Unit 4: BONDING

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Unit 4BONDING

• Why do elements form bonds????

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I. Energy and Bonds

• Elements form bonds to become more stable
• Forming bonds releases energy
• Breaking bonds absorbs energy
• Therefore
• Forming a bond is
• Breaking a bond is

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II. Types of Bonds

• Bond attractive force that hold elements
  together
• There are 3 major types of bonds formed between
  elements
• Each type of bond has different attractions and
  different properties

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Identifying Bond Types

• A Metallic bonds present within a metal
• B Ionic bonds metal nonmetal
• Cations and anions form neutral substances
• Electrons are given/taken to form ions
• C Covalent nonmetals sharing electrons
• No actual charges formed

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III. Lewis Dot Structures

• Lewis dot diagrams show number and relative
  placement of valence electrons
• Uses element symbol and dots in pattern
• 1 2
• 8 3
• 5 6
• 7 4

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A. Single Elements

• Count number of valence electrons look at s
  and p electrons
• Place in pattern around element symbol
• Ex.

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B. Ions and Ionic Compounds

• Determine the charge of the ion within the
  compound look at oxidation numbers on PT
• Positive ions have NO valence electrons!
• Negative ions have 8 valence electrons!
• Arrange with opposite charges connecting
• Ex. NaCl MgCl2

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C. Covalent Structures

• Determine the number of valence electrons for
  each element involved
• Choose a central atom least popular
• Organize remaining atoms symmetrically
• Form bonds to provide each element with 8 valence
  electrons
• May use multiple bonds for each element to see 8
• Ex. H2 O2 N2

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D. The Octet Rule

• Octet eight valence electrons on an atom
• valence electrons are those in s and p
  sublevels
• Elements with 8 valence electrons are very stable
  and usually not reactive!
• What group has all 8 valence electrons naturally?
• Which group of metals is most reactive?
• Which nonmetals are most reactive?

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Octets, continued

• Since all elements want 8 electrons, each atom
  will gain or lose electrons to see 8 valence
  electrons
• Metals lose electrons
• nonmetals gain electrons
• Ex. NaCl MgO

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Exceptions to Octet Rule

• Some need less than 8
• H, He, B
• Some can take more than 8, creating an expanded
  octet
• S, P, etc.

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IV. Metallic Bonds

• Metallic Bonds special bonds between the atoms
  within a metal sample
• Have fixed nuclei with mobile electrons
• Sea of Mobile Electrons
• Give metals special properties
• Malleability -- Good Conductivity
• Ductility
• http//micro.magnet.fsu.edu/electromag/java/ruther
  ford/

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Diagram of Metallic Bonding
http//www.drkstreet.com/resources/metallic-bondin
g-animation.swf metallic bonding online demo
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V. Covalent Bonds

• Covalent bonds are formed between NONMETALS who
  share electrons
• Some nonmetals can form more than one bond
  between the same 2 elements
• Different types of covalent bonding, due to
  symmetry and electronegativity values
• No formal charge, or ions, formed

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Properties of Covalent Compounds

• All phases present at STP
• Low boiling and melting points
• Low density
• High vapor pressure, or volatility
• Poor conductors of heat and electricity

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a. Nonpolar Covalent Bonds

• Nonpolar equal sharing
• Electrons shared within a bond are seen equally
  by both atoms
• Between same atoms ONLY!

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b. Multiple Bonds

• 1Single Bonds 2 e- shared between 2 atoms
• 1 e- from each element
• Not very strong
• Longest covalent bond
• Examples
• some diatomics

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Multiple Bonds, cont.

• 2 Double Bonds 4 e- shared between 2 atoms
• 2 e- from each element
• Stronger and shorter than single bonds
• Examples
• O2

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Multiple bonds cont.

• 3 Triple bonds 6 e- shared between 2 atoms
• 3 e- from each element
• Strongest and shortest bond type
• Examples
• N2

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Online resources

• Online resources for understanding bonding
  geometries of common compounds
• https//phet.colorado.edu/en/simulation/molecule-
  shapes
• https//phet.colorado.edu/en/simulation/molecule-
  shapes-basics
• http//www.pbslearningmedia.org/asset/lsps07_int_
  covalentbond/

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c. Polar Bonds

• Polar bonds are formed between atoms having
  differences in EN
• Atoms of different EN will have different
  attractions for the bonding electrons
• The atom with the higher EN will have a stronger
  attraction for the bonding electrons

• Most polar bonds are polarized meaning that the
  electrons spend more time closer to the atom with
  the higher EN and less time near the atom of a
  lower EN

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Polar Bonds cont

• Molecules with polar bonds will have Dipoles
• Dipoles a charge imbalance within a bond
  created by different attractions for the bonding
  electrons

