Sheets Page 1 Lecture 8

Lecture 8: Chemical bonding 1 Read: BLB 2.6–2.8; 8.1–8.4 HW: BLB 2:51,53,57,59,65,67; 8:16,22,25,29,37,39 Sup 2:4 Know: • chemical bonding ionic bonding covalent bonding metallic bonding • Lewis symbols • lattice energy Exam 1: Monday, Feb 9 @ 6:30!!! start preparing now!! only non-text programmable calculators are allowed—no PDAs, blackberries, cell phones, etc. will be permitted. Bring: pencils, student ID and a calculator—Absolutely NO text-programmable calculators or wireless devices (will be checked) Form a study group, use the CRC, take advantage of SI (info on web), use the online resources, and work those problems—practice, practice, practice Bonus deadline for Skill check test 5 is Thursday, 2/5 & Skill check test 6 is SUNDAY, 2/8

Sheets’s office hours: Mondays 12:30-2; Tuesdays 10:30-12 in 324 Chem (or 326 Chem).

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Chemical bonding • what is a chemical bond, anyway?

⇒ force holding atoms or ions together & it is potential energy that can be exploited

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3 types of chemical bonds ionic • valence electrons traded to form separate ions; attraction between positive & negative ions • each ion has noble gas e– configuration covalent • valence electrons between a few nuclei • distinct molecules; molecules = molecular compounds metallic (will see in Chem 112) • valence electrons shared among all nuclei = “electron sea”; conduction; found in metals

a

Na+ Cl–

a

valence e–

core e–

a

valence e–

core e–

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Dots & octets • valence electrons are involved in bonding • a dot (•) = a e– . C: 1s22s22p2 · C · · 4 valence electrons . S: [Ne] 3s23p4 :S: · 6 valence electrons • octet rule: elements tend to gain, lose or share electrons to achieve an inert gas configuration; compounds form to achieve octet for all atoms • ns2np6 valence = octet

⇒ inert gases have 8 valence electrons

(“duet” for H and He)

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• knowing the valence electrons, you should be able to easily draw Lewis structures; DO NOT memorize this table! You can figure it out!! :) • more about Lewis structures in next lecture (talk about a cliffhanger…)

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Transition metals (talked about last time in Lecture 7…)

• TM are exceptions to the octet rule! • lose s electrons first • lose d electrons only after valence s gone • transition metal ions can have variable charges examples Fe: [Ar] 4s23d6 Fe2+: [Ar] 3d6 Fe3+: [Ar] 3d5 Ag: [Kr] 5s14d10 Ag+: [Kr] 4d10

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Ionic bonding • electrostatic attraction between ions of opposite charge

• Coulomb's law: E !Q1Q2

d

Q1 : charge on 1st atom Q2 : charge on 2nd atom d : distance between charges

• e–ʼs exchanged to form separate ions with complete octets Na+ Cl− 1s22s22p6 1s22s22p63s23p6 [Ne] [Ar]

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This should be a review for you; we wonʼt discuss in class. Look at placement in periodic table and relate to charge of cation or anion. We will refer to these by name throughout the semester. Know these!!!

Common cations Common anions 1+ 1–

H+ hydrogen ion H– hydride Na+ sodium F– fluoride K+ potassium Cl– chloride Ag+ silver Br– bromide

NH4+ ammonium I– iodide

OH– hydroxide

2+ NO3– nitrate

Mg2+ magnesium CN– cyanide Ca2+ calcium Ba2+ barium 2– Co2+ cobalt(II) O2– oxide Cu2+ copper(II) S2– sulfide Fe2+ iron(II) CO3

2– carbonate Pb2+ lead(II) SO4

2– sulfate Mn2+ manganese Hg2+ mercury(II) 3– Zn2+ zinc PO4

3– phosphate

3+ Al3+ aluminum Cr3+ chromium(III) Fe3+ iron(III)

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Properties of ionic compounds • distinct + & − ions • ∞ crystal lattice • strong omnidirectional bonds • characteristics

• low electrical conductivity as solids (high as liquids)

• very high melting, boiling points • hard but brittle • soluble only in polar solvents (water): electrolytes (more ~Lecture 28)

• formed from what???? metal + non-metal

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Strength of ionic bonds (lattice energy)

• lattice energy: the change in energy when an ionic solid is separated into isolated ions in the gas phase NaCl(s) → Na+(g) + Cl–(g) ΔE = 788 kJ = lattice energy • lattice energy cannot be directly determined experimentally [• therefore use Hessʼs Law [Chap 5] (and known atomic properties) to get the value of the lattice energy (more on Hessʼs Law [~Lectures 38 & 39]) • Born-Haber cycle: a thermochemical cycle used to analyze factors contributing to stability of ionic compounds (BLB Fig. 8.4)]

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More on lattice energy NaCl E = 786 kJ MgO E = 3934 kJ ⇐ factor of >4! for MgO • why the big difference??? • different salts have different Qʼs and d NaCl MgO • radii Na+ 0.97 Å Mg2+ 0.66 Å

Cl − 1.81Å O2– 1.40Å

d 2.78Å 2.06Å d is smaller for • thus, lattice energy for MgO is >4× that of NaCl

E!Q1Q2

d

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Lattice energy of salts • melting points are indicative of lattice energy; higher M.P. ⇒ L.E. (MP ≠ LE) • relative magnitudes of lattice energy predictable based on charges (Q1, Q2) and separation (d) • recall(!) ion size predictions from periodic trends anion size MP NaF 993°C NaCl 801°C NaBr 747°C ⇒ d NaI 661°C ⇒ L.E. MgO 2800°C higher charge ⇒ Q1Q2 ⇒ L.E.

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Example: Which one of the following expresses the correct relationship between the melting points of the 3 ionic solids: NaCl, MgS, and BaS?

A. NaCl > MgS > BaS B. NaCl > BaS > MgS C. MgS > BaS > NaCl D. MgS > NaCl > BaS E. BaS > MgS > NaCl

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Covalent bonding • when neither atom is “willing” to give up electrons (completely) • atoms electrons: each atom has a noble gas configuration

H H+ H H

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Covalent bonding and Lewis structures • covalent bond: build-up of electron density between 2 nuclei or “sharing” of electron pair • Lewis structures: show how electrons are shared (atom connectivity);

shared e– pairs ⇒ BOND also display unshared e– pairs, or pairs example NH3

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More on covalent bonding • more electrons shared, the bond

bond strengths ~100–1000 kJ/mole • more electrons shared, the bond • single, double & triple bonds allowed & these multiple bonds have distinctive properties N−N N=N N≡N l 1.47 Å 1.24 Å 1.10 Å E 163 418 941 kJ/mole bond length, l > > bond strength, E (kJ/mol) < < weʼll talk more about multiple bonds in the coming lectures!!

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Covalent bonding (cont.) • to understand covalent bonding at molecular (atomic) level, we will explore the concept at three levels:

Lewis structures: tells where e– are (analogous to a census telling where people are) VSEPR theory: (Chap 9) shows how to use Lewis structures to predict shapes of molecules valence bond theory: (Chap 9) explains how covalent bonds are the result of orbital overlap (σ and π bonds)

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Before next class: Read: BLB 2.6–2.8; 8.1–8.4 HW: BLB 2:51,53,57,59,65,67; 8:16,22,25,29,37,39 Sup 2:4 Know: • electronegativity • Lewis structures • formal charges Answer: p 13: C