Chapter 6: Bonding or A Scientific Drama

Quick Review  What is a molecule?  Two or more atoms joined by bonds  What is a compound?  A molecule made up of more than one kind of atom  Binary compounds: 2 atoms  Ternary compounds: 3 atoms  But what is a bond?  What is a molecule?  Two or more atoms joined by bonds  What is a compound?  A molecule made up of more than one kind of atom  Binary compounds: 2 atoms  Ternary compounds: 3 atoms  But what is a bond? A ‘diatomic’ chlorine molecule Buckminsterfullerene

Bonds  Bonds hold molecules together  Different types of bonds give different properties  Bond: force of attraction between protons (nucleus) of one atom and electrons of another  Takes two electrons & two nuclei  Bonds hold molecules together  Different types of bonds give different properties  Bond: force of attraction between protons (nucleus) of one atom and electrons of another  Takes two electrons & two nuclei

Bond Formation  Electron orbitals overlap and hold two electrons in place  Bonds form when attraction > repulsion  Electron orbitals overlap and hold two electrons in place  Bonds form when attraction > repulsion

Making Bonds Bond formation is SPONTANEOUS :  System goes from high to low energy, by releasing energy  Creates stability Bond formation is SPONTANEOUS :  System goes from high to low energy, by releasing energy  Creates stability

Breaking Bonds Bond breaking is NOT SPONTANEOUS  System goes from low to high energy, needs to get that energy from somewhere else Bond breaking is NOT SPONTANEOUS  System goes from low to high energy, needs to get that energy from somewhere else

Bonds and Energy Levels  Kinetic and potential energies both decrease when bonds are formed  Kinetic and potential energies both increase when bonds are broken  KE is your temperature, PE reflects that  In comparisons, more energy lost means a more stable compound  Because energy out = energy in  Kinetic and potential energies both decrease when bonds are formed  Kinetic and potential energies both increase when bonds are broken  KE is your temperature, PE reflects that  In comparisons, more energy lost means a more stable compound  Because energy out = energy in

 Which kind of energy is stored in a chemical bond? 1) potential 2) kinetic 3) activation 4) ionization  As energy is released during the formation of a bond, the stability of the chemical system generally: 1) decreases 2) Increases 3) remains the same  Which kind of energy is stored in a chemical bond? 1) potential 2) kinetic 3) activation 4) ionization  As energy is released during the formation of a bond, the stability of the chemical system generally: 1) decreases 2) Increases 3) remains the same

Valence Electrons  Valence electrons are in the outermost energy level of an atom, so higher energy  Elements in same period have same # valence electrons, so behave similarly  Valence involved in bonding, others stay back and relax  Valence electrons are in the outermost energy level of an atom, so higher energy  Elements in same period have same # valence electrons, so behave similarly  Valence involved in bonding, others stay back and relax

Major League Valence Electrons steady, reliablea little less predictable wild, unpredictable

The Octet Rule  Noble gases are inert (inactive), very stable  8 valence electrons, complete octet  8 is most electrons that can be held in valence shell  All atoms want 8, and work to get there by interacting with others  Bonding is sharing or giving/receiving electrons  Stability found in complete octet  Noble gases are inert (inactive), very stable  8 valence electrons, complete octet  8 is most electrons that can be held in valence shell  All atoms want 8, and work to get there by interacting with others  Bonding is sharing or giving/receiving electrons  Stability found in complete octet

 Exceptions to octet rule:  H & He: only 2 electrons  B & Al: will try for 8 electrons, but fine with 6  N: easy-going, can have less or more than 8  F: really electronegative, can take more than 8e-, bonds with nonmetal in Period 3 or lower  Somewhat reactive noble gases (with F)  Exceptions to octet rule:  H & He: only 2 electrons  B & Al: will try for 8 electrons, but fine with 6  N: easy-going, can have less or more than 8  F: really electronegative, can take more than 8e-, bonds with nonmetal in Period 3 or lower  Somewhat reactive noble gases (with F)

