0
Chapter 1
Atoms, Molecules & Chemical bonding

Introduction
Quantum Theory
Atomic Orbitals
Electronic Configuration

Prepa...
A Molecular Comparison of
Liquids and Solids
• Physical properties of substances understood in terms
of kinetic molecular ...
A Molecular Comparison of
Liquids and Solids

Prentice Hall © 2003

Chapter 11
A Molecular Comparison
of Liquids and Solids

Prentice Hall © 2003

Chapter 11
A Molecular Comparison
of Liquids and Solids
• Converting a gas into a liquid or solid requires the
molecules to get close...
Intermolecular Forces
• The covalent bond holding a molecule together
is an intramolecular forces.
• The attraction betwee...
Intermolecular Forces
Intermolecular Forces
Ion-Dipole Forces
• Interaction between an ion and a dipole (e.g. water).
• Strongest of all intermo...
Intermolecular Forces
Dipole-Dipole Forces
• Dipole-dipole forces exist
between
neutral
polar
molecules.
• Polar molecules...
Intermolecular Forces
Dipole-Dipole Forces
Intermolecular Forces
London Dispersion Forces
• Weakest of all intermolecular forces.
• It is possible for two adjacent n...
Intermolecular Forces
London Dispersion Forces
• One instantaneous dipole can induce another
instantaneous dipole in an ad...
Intermolecular Forces
London Dispersion Forces
• Polarizability is the ease with which an electron cloud can
be deformed.
...
Intermolecular Forces
London Dispersion Forces
Intermolecular Forces
Hydrogen Bonding
• Special case of dipole-dipole forces.
• By experiments: boiling points of compoun...
Hydrogen Bonding
Intermolecular Forces
Hydrogen Bonding
• H-bonding requires H bonded to an
electronegative element (most important for
com...
Hydrogen Bonding
Intermolecular Forces
Hydrogen Bonding
• Hydrogen bonds are responsible for:
– Ice Floating
•
•
•
•
•
•
•
•

Solids are us...
Intermolecular Forces
Hydrogen Bonding
Lewis Structures, Octet Rule,
Resonance & Formal Charge
 The number and arrangements of electrons in the outermost
shells...
The Octet Rule
“In

most compounds, the representative elements in the
compounds achieve noble gas configurations”
• This ...
How to write Lewis structures for
molecules & polyatomic ions?
• Step 1 : Count the total number of valence
electrons in t...
• Step 3 : Place the valence electrons until every
atom has a stable octet by:
i) drawing bond (either „ ‟ , „ ‟ or „ ‟)
a...
Now, let’s practise writing Lewis
structures!!!
Draw Lewis structures for:
1) CCl2F2
2) NF3
3) N2
4) CO32-
Exceptions to the Octet Rule
1) Electron deficient structures



Gaseous molecules containing either beryllium or boron ...
Exceptions to the Octet Rule
2) Odd – electrons molecules




A few molecules contain a central atom with odd number of ...
Exceptions to the Octet Rule
3) Expanded valence shells/octet




Molecules and ions have more than eight valence elect...
Resonance structures
• A molecule or polyatomic ion for which two or more Lewis
formulas with the same arrangements of ato...
Now, let’s practise!!!

• Draw resonance structures for CO32- and NO3– .

David P. White
Prentice Hall © 2003
Formal Charge
• Formal charge is the hypothetical charge on an
atom in a molecule or polyatomic ion.
• The concept of form...
Formal Charge
Formal charge of atom
totalnumber of valence e in free atom totalnumber of nonbonding e

1
( totalnumber of ...
David P. White
Prentice Hall © 2003
Now, let’s practise!!!

