Ch 4 Chemical Bonding and Molecular Geometry Dr

Ch. 4: Chemical Bonding and Molecular Geometry Dr. Namphol Sinkaset Chem 200: General Chemistry I

I. Chapter Outline I. III. IV. V. VII. Introduction Ionic Bonding Covalent Bonding Chemical Nomenclature Lewis Dot Symbols and Structures Formal Charges and Resonance Molecular Structure and Polarity

I. Introduction • When elements form compounds, the original properties of the elements are lost.

I. To Lower Potential Energy • A chemical bond is the force that holds atoms together in a compound. • But why would atoms want to join with other atoms? • It all comes back to positive-negative attractions between particles in the atom which lead to lower PE!

I. Two Main Ways to Lower PE

I. Metal + Nonmetal = Ionic

I. Nonmetal + Nonmetal = Covalent • Instead of transferring e-’s, covalent bonding occurs via sharing of e-’s. • Attraction to two nuclei lowers PE.

II. Ionic Bonding • In ionic bonding, metal transfers e- to the nonmetal. • Metals have low IE’s; nonmetals have high EA’s. • Resulting ions attracted by +/- charge. • Number of cations and anions must lead to a neutral compound. § Charge on one becomes subscript on other.

II. Not Just One Ionic Bond

II. Cation Electronic Structure • Cations are often isoelectronic with noble gases. • Lower groups have partial loss of e-’s due to inert pair effect – stability of paired s e-’s. • Transition metals lose higher n e-’s first!

II. Anion Electronic Structure • Most monatomic ions form when valence s and p orbitals are “filled. ” • Gaining e-’s makes anions isoelectronic with the next noble gas. § e. g. Se 2 -: [Ar] 3 d 104 s 24 p 6 = Kr • Recall that charges for main-group monatomic ions can be deduced from location on periodic table.

II. Properties of Ionic Compounds • Ionic compounds have some distinguishing characteristics: § § Have crystalline structure Are rigid and brittle Have high melting and boiling points Conduct electricity when molten or when dissolved in water

II. Ionic Compounds Melt at High Temperatures • Why are such high temperatures needed?

II. Electrical Conductivity

III. Covalent Bonding • In covalent bonding, nonmetals share some (or all) of their valence electrons. • They share because the atoms involved have identical or similar IE’s and EA’s. • The formation of the bond can be explained energetically.

III. Formation of H 2

III. Fully Equal Sharing • If atoms involved in a covalent bond are the same, then e-’s will be shared equally. § Called a pure covalent bond. • Examples include all the diatomics like H 2, Cl 2, Br 2, etc.

III. Unequal Sharing • If atoms in a covalent bond are different, then e-’s aren’t shared equally. • The e-’s will be more attracted to one of the atoms resulting in a polar covalent bond.

III. Determining Bond Polarity • Is there a way to predict bond polarity? • Linus Pauling derived electronegativity values from looking at energies needed to break different types of bonds.

III. Electronegativity • To determine which atom in a polar covalent bond holds the partial negative charge, we use electronegativity. • Electronegativity is the relative ability of a bonded atom to attract shared e-. § It can be thought of as how greedy an atom is for e- when it is sharing them. • Note: electronegativity ≠ electron affinity.

III. EN Values

III. Using ΔEN • Differences in electronegativity can be used to determine the bond type.

III. Properties of Covalent Compounds • Covalent compounds have some distinguishing characteristics: § Have low melting and boiling points relative to ionic compounds § Tend to be insoluble in water § Are poor electrical conductors in any state

IV. Chemical Nomenclature • Naming inorganic compounds is systematic. • Key is identify what type of compound you are naming. • Variables include: § § Ionic or covalent? Single or multiple charge states? Polyatomic ion(s) involved? Acid?

IV. Ionic Nomenclature • Ionic compounds are named systematically, broken into two groups.

IV. Type I Compounds • Type I compounds are ionics that have a metal from Groups 1 or 2 and a nonmetal from Groups 14 -17. • Examples: § Na. Cl = sodium chloride § Mg. Br 2 = magnesium bromide § K 2 S = potassium sulfide

IV. Type I Compounds • To get a formula from a name, remember that a compound must be neutral. • Ion charges can be found by locating the element on the periodic table. • “The charge on one becomes the subscript of the other. ”

IV. Type I Compounds • e. g. What are the formulas for sodium nitride, calcium chloride, potassium sulfide, and magnesium oxide?

