Covalent Bonding C H A P T E R 9

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1 Covalent Bonding C H A P T E R 9

2 I. The Covalent Bond A. Covalent bond: chemical bond resulting from the sharing of valence electrons 1. Occurs when atoms have similar attractions for electrons neither atom can gain or lose electrons, so they share Covalent bond:

3 B. Molecules are formed when two or more atoms bond covalently; considered individual units (ex: H 2 O is one molecule) 1. Diatomic molecules: molecules composed of only two atoms a. May be composed of atoms of 2 elements (ex: NaCl, CO, NO) b. May be composed of 2 atoms of the same element i. H O N Cl Br I F hydrogen nitrogen bromine fluorine H 2 N 2 Br 2 F 2 oxygen chlorine iodine O 2 Cl 2 I 2

4 2. Polyatomic molecules: composed of three or more atoms a. Ex: H 2 O, CCl 4, CO 2 C 2 H 4 polyatomic FeO Cl 2 P 4 diatomic diatomic *(with ionic bonding) polyatomic HCl diatomic

5 C. Single Covalent Bonds 1. Formed when one pair of electrons is shared 2. Ex: H 2 H + H H:H can be written as H:H or H-H 3. H:H or H-H is an example of a Lewis structure a. Use electron-dot diagrams to represent bonding in molecules b. One line or one pair of dots represents one covalent bond 4. Single bonds are also called sigma (σ) bonds a. Form when orbitals overlap: can be two s, one s and one p, or two p

6 D. Multiple Covalent Bonds 1. Formed when two or more pairs of electrons are shared 2. Examples: double bond triple bond triple bond 3. Double & triple bonds contain pi bonds, when parallel orbitals overlap to share electrons a. Double bonds: one sigma, one pi (two shared pairs) b. Triple bonds: one sigma, two pi (three shared pairs)

7 E. Strength of Covalent Bonds 1. Bond length: distance between bonded atoms 2. Rule: shorter bond length = stronger bond 3. Bond length is longest in single bonds shorter in double bonds, shortest in triple bonds 4. So. weakest stronger strongest

8 5. Bond dissociation energy: amount of energy required to break a specific covalent bond a. Ex: F 2 (single) 159 kj/mol O 2 (double) 494 kj/mol N 2 (triple) 945 kj/mol b. Amount of energy available in a compound is the sum of all bond dissociation energies of all bonds in the compound c. In chemical rxns, bonds are formed, which gives off energy, and broken, which takes in energy leads to overall loss or gain of energy i. If more energy is taken in (gained): endothermic ii. If more energy is given off (lost): exothermic

9 Classwork p. 244 #1-5 (see example problem 9-1 on same page for guidance) p. 247 #6-12

10 II. Naming Molecules A. Binary Molecular Compounds 1. First element in the formula is first use entire element name 2. Second element is named using the root of the element name & adding suffix ide 3. Prefixes are used to indicate the number of atoms of each type present in the compound

11 4. Examples: a. P 2 O 5 b. CO c. CO 2 5. Some molecular compounds were named before the modern naming system & are still known by their common names a. Examples: i. Water: H 2 O, dihydrogen monoxide ii. iii. Ammonia: NH 3, nitrogen trihydride laughing gas: N 2 O, dinitrogen monoxide, nitrous oxide

12 B. Acids 1. Formed by dissolving a dry molecular compound in water, producing H+ ions 2. Naming Binary Acids a. Binary acids contain hydrogen & one other element b. One word, followed by acid c. Use the prefix hydro- to represent hydrogen, then the root of the second element with the suffix -ic d. Examples: i. HCl: hydrochloric acid ii. iii. HBr: hydrobromic acid HF: hydrofluoric acid

13 3. Naming Oxyacids a. Oxyacids contain hydrogen and an oxyanion (a polyatomic ion containing oxygen) b. One word, followed by acid c. Look up the name of the oxyanion (on PT) d. Use the root of the oxyanion followed by a suffix: i. -ic if the suffix of the ion is -ate ii. e. Examples: -ous if the suffix of the ion is -ite i. HNO 3 : nitric acid ii. HNO 2 : nitrous acid

14 4. Common Acids (MUST KNOW!) a. HCl: hydrochloric acid (stomach acid) b. H 2 SO 4 : sulfuric acid (acid rain) c. HNO 3 : nitric acid (fertilizer production) d. H 2 CO 3 : carbonic acid e. H 3 PO 4 : phosphoric acid (pop) f. HClO 4 : perchloric acid (explosive!) g. CH 3 COOH: acetic acid (vinegar!) h. (COOH) 2: oxalic acid (organic plants)

