mg+2 + S-2 MgS Al+3 + Al+3 + O-2 + O-2 + O-2 Al2O3 Give the formulas for the following: Na & Br Ca & O Ba & I Li & O Al & F Mg & N Many transition metals form ionic bonds & can have several charges such as Fe+2 = Iron (II); Fe+3 = Iron (III); Cu+2 = Copper (II); Cu+1 = Copper (I)

I. Chemical Bonds C. Electron Dot (Lewis) Structures - A Lewis electron dot structure is a symbol in which the valence electrons are shown as dots. - Examples: Na. Mg: Na+ Ca2+ H:1- (Called Hydride) :C: :Si: - How many valence electrons (dots) would N3- O2- F- or Ne have? What about Mg+2? 8 8 8 8 0

II. Covalent Bonds A. Introduction EN = electronegativity - Definition of a covalent bond: A bond formed by the sharing of two electrons. When two atoms of similar EN combine, neither has the “pull” to take electrons away & a sharing of electrons results. This occurs when NONMETALS, including H, combine with NONMETALS. Example: H. + H. ---) H—H = H2 - The atoms share valence electrons to get stable group VIIIA e- configurations.

II. Covalent Bonds A. Introduction Covalent bond = sharing of 2 electrons. 2 shared electrons with (Single Bond). 4 shared electrons with (Double Bond). 6 shared electrons with (Triple Bond). We frequently show the structure as a Lewis Structure - covalent bonds with lines and nonbonding valence electrons as dots. - Note: Group IVA usually forms 4 bonds; VA three bonds; VIA two bonds; and VIIA (along with H) one bond.

II. Covalent Bonds B. Examples H. + F::: ---) H F::: H. + O + .H ---) H O H :N + N: ---) :N N: :::Cl. + .O. + .Cl::: ---) :::Cl O Cl::: ::O: + :C: + :O:: -----) ::O = C = O::

II. Covalent Bonds C. Lewis Structures 1 II. Covalent Bonds C. Lewis Structures 1. Rules for Drawing Lewis Structures 1. Calculate the total # of valence electrons; take into account charge if the sample is an ion. 2. Place atom that forms most bonds at center (Closest to Group IVA & Lowest if in same group). If there is a charge, then add or subtract the appropriate number of electrons on the central atom. 3. Arrange other atoms around central atom & allow sharing so that each atom has stable electron configuration. Show bonding pairs as dashes & nonbonding valence e- as dots. 4. Double check: a) each atom has a stable electron configuration & b) have the same total number of valence electrons as in step 1.

II. Covalent Bonds C. Lewis Structures 2. Examples H I H2O NH4+ H2O2 CH4 SO2 AlCl4- NO2- CN-

Bonding Summary Two General Bonding Types Ionic: Compound containing metallic element. Atoms lose/gain e to look like nearest inert gas. Add together ions such that neutralize charge. Ia IIa IIIa Va VIa VIIa +1 +2 +3 -3 -2 -1 2. Covalent: Compound containing nonmetals. Atoms obtain inert gas configuration by sharing valence electrons. : :: :::

II. Covalent Bonds – Organic Compounds Can write organic structures several ways. Example – Butane (Note the five ways of presenting) Note: Carbon always has four bonds. C4H10 CH3CH2CH2CH3 CH3-CH2-CH2-CH3 H H H H H – C – C – C – C – H H H H H

II. Covalent Bonds – Organic Compounds Cyclic Organics: Example of Cyclopropane Aromatics Contain Benzene, C6H6

II. Covalent Bonds – Organic Compounds C7H6O3 MW = 138g C9H8O4 MW = 180 g

II. Covalent Bonds – Organic Compounds – Aspirin Lab 1) Equation & Conversion Factors: 1 Salicylic Acid + 1 Acetic Anhydride -----) 1 Aspirin + 1 Acetic Acid 1 = molecules or moles; 1 mole = formula weight in grams = 6.0x1023 molecules 2) Lab Calculations (questions 2 & 3): 2.0 g SA x 1 mole SA = 0.014 mole SA 138 g SA 0.014 mole SA x 1 mole Aspirin = 0.014 mole Aspirin 1 mole SA From the coefficients in the balanced chemical equation above.