IV. Transition Metals • Transition metals are found in the “Valley, ” Groups 3 -12, of the periodic table. • Transition metal cations often carry different charges, e. g. Fe 2+ and Fe 3+. • Thus, a name like “iron chloride” is ambiguous.

IV. Type II Compounds • Type II compounds are ionics that have a transition metal (Groups 3 -12) and a nonmetal (Groups 14 -17). • Examples: § Fe. Cl 2 = iron(II) chloride § Fe. Cl 3 = iron(III) chloride

IV. Type II Compounds • e. g. Give the correct name or formula for the compounds below. a) b) c) d) e) Mn. O 2 copper(II) chloride Au. Cl 3 molybdenum(VI) fluoride Hg 2 Cl 2

IV. Type II Compounds • An archaic naming system uses common names for transition metal cations of different charge. § Higher charge given –ic suffix § Lower charge given –ous suffix • Fe. Cl 3 = ferric chloride • Fe. Cl 2 = ferrous chloride

IV. Additional Complications • To make naming ionic compounds harder, sometimes polyatomic ions are involved. • polyatomic ion: an ion composed of two or more atoms

IV. Common Polyatomic Ions Name Formula ammonium NH 4+ hydrogen sulfate HSO 4 - permanganate Mn. O 4 - sulfite SO 32 - peroxide O 22 - hydrogen sulfite HSO 3 - hydroxide OH- phosphate PO 43 - acetate CH 3 COO- hydrogen phosphate HPO 42 - cyanide CN- dihydrogen phosphate H 2 PO 4 - azide N 3 - perchlorate Cl. O 4 - carbonate CO 32 - chlorate Cl. O 3 - bicarbonate HCO 3 - chlorite Cl. O 2 - nitrate NO 3 - hypochlorite Cl. O- nitrite NO 2 - chromate Cr. O 42 - sulfate SO 42 - dichromate Cr 2 O 72 -

IV. Oxyanion Families • Oxyanions are anions that contain oxygen and another element. • There are families of oxyanions, and they have a systematic naming system. • Have either two- or four-member families. § e. g. NO 2 - and NO 3§ e. g. Cl. O-, Cl. O 2 -, Cl. O 3 -, and Cl. O 4 -

IV. Two-Member Families • For a two-member family, oxoanion with fewer O atoms is given the “–ite” suffix while the one with more O atoms is given the “–ate” suffix. § e. g. NO 2 - = nitrite and NO 3 - = nitrate

IV. Four-Member Families • For the four-member families, the prefixes “hypo-” and “per-” are used to indicate fewer or more oxygen atoms. • e. g. the chlorine oxoanions § § Cl. O- = hypochlorite Cl. O 2 - = chlorite Cl. O 3 - = chlorate Cl. O 4 - = perchlorate

IV. Hydrated Ionic Compounds • Ionics with trapped waters are called hydrates. • Greek prefixes are used to indicate #’s of trapped waters. • e. g. cobalt(II) chloride hexahydrate.

IV. Naming Practice • e. g. Give names or formulas for the following compounds. a) b) c) d) e) f) Na 2 CO 3 magnesium hydroxide Cu. SO 4· 5 H 2 O Co. PO 4 nickel(II) sulfate Na. Cl. O 2

IV. Covalent Nomenclature • For covalent compounds, many different compounds can exist from the same two elements. § e. g. NO, NO 2, N 2 O 3, N 2 O 4, N 2 O 5! • Therefore, we need a systematic naming method.

IV. Type III Compounds • • Type III compounds are covalent (nonmetal bonded to nonmetal). Naming rules: 1) Element w/ lower group # is named 1 st using the normal element name EXCEPT when halogens are bonded to oxygen. 2) If elements are in the same group, lower element named first. 3) Second element is named using its root and the “ -ide” suffix. 4) #’s of atoms indicated with Greek prefixes EXCEPT when there is only one atom of the first element.

IV. Greek Prefixes

IV. Type III Compounds • Some examples: § § Cl. O 2 = chlorine dioxide N 2 O 5 = dinitrogen pentoxide S 2 Cl 2 = disulfur dichloride Se. F 6 = selenium hexafluoride

IV. Acids • There are many definitions of acids and bases, but we will use the Arrhenius definitions for now. • Acids are molecular compounds that produce H+ ions in aqueous solution. • We will look at the chemistry of acids in a later chapter.

