Molecules, Compounds & Chemical Reactions

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1 Molecules, Compounds & Chemical Reactions
Overview Compounds Formulas Ionic compounds Molecular (covalent) compounds Molecular weight/molar mass Covalent Bonds Bonding models: Lewis, VSEPR, etc. Molecular geometry and chemical properties Chemical reactions Balancing equations Molar relationships

2 Chemical Representations
Representational: chemical formulas, equations CH4 + 2O2  CO H2O Chemistry Macroscopic: experimental Submicroscopic: atoms, molecules

3 Compounds Atoms are typically bound together in molecules
Noble gases exist as individual atoms, most other elements found as compounds More stable Forming a bond: releases energy, less stable to more stable Breaking a bond: requires energy, more stable to less stable (misconception)

4 Chemical Formulas AxByCz A, B, C = elemental symbols
x, y, z = relative number of atoms in molecule Problem 1: How many of each atom are in the following molecules? C6H12O6 (glucose) NH3 (ammonia) NaHCO3 (baking soda)

5 Problem 1: How many of each atom are in the following molecules?
C6H12O6 b) NH3 c) NaHCO3 (baking soda)

6 Problem 1: How many of each atom are in the following molecules?
C6H12O6 6 carbons, 12 hydrogens, 6 oxygens b) NH3 1 nitrogen, 3 hydrogens c) NaHCO3 (baking soda) 1 sodium, 1 hydrogen, 1 carbon, 3 oxygens

7 Ionic Compounds Ionic compounds are made up of ions held together by electrostatic forces: ionic bonds Cation (+) & anion (–) Cation (metal) + anion (nonmetal: monatomic or polyatomic) Na+ + Cl–  NaCl

8 Crystal Lattice—NaCl NaCl: ratio of atoms in the lattice

9 Ionic Compounds Dissociate when they dissolve in water
NaCl Na+(aq) + Cl–(aq) water

10

11 Ionic Compounds Monatomic ions: Cation: group number
Li+, Na+, K+, Rb+, Cs+ Mg2+, Ca2+ Al3+ Transition metals often variable: Cu2+, Cu+ (copper II and copper I) and Fe3+, Fe2+ (iron III and iron II)

12 Ionic Compounds Monatomic ions: Anion: 8 – group number
add to electrons to make 8 (octet) O  O2– oxide S  S2– sulfide Cl  Cl– chloride P  P3– phosphide

13 Ionic Compounds Balance anion with cation to a neutral charge:
magnesium oxide: Mg2+ with O2– MgO aluminum chloride: Al3+ with Cl– AlCl3

14 Problem 2 Write the ionic formulas for each of the following:
The compound that magnesium makes with chlorine The compound that Fe3+ makes with oxygen Sodium fluoride Cesium oxide Copper (I) sulfide

15 Problem 2 Write the ionic formulas for each of the following:
The compound that magnesium makes with chlorine The compound that Fe3+ makes with oxygen Sodium fluoride Cesium oxide Copper (I) sulfide MgCl2 Fe2O3 NaF Cs2O Cu2S

16 Polyatomic Ions CO32– carbonate ion Travels as a unit (Texans)
Na2CO3 sodium carbonate Fe2(CO3)3 iron (III) carbonate

17 Common Polyatomic Ions
nitrate NO3– nitrite NO2– sulfate SO42– sulfite SO32– phosphate PO43– carbonate CO32– hydroxide OH– ammonium NH4+

18 Problem 3: Write the ionic formulas for each of the following:
Magnesium phosphate Ammonium carbonate Sodium sulfite Problem 4: Write the names for the following: NaNO3 K2SO4 FePO4

19 Problem 3: Write the ionic formulas for each of the following:
Magnesium phosphate Mg2+ with PO43– Mg3(PO4)2 b) Ammonium carbonate NH4+ with CO32– (NH4)2CO3 c) Sodium sulfite Na+ with SO32– Na2SO3

20 Problem 4: Write the names for the following:
a) NaNO3 sodium nitrate b) K2SO4 potassium sulfate c) FePO4 iron (III) phosphate