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d. Coordinate Covalent Bonds

• Coordinate covalent bonds
• bonds formed when only one element contributes
  electrons to the bond
• Only in special cases

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d. Network Covalent Bonds

• Network covalent bonds these are very strong
  bonds formed within a network solid between atoms
  of the same element or molecule
• Special cases

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VI. Molecular Structures

• Molecular shapes depend upon the distribution of
  electrons number of bonds formed
• Shapes are 3-Dimensional
• http//www2.chemistry.msu.edu/faculty/reusch/virtt
  xtjml/models.htmstart

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a. Nonpolar Molecules

• Nonpolar bonds are formed only between atoms
  having the same EN
• Only diatomic elements have true nonpolar bonds
• All bonding electrons are shared equally between
  atoms of the same EN

• Ex. Diatomic molecules

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Nonpolar, cont

• Even polar bonds can create nonpolar molecules
• Nonpolar molecules are SYMMERTICAL!
• Electrons are evenly distributed throughout the
  molecule, making it nonpolar!

• Symmetrical
• Nonpolar
• Ex. CF4

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b. Polar Molecules

• Polar Molecules have an asymmetric pull of
  electrons throughout the molecule
• Nonbonding electrons from lone pairs also create
  an asymmetric pull within the molecule

• Asymmetric
• Polar
• Ex. H2O

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Polar or Nonpolar Molecule???

• Examples
• a. CO2   
• b. OF2    
• c. CCl4    
• d. CH2Cl2   
• e. HCN

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c. Molecular Shapes, in 3D!

• Atoms are 3-dimensional substances that create
  3-D structures when bonding
• Both the bonds and the lone pair nonbinding
  electrons play a role in determining the shape of
  a molecule

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Bonding/Molecular Shape Terms

• Domain placement of electrons around an atom
• Bonding Domain includes all electrons
  participating in a bond counts as one area of
  space
• Nonbonding Domain space occupied by a lone pair
  of electrons nonbonding

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Additional secret informationthe VSEPR Theory

• VSEPR Valence Shell Electron Pair Repulsion
• This theory explains why the electrons within the
  bonds and the nonbonding electrons move as far
  apart as possible, creating a structure in
  3-dimensional space
• Nonbonding pairs sometimes have a greater effect
  than single bonds lets see!

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B. Shapes and Bond Angles

• http//intro.chem.okstate.edu/1314F97/Chapter9/VSE
  PR.html
• Or
• http//en.wikipedia.org/wiki/Molecular_geometry

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1. Linear

• 1 or 2 bonding domains
• 180o bond angle
• Symmetric if same elements, or distributed evenly
• Asymmetric if different atoms
• Examples Diatomics, CO2, HCl

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2. Trigonal Planar

• 3 bonding domains
• 120o bond angle
• Symmetric if all same elements
• Flat molecule!
• Examples BF3, SO3

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3. Trigonal Pyramidal

• 3 bonding domains, 1 nonbonding domain
• 107o bond angle
• Asymmetric due to lone pair electrons
• Examples NH3, PCl3

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4. Bent

• 2 bonding domains, 2 nonbonding domains
• 104.5o bond angle
• Asymmetric due to two lone pairs of electrons
• Examples H2O, SCl2

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5. Tetrahedron

• 4 bonding domains
• 109.5o bond angle
• Symmetric if all the same atoms bonded
• Asymmetric if different atoms
• Examples CH4, CCl4

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6. Trigonal Bipyramidal

• 5 bonding domains
• Expanded octet of 10 electrons
• 120o and 90o bond angle
• Symmetric if all the same atoms bonded
• Asymmetric if different atoms
• Examples PF5

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7. Octahedral

• 6 bonding domains
• Expanded octet of 12 electrons
• 90o bond angle
• Symmetric if all the same atoms bonded
• Asymmetric if different atoms
• Examples SF6

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VII. Ionic Bonds

• Ionic bonds a bond formed due to the transfer
  of electrons between metals and nonmetals
• Attractions bonds occur between ions charged
  atoms that have gained/lost electrons

• Cations positive ions have _______e-
• Metals form cations
• Anions negative ions have ______e-
• Nonmetals form anions

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Properties of Ionic Bonds

• High melting/boiling points
• Hard, but brittle crystals solids
• Dissolve in polar solvent
• Conduct electricity as liquid or in solution, but
  NOT as a solid

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Properties cont.