15 Homework Copy Figure 9-9 (p. 251) p. 249 #13-17 p. 250 #18-22 p. 251 #23-29 Worksheet!

16 III. Molecular Structures A. Structural Formulas 1. Definition: molecular model using letter symbols & bonds to show relative positions of atoms 2. Can be predicted by drawing the Lewis structure B. Steps to draw a Lewis structure: 1. Find the total number of valence electrons this is the total number of electrons available for bonding a. Divide # above by 2 to find the total number of electron pairs 2. Determine atom locations a. The central atom will be the one that can form multiple bonds b. The terminal atoms (at the end) will usually only form single bonds. Hydrogen is ALWAYS a terminal atom

17 3. Draw the skeleton structure the symbols of atoms in their predicted arrangement 4. Place single bonds between the central atom and each of the terminal atoms a. Subtract the # of pairs used from the total # of bonding pairs found in step 1 this # represents all of the lone pairs and any double or triple bonds b. Place lone pairs around the terminal atoms if necessary to satisfy the octet rule if there are odd numbers on any terminal atoms when you ve done this, that means there will be a double bond between that atom & the central atom 5. Check how many electron pairs are around the central atom. If there are not 4 pairs, there will need to be double or triple bonds. a. Change one or two of the pairs on a terminal atom to a double or triple bond between that terminal atom & the central atom to satisfy the central atom s octet

18 6. Examples: a. Water: H 2 O b. Carbon dioxide: CO 2 c. Phosphate ion: PO 4 3-

19 Homework Copy Figure 9-10 (p. 252) p. 255 #30-34 p. 874 #1 a-d under Section 9-1 p. 875 #4 a-f under Section 9-3

20 C. Resonance Structures 1. Some compounds can have more than one valid Lewis structure most often when there are both double & single bonds 2. In these cases, the bonds resonate or switch off between structures 3. Results in a more stable configuration with average bond strengths & average bond lengths (bond length is not actually single or double it s in between) 4. Differ only in the position of electron pairs, not the atom location

21 D. Exceptions to Octet Rule (three possible situations) 1. Molecule has odd number of valence electrons & cannot pair each electron a. Example: NO 2, ClO 2, NO 2. Some compounds form with fewer than 8 electrons present around an atom, with electrons still paired up (even # < 8) a. Example: BH 3 b. Coordinate covalent bond: when one atom that has lone pairs will contribute an electron pair to another atom, which is short on electrons

22 3. Expanded octet: central atom has more than 4 electron pairs a. These molecules will often have 5 or 6 pairs around the central atom b. Common in highly reactive atoms c. Example: PCl 5, SF 6

23 Homework p. 256 #35-38 p. 258 #39-48 Worksheet!

24 Lab 8.1 Tips Groups 1 & 5: start building your models Groups 2 & 6: start with aspirin Groups 3 & 7: start with acetaminophen Groups 4 & 8: start with ibuprofen Rotate models to the group with the next number (4 should rotate to 1, 8 should rotate to 5) Each group has their own modeling kit, so you do NOT need to rotate that just the 3 models that I built! Long gray connectors are double bonds, short gray are single bonds When you are finished, take your models apart & put them back in the bag count how m any of each part you have, check off on the index card in your bag, and each group member should initial at the bottom of the card confirming that all parts are back in place

25 IV. Molecular Shape A. VSEPR Model 1. Valence Shell Electron Pair Repulsion: model based on an arrangement that minimizes the repulsion of shared & unshared pairs of electrons around the central atom 2. Atoms exist at fixed angles to one another 3. All electron pairs repel each other a. Shared pairs repel each other equally b. Lone pairs occupy a slightly larger orbital than shared pairs, pushing the shared pairs toward one another

26 B. Possible Geometric Configurations 1. Linear: atoms in a straight line or form 180 angle a. 2 pairs of shared electrons; no lone pairs b. Examples: Cl 2, BeF 2 2. Trigonal planar: atoms form a triangle at 120 from each other a. 3 pairs of shared electrons; no lone pairs b. Examples: BF 3, AlCl 3 3. Tetrahedral: atoms form a three-sided pyramid at from each other a. 4 pairs of shared electrons; no lone pairs b. Example: CH 4

27 4. Trigonal pyramidal: atoms form an uneven pyramid with bond angles of a. 3 pairs of shared electrons; 1 lone pair b. Examples: PH 3, NH 3 5. Bent: atoms form a bent shape with bond angles of a. 2 pairs of shared electrons; 2 lone pairs b. Example: H 2 O 6. Trigonal bipyramidal: three atoms around central in the same horizontal plane, and one atom directly above & one atom directly below the central in the vertical plane a. Forms 2 pyramids (top & bottom); bond angles are 90 between the vertical and horizontal atoms, and 180 between atoms in the same plane b. 5 pairs of shared electrons; no lone pairs c. Examples: NbBr 5, PCl 5

28 7. Octahedral: four atoms around central in the same horizontal plane, and one atom directly above & one atom directly below the central in the vertical plane a. Bond angles are 90 between the central atom and each terminal atom b. 6 pairs of shared electrons; no lone pairs c. Example: SF 6 HOMEWORK: copy Table 9-3 on p. 260 add a column between Example and Total Pairs for Lewis structures & draw those!