21 Covalent Compounds Two or more nonmetals typically bond in covalent bonds Sharing of electrons (not always equal sharing) Molecular formula: actual atoms in molecule (rather than ratio) C6H12O6: glucose CH2O: formaldehyde Molecular formula correlates to molar ratio

22 Molecular Formula H2O 2 H’s and 1 O per molecule 500 H2O molecules:
How many H’s? 1000 H atoms How many O’s? 500 O atoms Problem 5 How many Cl atoms are in 40 CCl4 molecules? How many C atoms are in 30 C6H12O6 molecules? How many H atoms? How many O atoms?

23 Problem 5 CCl4 40 molecules CCl4 = 160 Cl atoms
4 x 40 = 1 x 160 Cl atoms

24 Problem 5 b) C6H12O6 30 molecules C6H12O6 = 180 C atoms
30 molecules C6H12O6 have = 360 H atoms 30 molecules C6H12O6 have = 180 O atoms

25 Problem 6 How many moles of H are in 5.0 moles of BH3?
How many moles of Cl are in 12.0 moles of C2H4Cl2?

26 Problem 6 How many moles of H are in 5.0 moles of BH3?
How many moles of Cl are in 12.0 moles of C2H4Cl2? a) 5.0 mol BH3 = 15.0 mol H b) mol C2H4Cl2 = 24.0 mol Cl

27 Covalent Bonding Sharing of valence electrons Non-metals Many models
Lewis structures: number and types of bonds VSEPR (valence shell electron pair repulsion): empirical model that predicts molecular geometry Valence Bond model: describes nature of bonds and predicts reactivity Molecular Orbital theory: gold standard of understanding bonding, but requires high level of mathematics (calculus, group theory)

28 Bonding Why do bonds form? A + B vs. A—B lower E higher E NO BOND
higher E lower E BOND

29 Bonding Why do bonds form? A + B vs. A—B lower E higher E NO BOND
He + Cl vs. He—Cl

30 Bonding Why do bonds form? A + B vs. A—B lower E higher E NO BOND
He + Cl vs. He—Cl

31 Bonding Why do bonds form? A + B vs. A—B higher E lower E BOND
H + Cl vs. H—Cl preferred

32 Valence electrons 1e– 2e– 3e– 4e– 5e– 6e– 7e– 8e– H He Li Be B C N O F
Ne Na Mg Al Si P S Cl Ar

33 Lewis “dot” structures
F 7 valence electrons F– 8 valence electrons F Ne F N 5 valence electrons N3– 8 valence electrons N N Ne Put electrons in singly b/f pairing: C N

34 Problem 7 How many valence electrons would you expect each of the following to have? Draw a Lewis structure for each one. Sr Se I (iodine) K Cs

35 Problem 7 a) Sr: 2 valence electrons b) Se: 6 valence electrons
I c) I: 7 valence electrons d) K: 1 valence electron K e) Cs: 1 valence electron Cs

36 Problem 8 Draw Lewis dot structures for C P P3– Se2–

37 Problem 8 C a) C: 4 valence electrons P b) P: 5 valence electrons
c) P3–: 8 valence electrons P d) Se2-: 8 valence electrons Se

38 Covalent Bonding Sharing of electrons H O H O H H O H H2O

39 Covalent Bonding Sharing of electrons H O H like Ne

40 Covalent Bonding Sharing of electrons H O H like He

41 Covalent Bonding Nonmetal + nonmetal H C

42 Covalent Bonding Nonmetal + nonmetal H H C H CH4

43 Problem 9 Write the Lewis structure for each atom in the group of atoms given. Then figure out how they might bond together covalently to form a stable compound (octet/duet). C, Cl, F, Cl, H (central C atom) N, 2 H’s, Cl (central N atom) P and H (use as many H’s as you need) 2 C’s and 6 H’s O, F, H (central O)

