Chemical Bonds Chapter 6 1
Ch. 6 Chemical Bonding I. How and Why Atoms Bond A. Vocabulary B. Chemical Bonds - Basics C. Chemical Bonds Types D. Chemical Bonds Covalent E. Drawing Lewis Diagrams F. Bond Polarity 2
A. Vocabulary (review) CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula NaCl CO 2
A. Vocabulary (review) COMPOUND 2 elements more than 2 elements Binary Compound Ternary Compound NaCl NaNO 3
A. Vocabulary (review) ION 1 atom 2 or more atoms Monatomic Ion Polyatomic Ion Na + NO 3 -
A. Vocabulary Chemical Bond Attractive force between atoms or ions that binds them together as a unit Bonds form in order to. Decrease potential energy Increase stability (fill valence shell according) 6
A. Vocabulary Chemical Formula The number and type of atoms in a compound The number of atoms in a molecular element ex. NaCl & CO 2 7
A. Vocabulary Compound A chemically bonded substance made up of more than one element BINARY COMPOUND: 2 elements NaCl TERTIARY COMPOUND: 3+ elements NaNO 3 8
A. Vocabulary Ion An atom or molecule with differing numbers of protons and electrons More protons = positive cation More electrons = negative anion Monatomic Ion: 1 atom Na + Polyatomic Ion: 2+ atoms NO 3-9
B. Chemical Bonds - Basic Most stable atoms have 8 valence electrons Octet rule Does not apply to d and f sublevels Atoms can gain, lose, or share electrons Generally: More than 4 valence electrons = gain electrons Less than 4 valence electrons = lose electrons 10
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B. Chemical Bonds - Basic Bonding depends on electronegativity (EN) 3 BASIC combinations 1. IONIC BOND - metals & nonmetals Highly EN nonmetals attract electrons Valence electrons of metal transfer to nonmetal Metal becomes cation, nonmetals becomes an anion Ionic bond results 12
B. Chemical Bonds - Basic 2. COVALENT BOND - nonmetals & nonmetals Large EN of both atoms equally attract electrons Neither gives up electrons and they are shared Covalent bond results 3. METALLIC BOND - metals & metals Neither metal strongly attracts electrons Electrons shared Metallic bond results THERE ARE MANY EXCEPTIONS TO THESE COMBINATIONS 13
Bond Formation Type of Structure Melting Point Solubility in Water Electrical Conductivity Other Properties C. Chemical Bonds - Types IONIC COVALENT e - are transferred from metal to nonmetal crystal lattice high yes yes (solution or liquid) e - are shared between two nonmetals true molecules Physical State solid liquid or gas low usually not no odorous
C. Chemical Bonds - Ionic Ionic Bonding - Crystal Lattice Lattice energy is the amount of energy needed to separate, or dissolve, a crystal Formula Units are the relative number of atoms in a compound. Simplest whole number ratio between atoms in a compound
NaCl crystals 16
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C. Chemical Bonds - Covalent Double covalent bonds atoms share 2 pairs of electrons Ex. Sulfur monoxide (S-O) Triple covalent bonds atoms share 3 pairs of electrons Ex. Atmospheric nitrogen (N N) Covalent bonding creates true molecules, individual units of bound substances 18
C. Chemical Bonds - Types Bond Formation Type of Structure Physical State Melting Point Solubility in Water Electrical Conductivity Other Properties METALLIC e - are delocalized among metal atoms electron sea solid very high no yes (any form) malleable, ductile, lustrous
D. Chemical Bonds Electron-Dot Structures Electron-Dot Structures (Lewis Dot Structures) Used to show how electrons are shared between atoms 20
D. Chemical Bonds Electron-Dot Structures Diagram of covalent bonds (-) Ex. CH 3 OH 21
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D. Chemical Bonds Electron-Dot Structures Resonance structures the different ways structures are arranged 23
D. Chemical Bonds Electron-Dot Structures Rules for drawing electron-dot structures Rule 6: Check to make sure all atoms have 8 electrons 24
E. Drawing Lewis Diagrams Examples on next slides: CF 4 BeCl 2 Standard Form CO 2 ClO 4 - NH 4 + SO 3 Polyatomic Form Resonance 25
Drawing Lewis Diagrams CF 4 1 C 4e - = 4e - 4 F 7e - = 28e - 32e - - 8e - 24e - F F C F F
Drawing Lewis Diagrams BeCl 2 1 Be 2e - = 2e - 2 Cl 7e - = 14e - 16e - - 4e - 12e - Cl Be Cl
Drawing Lewis Diagrams CO 2 1 C 4e - = 4e - 2 O 6e - = 12e - 16e - - 4e - 12e - O C O
Polyatomic Ions To find total # of valence e-: Add 1e- for each negative charge. Subtract 1e- for each positive charge. Place brackets around the ion and label the charge.
Polyatomic Ions ClO - 4 1 Cl 7e - = 7e - 4 O 6e - = 24e - 31e - + 1e - 32e - - 8e - 24e - O O Cl O O
Polyatomic Ions NH 4 + 1 N 5e - = 5e - 4 H 1e - = 4e - 9e - - 1e - 8e - - 8e - 0e - H H N H H
Resonance Structures Molecules that can t be correctly represented by a single Lewis diagram. Actual structure is an average of all the possibilities. Show possible structures separated by a double-headed arrow.