Lewis Electron Dot Structures or Lewis Structures, or Lewis Diagrams  Show valence electrons for bonding  Chemical symbol surrounded by 1 to 8 “ VE ” dots  Symbol represents nucleus and non- VE, called a kernel  Show valence electrons for bonding  Chemical symbol surrounded by 1 to 8 “ VE ” dots  Symbol represents nucleus and non- VE, called a kernel  Two on top in S, then one at a time

Draw Lewis Structures for:  K  Mn  Sn  Al  At  Kr  C  Si  Pb Draw Lewis Structures for: KK  Mn  Sn  Al  At  Kr CC  Si  Pb

Important Clarification About Lewis Structures  When placing electrons dots around a chemical symbol, the first two always go on top  BUT those are simply the first 2 of however many valence electrons are present  They are counted with the valence electrons, not separate  If an atom has 1 VE, only draw 1 dot, not 2 on top and then 1  When placing electrons dots around a chemical symbol, the first two always go on top  BUT those are simply the first 2 of however many valence electrons are present  They are counted with the valence electrons, not separate  If an atom has 1 VE, only draw 1 dot, not 2 on top and then 1

3 Types of Bonding  Covalent: sharing of electrons  Ionic: giving and receiving of electrons  Metallic: distribution of electrons in metals  Covalent: sharing of electrons  Ionic: giving and receiving of electrons  Metallic: distribution of electrons in metals

Types of Bonds: Covalent Bonding

Covalent bonds  Covalent bonds: atoms share electrons to achieve a stable octet  2 nonmetals, same or different elements  Remember: 2 electrons per bond  8 electrons (octet) = up to 4 bonds can be formed (with exceptions)  Single/double/triple bonds  Extra electrons as lone pairs  Covalent bonds: atoms share electrons to achieve a stable octet  2 nonmetals, same or different elements  Remember: 2 electrons per bond  8 electrons (octet) = up to 4 bonds can be formed (with exceptions)  Single/double/triple bonds  Extra electrons as lone pairs

Know That Lines Are Bonds! 7 single bonds 1 double bond & 2 single bonds 10 single bonds Looks confusing, but same basic principle

- Reminder -  Electronegativity is a measure of the attraction of a nucleus for a bonded electron; how much an atom “wants” electrons

Rules for Lewis Diagrams of covalent bonds 1. Count total # valence electrons in atoms of compound 2. Arrange atoms Central usually has lowest electronegativity, only present once, and isn’t H 3. Place single bonds 4. Add lone pair electrons to outsides, then center 5. Make more bonds if needed, using lone pairs 1. Count total # valence electrons in atoms of compound 2. Arrange atoms Central usually has lowest electronegativity, only present once, and isn’t H 3. Place single bonds 4. Add lone pair electrons to outsides, then center 5. Make more bonds if needed, using lone pairs

 Draw Lewis structures for: C 2 H 2 Water  N 2 Ammonia  Draw Lewis structures for: C 2 H 2 Water  N 2 Ammonia 1. Count total # valence electrons in atoms of compound 2. Arrange atoms Central usually has lowest electronegativity, only present once, and isn’t H 3. Place single bonds 4. Add lone pair electrons to outsides, then center 5. Make more bonds if needed, using lone pairs

Draw the Lewis structure for formaldehyde, CH 2 O 1. Count total # valence electrons in atoms of compound 2. Arrange atoms Central usually has lowest electronegativity, only present once, and isn’t H 3. Place single bonds 4. Add lone pair electrons to outsides, then center 5. Make more bonds if needed, using lone pairs 1. Count total # valence electrons in atoms of compound 2. Arrange atoms Central usually has lowest electronegativity, only present once, and isn’t H 3. Place single bonds 4. Add lone pair electrons to outsides, then center 5. Make more bonds if needed, using lone pairs