IV. Acid Nomenclature • There are two categories of acids that have different naming rules. 1) Binary acids 2) Oxyacids • You can easily recognize acids because their formula has H as the first element.

IV. Naming Binary Acids • Binary acids contain a nonmetal anion. § § HCl = hydrochloric acid HBr = hydrobromic acid H 2 Se = hydroselenic acid HI = hydroiodic acid

IV. Naming Oxyacids contain an oxyanion. • Set 1 • Set 2 § § § § HNO 3 = nitric acid H 2 SO 4 = sulfuric acid HCl. O 3 = chloric acid HCl. O 4 = perchloric acid HNO 2 = nitrous acid HCl. O 2 = chlorous acid HCl. O = hypochlorous acid H 2 SO 3 = sulfurous acid

IV. Naming Practice • e. g. Give the correct formula or name of the compounds below. a) b) c) d) e) f) g) h) i) j) k) l) m) n) Co. Cl 3 dichlorine heptaoxide HBr(aq) Sr. O magnesium hydroxide carbon tetrachloride Mg. SO 4· 7 H 2 O sodium hydride HIO 3 V 2 O 5 Ru(Cl. O 4)3 NI 3 titanium(IV) oxide N 2 F 2

V. Lewis Dot Symbols • • • Valence e-’s are the most important e-’s in bonding. Lewis dot symbols are a way to depict the valence e-’s of atoms. Lewis dot symbols have two parts: 1) element symbol: represents nucleus and core e 2) dots around symbol: represent valence e-’s

V. Lewis Dot Symbols • The number of valence e- is given by the element’s group number!!

V. Depicting Ionic Bonds

V. The Octet Rule • Lewis dot structures are more often used for covalent bonding. • Main group elements tend to form enough bonds to have 8 valence e-’s. • Known as the octet rule (duet for H and He). • Using Lewis dot symbols, it becomes like a puzzle.

V. Single Bonds • Single bonds are the basic connections that hold covalent compounds together. • Single bonds form when one pair of e-’s are shared. • Note that lone pairs are e- pairs that don’t participate in bonding.

V. Double Bonds • When more e-’s need to be shared to reach an octet, a double bond is possible.

V. Triple Bonds • When even more e-’s need to be shared, a triple bond is possible. § e. g. molecular nitrogen, : N≡N: • As the bond order increases, the bond gets stronger and shorter.

V. Steps for Drawing Lewis Structures 1) Determine total # of valence e-. 2) Place atom w/ lower Group # (lower electronegativity) as the central atom. 3) Attach other atoms to central atom with single bonds. 4) Fill octet of outer atoms. (Why? ) 5) Count # of e- used so far. Place remaining e- on central atom in pairs. 6) If necessary, form higher order bonds to satisfy octet rule of central atom. 7) Allow expanded octet for central atoms from Period 3 or higher (n ≥ 3).

V. Lewis Structure Practice • Draw correct Lewis structures for NF 3, CO 2, Se. Cl 2, PI 5, IF 2 -, IF 6+, and H 2 CO.

V. Exceptions to the Octet Rule • We’ve already discussed expanded valence (hypervalent molecule) cases, but there are other exceptions as well. § Compounds w/ odd # of e-’s: free radicals. Examples include NO and NO 2. § Incomplete octets: e- deficient atoms like Be and B, e. g. Be. Cl 2 and BF 3. § Expanded octets – when d orbitals are used to accommodate more than an octet.

VI. Multiple Valid Lewis Structures • Sometimes more than one Lewis structure can be drawn for the same molecule. • For example, ozone (O 3).

VI. Resonance Structures • Resonance structures are also known as resonance forms. • A resonance structure is one of two or more Lewis structures that have the same skeletal structure (atoms in same place), but different electron arrangements.

VI. Resonance Hybrid • Neither resonance form is a true picture of the molecule. • The molecule exists as a resonance hybrid, which is an average of all resonance forms. • In a resonance hybrid, e- are delocalized over the entire molecule.

VI. Sample Problem • Draw the resonance structures of the carbonate anion.

VI. Important Resonance Forms • If all resonance forms have the same surrounding atoms, then each contributes equally to the resonance hybrid. • If this is not the case, then one or more resonance forms will dominate the resonance hybrid. • How can we determine which forms will dominate?