44 Problem 9 C Cl F H Cl F Cl C H Cl

45 Problem 9 b) N Cl H

46 Problem 9 b) N Cl H Cl N H Cl—N —H H H

47 Problem 9 c) P x H H P H H—P —H H H

48 Problem 9 d) C C H H H H C C H H–C –C –H H H H H

49 Problem 9 e) O F H F O H H–O –F

50 Multiple bonds C C H H H H Some atoms form double and triple bonds

51 Multiple bonds C2H4 H C C H H H

52 Multiple bonds C2H4 H C C H H H

53 Multiple bonds C2H4 H C C H H H

54 Multiple bonds C2H4 H C C H H–C C–H H H H H

55 Multiple bonds triple bonds C2H2 C C H H

56 Multiple bonds C2H2 H C C H

57 Multiple bonds C2H2 H C C H

58 Multiple bonds C2H2 H C C H

59 Multiple bonds C2H2 H C C H H – C  C – H

60 Covalent Bonds Atom… has… typically forms…
C, Si 4 valence electrons 4 bonds N, P 5 valence electrons 3 bonds O, S, Se 6 valence electrons 2 bonds F, Cl, Br, I 7 valence electrons 1 bond

61 Problem 10 Draw the Lewis structure for each of the atoms in the formula below. Then draw the Lewis structure for the molecule. HCN (central C) CH2O (central C) C2Cl2 HNO (central N)

62 Problem 10 a) H C N

63 Problem 10 a) H C N H C N

64 Problem 10 a) H C N H C N H–C N H C N

65 Problem 10 b) H H C O H H H C O H C O H–C=O H

66 Problem 10 c) Cl C C Cl Cl C C Cl Cl C C Cl Cl–C C–Cl

67 Problem 10 d) H N O

68 Problem 10 d) H N O H N O H N O H–N=O

69 Molecular Geometry (VSEPR)
Valence Shell Electron Pair Repulsion Electron bonding pairs repel each other Adopt geometry to maximize their separation distance. Treat multiple bonds as if they were single electron pair. 109.5° CH4 tetrahedral

70 Tetrahedral Carbon Anytime carbon is bonded to four other atoms, it has tetrahedral geometry. So H3C–CH2–CH2–CH2–CH2–CH2–CH2–CH2–CH2–CH2–CH2–CH3 is really

71 VSEPR CO2 O=C=O 2 electron pairs: linear
CO2 O=C=O   

72 VSEPR 3 electron pairs 3 bonding: trigonal planar H–B–H H link

73 4 electron pairs H H–C–H H H–N–H H H–O–H 4 bonding: tetrahedral
link H 4 bonding: tetrahedral H–C–H H 3 bonding (1 nonbonding): trigonal pyramid H–N–H H     2 bonding (2 nonbonding): “bent” or nonlinear H–O–H  

74 Problem 11 Given the following Lewis structures, predict the three dimensional molecular geometry. a) H–N–Cl b) Cl–S • • • • • • • • • • • • • • • • • • • Cl Cl • • • •

75 Problem 11 a) Trigonal pyramid N Cl H Cl b) Bent or Nonlinear S Cl Cl

76 Multiple Bonds in VSEPR
Treated like single bonds H–C–H O 120° Trigonal planar

77 Problem 12 Given the following Lewis structures, predict the three dimensional molecular geometry. Draw a picture of the molecule. a) H–CN b) O=N–H c) H–C=C –H • • • • • • H H

78 Problem 12 a) H–CN Linear ·· b) O=N–H ·· O H

79 Problem 12 c) H–C=C–H H H

80 2 AX2 3 AX3 1 4 AX4 e– regions Formula B NB e – geom Molec. Geo.
Bond angles 2 AX2 linear 180º 3 AX3 Trig. planar 120º 1 Bent 4 AX4 Tetra-hedral 109.5º Trig. pyram.