III. Shapes Molecular Shapes play a major role in: 1) Physical Properties 2) Chemical Properties 3) Biochemical Properties To Obtain the shape of a molecule one draws the Lewis Structure, counts the number of “things” around the central atom, and uses simple geometry to predict the shape.

III. Shapes C. Simplified Examples Bond angle = 180o Bond angle = 120o Bond angle = 109o

IV. Nomenclature A. Introductions There are common & systematic names for chemicals. A chemical may have scores of common names. A systematic name must allow one to both obtain the formula and derive the name from the formula. There are two general rules for naming inorganic compounds. Ionic compounds use Rule #1. Molecular or Covalent compounds use Rule #2.

IV. Nomenclature B. Ionic Compounds Rule #1 for ionic compounds: Name the + element, then the – element and change the ending to “ide.” Examples: NaCl = Sodium Chloride Na2O = Sodium Oxide

IV. Nomenclature B. Ionic Compounds Rule #1 – “ide” names Negative atoms have an “ide” ending. Atom Anion Name Chlorine Cl1- Chloride Oxygen O2- Oxide Fluorine F1- Fluoride Sulfur S2- Sulfide Nitrogen N3- Nitride Iodine I1- Iodide Bromine Br1- Bromide Phosphorus P3- Phosphide

IV. Nomenclature B. Ionic Compounds Examples NaCl Na2O AlF3 Be3N2 Calcium Sulfide Barium Iodide Barium Oxide Magnesium Nitride Sodium Chloride Sodium Oxide Aluminum Fluoride Beryllium Nitride CaS BaI2 BaO Mg3N2

IV. Nomenclature C. Molecular Compounds Rule #2 When nonmetals & H combine with each other through sharing electrons (covalent bonds), they form molecules; there are no ions. Rule #2 – When both elements are nonmetals (molecular compounds), then Name the + & the - & change ending to “ide” as before. Use prefixes of di, tri, tetra, penta, etc to tell how many of each element is present.

IV. Nomenclature C. Molecular Compounds CO2 = Carbon Dioxide CCl4 = Carbon Tetrachloride N2O = Dinitrogen Oxide P2S5 = Diphosphorus Pentasulfide PBr3 = Phosporus Tribromide BI3 = Boron Triiodide Notes: (1) Organic compounds like CH4 use their own rules which we won’t cover. (2) diatomic molecules named with the element name. O2 = Oxygen

V. Polyatomic Ions Previous compounds formed from two elements. Frequently have compounds formed from three or four elements. When this happens, then usually have a polyatomic ion present. Polyatomic ions: stable ions formed from two or more elements; held together by covalent bonds. Examples: SO4-2 = Sulfate NO2- = nitrite PO4-3 = Phosphate

V. Polyatomic Ions Polyatomic ions are held together by covalent bonds, and they form ionic bonds with metals. Examples: NaNO2 Na2SO4 Na3PO4 When have more than one polyatomic ion in a compound then use parentheses around the ion. Examples: Na2SO3 Ca(NO2)2 Ca3(PO4)2 Nomenclature: Simply use the polyatomic ion name. Example: Calcium Nitrite & Calcium Phosphate above Need to memorize the following polyatomic ions, their names and their charges.

V. Polyatomic Ions - Memorize the Names, Formulas and the Charges Formula Name Formula Name NH4+ Ammonium (The Only Positive One in this list) C2H3O2- Acetate CN- Cyanide NO3- Nitrate NO2- Nitrite OH- Hydroxide HCO3- Hydrogen Carbonate CO3-2 Carbonate Cr2O7-2 Dichromate SO4-2 Sulfate SO3-2 Sulfite PO4-3 Phosphate

V. Polyatomic Ions – Examples of Naming and Obtaining Formulas Aluminum Hydroxide Calcium Cyanide Barium Sulfate Ammonium Nitrate Ba(OH)2 LiNO2 KNO3 NaHCO3 Al2(SO4)3 Al(OH)3 Ca(CN)2 BaSO4 NH4NO3 Barium Hydroxide Lithium Nitrite Potassium Nitrate Sodium Hydrogen Carbonate Aluminum Sulfate

Naming: Mixed Examples NaF CS2 NI3 BaI2 K3PO4 Boron Trifluoride Sodium Sulfite Sodium Fluoride Carbon Disulfide Nitrogen Triiodide Barium Iodide Potassium Phosphate BF3 Na2SO3