Draw the Lewis structure for hypobromous acid, HOBr 1. Count total # valence electrons in atoms of compound 2. Arrange atoms Central usually has lowest electronegativity, only present once, and isn’t H 3. Place single bonds 4. Add lone pair electrons to outsides, then center 5. Make more bonds if needed, using lone pairs 1. Count total # valence electrons in atoms of compound 2. Arrange atoms Central usually has lowest electronegativity, only present once, and isn’t H 3. Place single bonds 4. Add lone pair electrons to outsides, then center 5. Make more bonds if needed, using lone pairs

Polarity of covalent bonds & The 1.7 “Rule”  Bond polarity is a measure of differences in electronegativity (EN) Nonpolar covalent: share electrons evenly, with same EN, like diatoms  Polar covalent: unequal sharing of electrons, different EN, bonded atoms become “+” and “-”  Bond polarity is a measure of differences in electronegativity (EN) Nonpolar covalent: share electrons evenly, with same EN, like diatoms  Polar covalent: unequal sharing of electrons, different EN, bonded atoms become “+” and “-” to ~ “RULE” NOT ALWAYS TRUE!

1.7 Rule

 Which of the following bonds is the most polar in nature? 1) Cl 2 2) HCl 3) HBr 4) HI  The bond in a diatomic nitrogen molecule (N 2 ) is best described as 1) polar 2) nonpolar ionic 3) nonpolar covalent 4) polar ionic  Which of the following bonds is the most polar in nature? 1) Cl 2 2) HCl 3) HBr 4) HI  The bond in a diatomic nitrogen molecule (N 2 ) is best described as 1) polar 2) nonpolar ionic 3) nonpolar covalent 4) polar ionic

Molecular Substances  Each atom of a molecular substance has the electron configuration of a noble gas  Can be solid/liquid/gas, depending on attractive forces  Properties associated with covalent bonding: generally soft, poor conductors of heat and electricity, low melting/boiling points  Each atom of a molecular substance has the electron configuration of a noble gas  Can be solid/liquid/gas, depending on attractive forces  Properties associated with covalent bonding: generally soft, poor conductors of heat and electricity, low melting/boiling points

 Molecules may contain polar bonds without being polar themselves C-O: 0.8C-Cl: 0.6O-H: 1.2H-Cl: 1.0N-H: 0.8 Symmetrical = nonpolar Asymmetrical = polar

 Which electron dot diagrams represent polar molecules?

Molecular Shapes  Linear  Bent  Pyramidal  Tetrahedral  Remember: lone pairs influence shape!  Symmetry = nonpolar  Linear  Bent  Pyramidal  Tetrahedral  Remember: lone pairs influence shape!  Symmetry = nonpolar

One Last Note About Covalent  Network solids are covalently bonded compounds without truly individual molecules  They go on forever in repeating patterns of atomic structure  Can be “divided” into unit cells, simplest representation of larger shape  Network solids are covalently bonded compounds without truly individual molecules  They go on forever in repeating patterns of atomic structure  Can be “divided” into unit cells, simplest representation of larger shape Unit cell… …of a net- work solid. NOT a net- work solid.

3 Types of Bonding  Covalent: sharing of electrons  Ionic: giving and receiving of electrons  Metallic: distribution of electrons in metals  Covalent: sharing of electrons  Ionic: giving and receiving of electrons  Metallic: distribution of electrons in metals

Types of Bonds: Ionic Bonding or “The Suspiciously Generous Stranger”

Ionic Bonding  Metal ion (+ cation) and a nonmetal ion (- anion)  Positive and negative ions held together by electrostatic attraction between opposite charges  Causes high melting/boiling points, hard/brittle substances, sometimes conductive  Metal ion (+ cation) and a nonmetal ion (- anion)  Positive and negative ions held together by electrostatic attraction between opposite charges  Causes high melting/boiling points, hard/brittle substances, sometimes conductive

