chapter 6: bonding or a scientific drama. quick review what is a molecule? two or more atoms...
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Chapter 6: Bonding
orA Scientific Drama
Chapter 6: Bonding
orA Scientific Drama
Quick ReviewQuick 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
BondsBonds
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
BondFormatio
n
BondFormatio
n 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 BondsMaking 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 BondsBreaking 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 LevelsBonds 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 ElectronsValence 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
Major League Valence Electrons
steady, reliable a little less predictable
wild, unpredictable
The Octet RuleThe 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
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 Mn Sn Al At Kr C Si Pb
Important Clarification About Lewis StructuresImportant 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 Bonding3 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 BondingTypes of Bonds:
Covalent Bonding
Covalent bondsCovalent 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!
Know That Lines Are Bonds!
7 single bonds 1 double bond & 2 single bonds
10 single bonds
Looks confusing,but same basicprinciple
- Reminder -- Reminder -
Electronegativity is a measure of the attraction of a nucleus for a bonded electron; how much an atom “wants” electrons
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
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:• C2H2 •Water N2 •Ammonia
Draw Lewis structures for:• C2H2 •Water N2 •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, CH2O
Draw the Lewis structure for
formaldehyde, CH2O1. 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
Draw the Lewis structure for
hypobromous acid, HOBr1. 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”
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 “-”
0.0
0.1 to ~1.7-1.9
“RULE” NOT ALWAYS TRUE!
1.7 Rule1.7 Rule
Electronegativity BondDifference Polarity
0.0 nonpolar covalent0.1 - 1.7 polar covalent
1.8 and up ionic
Which of the following bonds is the most polar in nature?
1) Cl2 2) HCl 3) HBr 4) HI
The bond in a diatomic nitrogen molecule (N2) 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) Cl2 2) HCl 3) HBr 4) HI
The bond in a diatomic nitrogen molecule (N2) is best described as
1) polar 2) nonpolar ionic
3) nonpolar covalent 4) polar ionic
Molecular SubstancesMolecular 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
Molecules may contain polar bonds without being polar themselves
C-O: 0.8 C-Cl: 0.6 O-H: 1.2 H-Cl: 1.0 N-H: 0.8
Symmetrical = nonpolar Asymmetrical = polar
Which electron dot diagrams represent polar molecules?
Which electron dot diagrams represent polar molecules?
Molecular ShapesMolecular 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 CovalentOne 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.
NOTa net-worksolid.
3 Types of Bonding3 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”
Types of Bonds:Ionic Bonding
or“The Suspiciously
Generous Stranger”
Ionic BondingIonic 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
IonsIons 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, soNegative charge = gain electronsPositive charge = loss electrons
ReminderReminder
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 ionsLewis Structures for ions
Electron loss/gain based on electronegativity values
Nonmetals pull electrons away from metals
Ions go Ions go in brackets, in brackets, with charge with charge superscriptsuperscript
Electron loss/gain based on electronegativity values
Nonmetals pull electrons away from metals
Ions go Ions go in brackets, in brackets, with charge with charge superscriptsuperscript
May use different colorsor symbols for “new”electrons
Lewis Structures for IonsLewis 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 W2+ 3+ 3- 2- 4+ 2- 6+
Where do all the [M+]valence e- go?
Lewis Structures for Ionic Bonding
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 PbI2
Polyatomic IonsPolyatomic 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
Lewis Structures for Polyatomic Ions
-Polyatomic ions are held together by covalent bonds, but interact with other things by ionic bonding
Combined rules
1.Count all valence electrons, plus or minus any charge
*Add for (-), subtract for (+)
2.Follow steps for covalent
bonding
3.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
1.Count all valence electrons, plus or minus any charge
*Add for (-), subtract for (+)
2.Follow steps for covalent
bonding
3.Add brackets and
superscript like ionic bonds
NO3 PO4- 3- - NH4
+ClF2
+HSO4
1.7 +
0.0
0.1-1.6
3 Types of Bonding3 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 BondsMetallic 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 PropertiesMetallic 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:
http://www.youtube.com/watch?v=XHV9LzCH2KA
Metallic bonding animation:
http://www.youtube.com/watch?v=XHV9LzCH2KA
Other BondsOther 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 BetweenBond Types
Distinguishing BetweenBond Types
Distinguishing BetweenBond Types
Distinguishing BetweenBond Types
-Nonmetal & nonmetal-Metal & nonmetal-Polyatomic ions & anything
Intermolecular ForcesIntermolecular 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
DipolesDipoles
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 BondingHydrogen 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 H2O 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 H2O its high boiling point (high for a small molecule)
Hydrogen in polar molecules is basically a bare proton, attracted by N/O/F
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