Resonance Structures SO 3 O O S O O O S O O O S O
F. Bond Polarity Most bonds are a blend of ionic and covalent characteristics Differences in electronegativity determines bond type. 34
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F. Bond Polarity Electronegativity trend Increases up and to the right Fluorine the most EN element 36
F. Bond Polarity Electronegativity Attraction an atom has for a shared pair of electrons Higher EN atom δ - Lower EN atom δ + 37
F. Bond Polarity Nonpolar covalent bond E- are shared equally Symmetrical e- density Usually identical atoms 38
F. Bond Polarity Polar covalent bond E- are shared unequally Asymmetrical e- density Results in partial charges (dipole) 39
F. Bond Polarity Ionic bond Electrons are pulled Ions are present Complete + and - charges 40
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F. Bond Polarity Nonpolar covalent no charge Polar covalent partial charge Which diagram represents a molecular with a nonpolar covalent bond? 42
Most bonds are a blend of ionic and covalent characteristics. Difference in electronegativity determines bond type. F. Bond Polarity
Ch. 6 Chemical Bonding II. Molecular Compounds A. Molecular Compounds vs. Formula Units B. Special Cases 44
A. Molecular Compounds vs. Formula Units 45
A. Special Cases Diatomic Molecules H NO Halogens H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2 H and Halogens form single bonds O forms double bond N forms triple bond Polyatomic Ions A positive charge means lost electron A negative charge means gained electron Square brackets with superscript [ ] +/- Overall charge of the compound is NEUTRAL 46
Ch. 6 Chemical Bonding III. Quantum Model A. Valence Bond Theory B. Hybridization 47
A. Valence Bond Theory Valence Bond Theory covalent bonds are formed when orbitals of different atoms overlap to form an area of high electron probability H-H s orbital of each hydrogen overlaps Sigma Bonds!!! F-F p orbitals of each fluorine overlaps 48
A. Valence Bond Theory Sigma bonds = end-to-end bonds S and s bonds S and p bonds P and p bonds Additional Bonding = more than one set of orbitals overlap Pi bonds = side-to-side bonds Double bond 1 sigma bond + 1 pi bond Triple bond 1 sigma + 2 pi bonds 49
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B. Hybridization Hybridization the process by which new kinds of orbitals with equal energies are formed from a combination of orbitals of different energies: If 1S and 3P orbitals hybridize = sp3 If 1S and 2P orbitals hybridize = sp2 (1 left over p orbital forms 1 pi bond) Id 1S and 1P orbital hybridize = sp (2 left over p orbitals form 2 pi bonds) 53
C forming 4 bonds C C in CH 4 54
B forming 3 bonds 55
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Ch. 6 Chemical Bonding IV. Molecular Shape A. VSEPR Theory B. Determining Molecular Shape C. Common Molecular Shapes 57
A. VSEPR Theory Valence Shell Electron Pair Repulsion Theory Electron pairs orient themselves in order to minimize repulsive forces.
A. VSEPR Theory Types of e - Pairs Bonding pairs - form bonds Lone pairs - nonbonding e - Lone pairs repel more strongly than bonding pairs!!!
A. VSEPR Theory Lone pairs reduce the bond angle between atoms. Bond Angle
B. Determining Molecular Shape Draw the Lewis Diagram. Tally up e - pairs on central atom. double/triple bonds = ONE pair Shape is determined by the # of bonding pairs and lone pairs. Know the 8 common shapes & their bond angles!
C. Common Molecular Shapes 2 total 2 bond 0 lone LINEAR BeH 2 180
C. Common Molecular Shapes 3 total 3 bond 0 lone BF 3 TRIGONAL PLANAR 120
C. Common Molecular Shapes 3 total 2 bond 1 lone SO 2 BENT <120
C. Common Molecular Shapes 4 total 4 bond 0 lone CH 4 TETRAHEDRAL 109.5
C. Common Molecular Shapes 4 total 3 bond 1 lone NH 3 TRIGONAL PYRAMIDAL 107
C. Common Molecular Shapes 4 total 2 bond 2 lone H 2 O BENT 104.5
C. Common Molecular Shapes 5 total 5 bond 0 lone PCl 5 BIPYRAMIDAL TRIGONAL 120 /90
C. Common Molecular Shapes 6 total 6 bond 0 lone SF 6 OCTAHEDRAL 90
D. Examples PF 3 4 total 3 bond 1 lone F P F F TRIGONAL PYRAMIDAL 107
D. Examples CO 2 2 total 2 bond 0 lone O C O LINEAR 180
Ch. 6 Chemical Bonding V. Nomenclature A. Molecular Nomenclature B. Ionic Nomenclature C. Naming Acids 72
A. Molecular Nomenclature Binary compounds Prefix naming system Least EN atom comes first Add prefixes to indicate # of atoms Omit mono-prefix on first element Change the ending of the second element to -ide 73
A. Molecular Nomenclature Pre-fix Number Mono- 1 Di- 2 Tri- 3 Tetra- 4 Penta- 5 Hexa- 6 Hepta- 7 Octa- 8 Nona- 9 Deca- 10 74
A. Molecular Nomenclature Examples: CCl 4 N 2 O SF 4 H 2 O 75
B. Ionic Nomenclature Ionic Formulas: Write each ion, cation first. Don t show charges in the final formula Overall charge must equal zero If charges cancel, just write symbols If not, use subscripts to balance charges Use parentheses to show more than one polyatomic ion Stock system Roman numerals indicate the ion s charge. 76
B. Ionic Nomenclature Ionic Names: Write the names of both ions, cations first Change the ending of monotomic ions to ide Polyatomic ions have special names Stock Systems: Use Roman numerals to show the ion s charge if more than one is possible. Overall charge must equal zero. 77
B. Ionic Nomenclature Examples: Potassium chloride K + Cl - KCl Magnesium nitrate Mg 2+ NO 3- Mg(NO 3 ) 2 Copper (II) chloride Cu 2+ Cl - CuCl 2 78
C. Naming Acids See text 79