29 C. Hybridization 1. Atomic orbitals merge to form hybrid/combination orbitals at bonding sites 2. Explains why all bonding orbitals around certain atoms are the same instead of having bonding s orbitals, bonding p orbitals, etc. 3. This involves moving electrons to higher orbitals in order to free up bonding sites

30 4. Variations of hybrid orbitals a. sp orbitals: 1 s + 1 p orbital merge; forms triple bonds; linear structure b. sp 2 orbitals: 1 s + 2 p orbitals merge; forms double bonds; planar structure c. sp 3 orbitals: 1 s + 3 p orbitals merge; forms single bonds; tetrahedral structure d. sp 3 d orbitals: 1 s + 3 p + 1 d orbital merge; forms single bonds; trigonal bipyramidal structure e. sp 3 d 2 orbitals: 1 s + 3 p + 2 d orbital merge; forms single bonds; octahedral structure

31 p. 262 #49-59 Homework

32 V. Electronegativity & Polarity A. Electronegativity 1. Electronegativity (EN): relative ability of an atom to attract electrons in a chemical bond a. EN is an average value; slightly different for different bonds b. Most active metals have lowest EN, most active nonmetals have highest EN c. EN is a dimensionless quantity (not positive or negative)

33 B. Bond Character 1. Atomic bonds are not just purely ionic or covalent; it is a continuum of bond types 2. Bond strength depends on the EN difference between the two atoms a. Greater difference stronger bond b. When EN difference is large, one atom has a higher attraction & pulls electrons from the other atom to itself forms ionic bond c. When EN difference is small, atoms have nearly the same attraction for electrons, so they share forms covalent bond

34 3. When EN difference is 1.70, a bond is 50% ionic & 50% covalent a. EN difference > 1.70 shows ionic character b. EN difference < 1.70 shows covalent character c. Examples: i. NaCl: Na = 0.93 Cl = 3.16 difference = = > 1.70 ionic ii. H 2 O: H = 2.20 O = 3.44 difference = = < 1.70 covalent iii. O 2 : O = 3.44 O = 3.44 difference = = zero zero << % covalent / nonpolar covalent

35 C. Polarity 1. Polar covalent bond: occurs when a shared pair of electrons is attracted more strongly to one of the atoms; referred to as unequal sharing a. Arrangement of polar bonds determines overall polarity of a molecule b. Polar bonds produce polar molecules unless the polar bonds are symmetrically arranged (ex: CCl 4 ) i. Check VSEPR shape to see if molecule is symmetric 2. Electronegativity is used to determine polarity; greater EN difference indicates greater polarity a. Example 1: H 2 (H = 2.20) b. Example 2: HF (H = 2.20, F = 3.98) c. Example 3: H 2 O (H = 2.20, O = 3.44) d. Example 4: CCl 4 (C = 2.55, Cl = 3.16) 3. Polar molecules are often called dipoles because they have positive & negative poles a. Delta (δ) with a + or - is used to represent the partial positive/negative charges on each pole

36 Homework p. 266 #60-63, p. 875 #9 under Section 9-5 To determine polarity: If bonds are nonpolar, the molecule is nonpolar If bonds are polar: Determine VSEPR structure If polar bonds are symmetric (charges are balanced) and contain the same atoms (ex: CCl 4 ), then a molecule is nonpolar If polar bonds are not symmetric (as in bent & trigonal pyramidal), a molecule is polar Use electronegativity table on p. 263 If you need help on which VSEPR structures are polar, see next slide

37

38 D. Properties of Covalent Compounds 1. Covalent properties are due to the weak attractions between positive & negative poles, called intermolecular / van der Waals forces a. Dispersion forces: weak forces contained in nonpolar substances b. Dipole-dipole forces: slightly stronger forces contained in polar substances c. Hydrogen bond: strong intermolecular force found between the H end of a dipole & the atom on the other dipole 2. Characteristics of covalent compounds: a. Low melting points b. Do not conduct electricity c. Volatile/easily reactive due to weak bonds

39 E. Covalent Network Solids 1. Composed of atoms interconnected by a series of covalent bonds 2. Examples: quartz, diamond 3. Characteristics of covalent network solids: a. High melting points b. Brittle, very hard c. Do not conduct electricity Diamond Quartz

40 Homework p. 267 #64-70 Worksheet Read Lab 8.2 (p in lab manual)

41 Chapter 9 Review p. 271: Vocabulary p. 272: Concept Map #71 *DRAW IT* p. 272: Mastering Concepts #72-87 *skip 75 & 84 p. 272: Mastering Problems # *remember you already did either staple in or recopy HONORS ONLY p. 274: Mixed Review # HONORS ONLY p. 274: Thinking Critically # p. 275: Standardized Test Practice #1-10

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