81 Polar and Nonpolar Covalent Bonds
Equal sharing of electrons: C—C C C Ionic bonds: Na+ Cl– Complete transfer of electrons

82 Polar Covalent Bonds Unequal sharing of electrons H—Cl
Different electronegativity Cl is more electronegative than H Electronegativity: ability to attract electron density in a covalent bond

83 Polar Covalent Bonds Unequal sharing of electrons H—Cl H Cl H—Cl
Area of low electron density Area of high electron density + – H—Cl

84 Polar Covalent Bonds electronegativity electronegativity

85 Problem 13 Which of the following pairs of atoms is the more electronegative? P or F Sr or Si Se or Cs

86 Problem 13 Which of the following pairs of atoms is the more electronegative? P or F Sr or Si Se or Cs

87 Problem 14 Indicate the polarity of the following bonds using the + and ‑ symbolism.   a) CO b) NP c) BrF

88 Problem 14 Indicate the polarity of the following bonds using the + and ‑ symbolism.   + – – + + – b) NP a) CO c) BrF

89 Molecular Polarity O=C=O Polar bonds that cancel  nonpolar molecule
+ – Overall molecule: nonpolar Polar bonds that do NOT cancel  polar molecule + – H O Polar molecules dissolve other polar molecules (some ionic compounds Nonpolar molecules dissolve other nonpolar molecules

90 Chemical Reactions Chemical reactions occur when bonds break and reform in different arrangements Can absorb or release heat Overall energy needs to be “downhill” (more stable products)

91 Chemical Equations (Reactions)
No coefficient: assume it is 1 Designates the total number of H2 molecules that react N H2  2NH3 Ratio of N:H is 1:3 in product molecule reactants product(s) Total molecules of product formed: 2 Designates the atoms of N in each molecule Designates the atoms of H in each molecule

92 Chemical Equations CH4 + O2  CO2 + H2O CH4 + O2  CO2 + H2O reactants
products Number and type of atoms on each side must be equal (balance) CH4 + O2  CO2 + H2O 1 C 4 H 2 O 1 C 2 H 3 O

93 Chemical Equations CH4 + O2  CO2 + H2O CH4 + O2  CO2 + 2H2O
reactants products Number and type of atoms on each side must be equal (balance) CH4 + O2  CO2 + 2H2O 1 C 4 H 2 O 1 C 4 H 4 O

94 Chemical Equations CH4 + O2  CO2 + H2O CH4 + 2O2  CO2 + 2H2O
reactants products Number and type of atoms on each side must be equal (balance) CH4 + 2O2  CO2 + 2H2O 1 C 4 H 4 O 1 C 4 H 4 O balanced Can only change coefficients Cannot change molecular formula

95 Chemical Equations as Conversion Factors
2Fe + O2  2FeO If we start with 2.0 mol of Fe, how many moles of O2 do we need, and how many moles of FeO are produced? If we start with 4.0 mol of Fe, how many moles of O2 do we need, and how many moles of FeO are produced? 4.0 mol Fe = 2.0 mol O2 4.0 mol Fe = 4.0 mol FeO

96 The Mole Chemists use “moles” as a way to count atoms
Molar ratio corresponds to ratio of atoms and molecules in balanced equation Mole (and thus molecules) linked to laboratory by mass in grams and molar mass (molecular weight)

97 Gram to Mole Conversions
Grams  molar mass = moles Moles x molar mass = grams How many moles in 35.5 g of PCl3? 35.5 g Ratio method: = mol

98 Gram to Mole Conversions
Grams  molar mass = moles Moles x molar mass = grams How many grams is 6.50 mol of PCl3? 6.50 mol = g

99 Problem 15 Calculate the number of moles in 24.4 g of H2.
If the H2 reacts with oxygen to make hydrogen peroxide: H2 + O2  H2O2 How many moles of H2O2 will form? c) How many g of H2O2?

100 Problem 15 a) number of moles in 24.4 g of H2. 24.4 = 2.02 (x) x = 12.1 mol 24.4 g = 12.1 mol H2 b) H2 + O2  H2O2 How many moles of H2O2 will form? 1:1 ratio 12.1 mol H2O2

101 Problem 15 c) How many g of H2O2? 34.02 g/mol
12.08 mol H2O2 (2(1.01 g/mol) + 2(16.00 g/mol)) 34.02 g/mol 12.08 mol H2O2 (34.02 g/mol) = 411 g H2O2