Ions  Formed when individual atoms lose or gain electrons to be like the closest noble gas  Metals low EN, nonmetals high EN  Metals lose valence e- (to the nonmetal) and become cations (+)  Nonmetals gain valence e- (from the metal) and become anions (-)  Oxidation numbers represent charges formed, and the closest noble gas configuration  Formed when individual atoms lose or gain electrons to be like the closest noble gas  Metals low EN, nonmetals high EN  Metals lose valence e- (to the nonmetal) and become cations (+)  Nonmetals gain valence e- (from the metal) and become anions (-)  Oxidation numbers represent charges formed, and the closest noble gas configuration Electrons (-) charged, so Negative charge = gain electrons Positive charge = loss electrons

Reminder  Ionization energy: the amount of energy needed to remove the most loosely bound electron from a neutral atom -Sort of the inverse of electronegativity  Metals have low IE, want to lose their electrons, don’t put up a fight  Nonmetals have high IE, want all the electrons they can get, hold on for dear life  Ionization energy: the amount of energy needed to remove the most loosely bound electron from a neutral atom -Sort of the inverse of electronegativity  Metals have low IE, want to lose their electrons, don’t put up a fight  Nonmetals have high IE, want all the electrons they can get, hold on for dear life

Nonmetals gain e- & look like the closest noble gas to the right. Metals lose e- to look like the closest previous noble gas.

Lewis Structures for ions  Electron loss/gain based on electronegativity values  Nonmetals pull electrons away from metals  Ions go in brackets, in brackets, with charge with charge superscript superscript  Electron loss/gain based on electronegativity values  Nonmetals pull electrons away from metals  Ions go in brackets, in brackets, with charge with charge superscript superscript  May use different colors or symbols for “new” electrons

Lewis Structures for Ions 1. Count starting valence e- for each atom 2. Add or subtract electrons for ionic charge 3. Draw Lewis structures with new total electrons ∙ Metals will have none, nonmetals will have 8 · Draw nonmetals’ added e- differently than original 4. Put brackets around the ion 5. Write charge as a superscript 1. Count starting valence e- for each atom 2. Add or subtract electrons for ionic charge 3. Draw Lewis structures with new total electrons ∙ Metals will have none, nonmetals will have 8 · Draw nonmetals’ added e- differently than original 4. Put brackets around the ion 5. Write charge as a superscript Co Cr As Te Po O W Where do all the [M+] valence e- go?

Lewis Structures for Ionic Bonding 1. Count starting valence e- for each atom 2. Add or subtract electrons for ionic charge 3. Draw Lewis structures with new total electrons ∙ Metals will have none, nonmetals will have 8 · Draw nonmetals’ added e- differently than original 4. Put brackets around the ion 5. Write charge as a superscript 1. Count starting valence e- for each atom 2. Add or subtract electrons for ionic charge 3. Draw Lewis structures with new total electrons ∙ Metals will have none, nonmetals will have 8 · Draw nonmetals’ added e- differently than original 4. Put brackets around the ion 5. Write charge as a superscript NaCl MgO KBr PbI 2

Polyatomic Ions  Polyatomic ions have multiple atoms in a charged compound.  Receives e- from elsewhere to become stable  All individual atoms are not necessarily charged  Associate with other things by ionic bonding, but held together by covalent bonds  Polyatomic ions have multiple atoms in a charged compound.  Receives e- from elsewhere to become stable  All individual atoms are not necessarily charged  Associate with other things by ionic bonding, but held together by covalent bonds

Lewis Structures for Polyatomic Ions -Polyatomic ions are held together by covalent bonds, but interact with other things by ionic bonding Combined rules  Count all valence electrons, plus or minus any charge *Add for (-), subtract for (+)  Follow steps for covalent bonding  Add brackets and superscript like ionic bonds -Polyatomic ions are held together by covalent bonds, but interact with other things by ionic bonding Combined rules  Count all valence electrons, plus or minus any charge *Add for (-), subtract for (+)  Follow steps for covalent bonding  Add brackets and superscript like ionic bonds NO 3 PO NH 4 + ClF 2 + HSO 4