• Ionic substances have high heats of vaporization
• Low vapor pressure not very volatile
• Most dissolve in water to form () and (-) ions,
  or electrolytes

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A. Electronegativity Differences

• Large differences in EN Ionic Bonds
• When there is a larger difference in EN, the
  element with the higher EN will most likely to
  see the bonding electrons more, or share them
  less
• Ionic bonds have the greatest differences in EN!
• reinforced by the fact that one of the elements
  will actually TAKE the electrons instead of
  sharing them

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Covalent Bonds and EN

•   Even though nonmetals have relatively low EN in
  general, they do have slight differences
•  The only time there is no EN difference between
  atoms is for Diatomic elements
•  This means that the electrons in the bond(s)
  between the diatomic elements will be shared
  equally

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Rankings of EN DifferencesSee Figure 6-11, page
107, and figure 6-14, page 109
 
0 0.3 1.7 gt1.9 Diatomic
Nonpolar Polar Ionic Elements
Covalent Covalent Bonds
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Polar bonds in Molecules

• Arrows point to the element with the highest EN
• Use lower case Greek letter delta to represent
  partial charges d or d-
• Partially negative more Electronegative atom!

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b. Polyatomic Ions

• See Table E!!!
• Have BOTH covalent and ionic properties
• Covalent bonds hold the atoms together within the
  ion
• Overall, the structure has lost/gained electrons
  to have a charge
• Share electrons within, has brackets and charges
  for the Lewis Structure

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E. Resonance Structures

• Lewis dot structures with double-bond electrons
  that rotate from one pair to another
• Overall structure hybrid of all resonance
  structures

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Other Resonance structures

• NO3-1
• C6H6
• SO3-2
• CO3-2

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ReviewBond Strengths

• Network Covalent
• Ionic
• Covalent Triple Bond
• Covalent Double Bond
• Covalent Single Bond
• More Stable Molecules Stronger Bonds

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Larger EN Differences Stronger Bonds

• Stronger Bonds
• Equal
• More Stable Molecules

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VIII. Intermolecular Forces

• INTRAmolecular forces forces between atoms
• Bonds forces between the atoms
• INTERmolecular forces forces between molecules
• Four major variations
• Depends on the type of molecules or ions involved

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1 Molecule-Ion Attractions

• Definition
• Invisible force of attraction holding
  polar molecules and ions together in a solution
• Need polar molecules as solvent and ionic
  compound create () and (-) ions

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Molecule-ion forces

• The strongest of all the intermolecular forces!
• Positive ion attracted to partially negative end
  of polar molecule
• Negative ion attracted to partially positive end
  of polar molecule
• Orientation of polar molecules important!!!!
• Ex. Solution of NaCl(aq)

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2 Dipole-Dipole Attractions

• Definition
• Partially positive and partially negative
  ends of polar molecules develop attractive forces
• Need polar molecules as liquid

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Dipole-Dipole forces

• Occur within a sample of polar molecules
• Attraction occurs between partially positive ends
  of several of same polar molecules
• Partially positive end of the molecule near the
  atom with lower EN
• bonding electrons pulled away from it
• Partially negative end of moleculenear the atom
  with the highest EN
• pulls bonding electrons towards it
• Ex. HCl, HBr, HI

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3 Hydrogen Bonding

• Definition
• Special type of dipole-dipole forces occurring
  between polar molecules containing hydrogen and
  fluorine, oxygen, or nitrogen
• Ex. HF, H2O, and NH3

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Hydrogen Bonding

• Stronger than Dipole-Dipole, but weaker than
  actual bonds forming
• Hydrogen partially () strongly attracted to
  F, O, or N partially (-) end of molecule
• F, O, and N have high EN, small radii, and strong
  pull on bonding electrons
• Responsible for
• Abnormally high boiling point of water
• Larger volume of water in liquid phase

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4 Weak, London Dispersion, or Van der Waals
Forces

• Definition
• Weak attractive forces present between nonpolar
  molecules
• Need
• Nonpolar, symmetric molecules

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Weak, LD, or VdW forces

• Weakest attractive forces
• Created when nonpolar atoms/molecules have small,
  temporary dipoles formed via distribution of
  electrons
• Change as
• Distance between molecules increase, forces
  decrease
• Mass of molecules increase, forces increase

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Special effects of weak/LD/VdW forces

• Reason why diatomic elements of group 17 have
  increasing boiling points from top to bottom
• Remember phases of group 17
• gas, gas, liquid, solid, solid
• Cause hydrocarbons of fossil fuels to have
  increasing boiling points as their size and mass
  increase
• Methane is a gas, gasoline is a liquid, grease
  is a solid at the same temperatures

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Strengths of IMFsStrongest to Weakest

• 1 Molecule-ion
• 2 Hydrogen bonding
• 3 Dipole-Dipole
• 4 London Dispersion/Van der Waals/Weak