VI. Formal Charge • formal charge: the charge an atom would have if bonding e- were shared equally formal charge = (# valence e-) – (# lone pair e-) – (½ # bonding e-)

VI. Formal Charges in O 3 • We calculate formal charge for each atom in the molecule. • For oxygen atom A (left structure), there are 6 valence e-, 4 lone pair e-, and 4 bonding e-. The formal charge for this O atom is 0. • NOTE: sum of all formal charges must equal the overall charge of the molecule!

VI. Using Formal Charges • Formal charges help us decide the most important resonance forms (or most likely structure): 1) Structure w/ all f. c. ’s equal to 0 preferable. 2) Structure with smallest nonzero f. c. ’s preferable. 3) Structure with adjacent f. c. ’s of 0 or of opposite sign preferable. 4) Structure w/ negative f. c. ’s on higher EN elements preferable.

VI. Sample Problem • Find the dominant resonance structures for the sulfate anion.

VI. Sample Problem • Draw the three possible Lewis structures for CO 2 and use formal charges to show which structure is most likely.

VII. Chemistry Happens in 3 D! • To understand how molecules react and interact with each other, we need to know how they look in 3 dimensions. • Important details include bond angle and bond length.

VII. VSEPR Theory • From a correct Lewis structure, we can get to the 3 D shape using this theory. • VSEPR stands for valence shell electron pair repulsion. • The theory is based on the idea that epairs want to get as far away from each other as possible!

VII. VSEPR Categories • There are 5 electron pair geometries from which all molecular structures derive.

VII. VSEPR Categories

VII. Drawing w/ Perspective • We use the conventions below to depict a 3 D object on a 2 D surface.

VII. Determining 3 D Shape • The 5 electron pair geometries (EG) are a starting point. • To determine the molecular structure (MG), we consider the # of atoms and the # of e- pairs that are associated w/ the central atom. • All the possibilities for molecular structure can be listed in a classification chart.

VII. Linear/Trigonal Planar Geometries • First, we have the linear and trigonal planar EG’s. EG Bonds Lone Pairs MG Linear 2 0 linear Trigonal planar 3 0 2 1 trigonal planar bent

VII. Tetrahedral Geometries EG Bonds Lone Pairs MG Tetrahedral 4 3 0 1 tetrahedral pyramidal 2 1 2 3 bent linear

VII. Trigonal Bipyramidal Geometries EG Trigonal Bipyramidal Bonds Lone Pairs 5 0 4 3 2 1 1 2 3 4 MG trigonal bipyramidal see-saw T-shaped linear

VII. Octahedral Geometries EG Octahedral Bonds Lone Pairs MG 6 0 octahedral 5 1 4 2 square pyramidal square planar 3 2 1 3 4 5 T-shaped linear

VII. Summary Chart

VII. Steps to Determine Molecular Structure 1) Draw Lewis structure. 2) Count # of bonds (double/triple count as one) and lone pair e-’s on the central atom. 3) Selectronic pair geometry. 4) Place e-’s and atoms that lead to most stable arrangement (minimize e- repulsions). 5) Determine molecular geometry.

VII. Trig Bipy is Special • In other EG’s, all positions are equivalent. • In trig bipy, lone pairs always choose to go equatorial first. • Why?

VII. Lone Pairs Take Up Space • Lone pair e-’s don’t have another nucleus to “anchor” them.

VII. Distortion of Angles • Lone pair e-’s take up a lot of room, and they distort the optimum angles seen in the EG’s.

VII. Some Practice • Draw the molecular geometries for SF 4, Be. Cl 2, Cl. O 2 -, Te. F 5 -, Cl. F 3, and NF 3.

VII. Larger Molecules

VII. Molecular Polarity • Individual bonds tend to be polar, but that doesn’t mean that a molecule will be polar overall. • To determine molecular polarity, you need to consider the 3 D shape and see if there’s a net dipole moment. § Net dipole moments arise when bond dipole moments (vectors) don’t cancel.

VII. Bond and Net Dipole Moments

VII. Sample Problem • Determine the molecular structure of IF 2 - and state whether it is polar or nonpolar.

VII. Polarity and Properties • Polarity is the result of a compound’s composition and structure. • Knowing that a compound is polar/nonpolar allows us to explain some of its properties.

VII. Alignment in Electric Field

VII. “Like Dissolves Like”