3 Types of Bonding  Ionic: giving and receiving of electrons  Covalent: sharing of electrons  Metallic: distribution of electrons in metals  Ionic: giving and receiving of electrons  Covalent: sharing of electrons  Metallic: distribution of electrons in metals

Metallic Bonds  Metals: few valence electrons and low ionization energies  Can mix, but don’t bond with other metals  Atoms fixed in crystalline lattice  Valence e- in a “sea of mobile electrons”  Metallic bonds: force of attraction of mobile VE for (+) atoms  Like metal ion Whack-a-Mole  Metals: few valence electrons and low ionization energies  Can mix, but don’t bond with other metals  Atoms fixed in crystalline lattice  Valence e- in a “sea of mobile electrons”  Metallic bonds: force of attraction of mobile VE for (+) atoms  Like metal ion Whack-a-Mole

Metallic Bond Properties  Good heat/electric conductivity  Bonds are strong: high melting and boiling points  Malleability: metals can be hammered into shapes  Atoms move to new positions, but electrons stay mobile Which element has a crystalline lattice throughout which electrons flow freely? 1) Bromine 2) Calcium 3) Carbon 4) Sulfur  Good heat/electric conductivity  Bonds are strong: high melting and boiling points  Malleability: metals can be hammered into shapes  Atoms move to new positions, but electrons stay mobile Which element has a crystalline lattice throughout which electrons flow freely? 1) Bromine 2) Calcium 3) Carbon 4) Sulfur

 Metallic bonding animation:  LzCH2KA  Metallic bonding animation:  LzCH2KA

Other Bonds  Savings bonds  Get stronger over time  James Bond’s  Vary in strength from one to the next  Barry Bonds  Also increase in strength, but  End up much stronger than should ever occur naturally  Savings bonds  Get stronger over time  James Bond’s  Vary in strength from one to the next  Barry Bonds  Also increase in strength, but  End up much stronger than should ever occur naturally

Distinguishing Between Bond Types

-Nonmetal & nonmetal-Metal & nonmetal -Polyatomic ions & anything

Intermolecular Forces  Act between molecules, not within them like “real bonds”  Only in covalent compounds, not ionic or metal  Strong IMF give high boiling and melting points  Electrostatic attractions exist between ionic compounds  Act between molecules, not within them like “real bonds”  Only in covalent compounds, not ionic or metal  Strong IMF give high boiling and melting points  Electrostatic attractions exist between ionic compounds

Dipoles  Polar molecules are dipoles (two distinct ends)  Opposite charges temporarily attract by dipole-dipole forces between covalent molecules  A dipole can “persuade” a nearby molecule to become an induced dipole  Dipole moment is the strength of attraction  Polar molecules are dipoles (two distinct ends)  Opposite charges temporarily attract by dipole-dipole forces between covalent molecules  A dipole can “persuade” a nearby molecule to become an induced dipole  Dipole moment is the strength of attraction

Hydrogen Bonding  Important enough to get its own slide  Hydrogen bonds act between an H atom and a nearby N, O, or F (3 very electronegative atoms)  Much stronger than dipole- dipole attractions  Hold water molecules together, give H 2 O its high boiling point (high for a small molecule)  Hydrogen in polar molecules is basically a bare proton, attracted by N/O/F  Important enough to get its own slide  Hydrogen bonds act between an H atom and a nearby N, O, or F (3 very electronegative atoms)  Much stronger than dipole- dipole attractions  Hold water molecules together, give H 2 O its high boiling point (high for a small molecule)  Hydrogen in polar molecules is basically a bare proton, attracted by N/O/F