• Write formal charges for the carbonate ion, CO32-.
• Write formal charges for the ammonia molecu...
Week 3 intermolecular forces&lewis structure
Week 3 intermolecular forces&lewis structure
Week 3 intermolecular forces&lewis structure
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Transcript of "Week 3 intermolecular forces&lewis structure"

  1. 1. Chapter 1 Atoms, Molecules & Chemical bonding Introduction Quantum Theory Atomic Orbitals Electronic Configuration Prepared by: Mrs Faraziehan Senusi PA-A11-7C Molecular Orbitals Bonding and Intermolecular Compounds David P. White Prentice Hall © 2003
  2. 2. A Molecular Comparison of Liquids and Solids • Physical properties of substances understood in terms of kinetic molecular theory: – Gases are highly compressible, assumes shape and volume of container: • Gas molecules are far apart and do not interact much with each other. – Liquids are almost incompressible, assume the shape but not the volume of container: • Liquids molecules are held closer together than gas molecules, but not so rigidly that the molecules cannot slide past each other. – Solids are incompressible and have a definite shape and volume: • Solid molecules are packed closely together. The molecules are so rigidly packed that they cannot easily slide past each other.
  3. 3. A Molecular Comparison of Liquids and Solids Prentice Hall © 2003 Chapter 11
  4. 4. A Molecular Comparison of Liquids and Solids Prentice Hall © 2003 Chapter 11
  5. 5. A Molecular Comparison of Liquids and Solids • Converting a gas into a liquid or solid requires the molecules to get closer to each other: – cool or compress. • Converting a solid into a liquid or gas requires the molecules to move further apart: – heat or reduce pressure. • The forces holding solids and liquids together are called intermolecular forces.
  6. 6. Intermolecular Forces • The covalent bond holding a molecule together is an intramolecular forces. • The attraction between molecules is an intermolecular force. • Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl). • When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).
  7. 7. Intermolecular Forces
  8. 8. Intermolecular Forces Ion-Dipole Forces • Interaction between an ion and a dipole (e.g. water). • Strongest of all intermolecular forces. • ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increase. Prentice Hall © 2003 Chapter 11
  9. 9. Intermolecular Forces Dipole-Dipole Forces • Dipole-dipole forces exist between neutral polar molecules. • Polar molecules need to be close together. • Weaker than ion-dipole forces. • There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. • If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.
  10. 10. Intermolecular Forces Dipole-Dipole Forces
  11. 11. Intermolecular Forces London Dispersion Forces • Weakest of all intermolecular forces. • It is possible for two adjacent neutral molecules to affect each other. • The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). • For an instant, the electron clouds become distorted. • In that instant a dipole is formed (called an instantaneous dipole).
  12. 12. Intermolecular Forces London Dispersion Forces • One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom). • The forces between instantaneous dipoles are called London dispersion forces. Prentice Hall © 2003 Chapter 11
  13. 13. Intermolecular Forces London Dispersion Forces • Polarizability is the ease with which an electron cloud can be deformed. • The larger the molecule (the greater the number of electrons) the more polarizable. • London dispersion forces increase as molecular weight increases. • London dispersion forces exist between all molecules. • London dispersion forces depend on the shape of the molecule. • The greater the surface area available for contact, the greater the dispersion forces.
  14. 14. Intermolecular Forces London Dispersion Forces
  15. 15. Intermolecular Forces Hydrogen Bonding • Special case of dipole-dipole forces. • By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. • Intermolecular forces are abnormally strong.
  16. 16. Hydrogen Bonding
  17. 17. Intermolecular Forces Hydrogen Bonding • H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N). – Electrons in the H-X (X = electronegative element) lie much closer to X than H. – H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X. – Therefore, H-bonds are strong.
  18. 18. Hydrogen Bonding
  19. 19. Intermolecular Forces Hydrogen Bonding • Hydrogen bonds are responsible for: – Ice Floating • • • • • • • • Solids are usually more closely packed than liquids; Therefore, solids are more dense than liquids. Ice is ordered with an open structure to optimize H-bonding. Therefore, ice is less dense than water. In water the H-O bond length is 1.0 Å. The O…H hydrogen bond length is 1.8 Å. Ice has waters arranged in an open, regular hexagon. Each + H points towards a lone pair on O.
  20. 20. Intermolecular Forces Hydrogen Bonding
  21. 21. Lewis Structures, Octet Rule, Resonance & Formal Charge  The number and arrangements of electrons in the outermost shells of atoms determine the chemical and physical properties of the elements as well as the kinds of chemical bonds they form.  We write Lewis dot formulas (or Lewis dot representations, or just Lewis formulas) as a convenient bookkeeping method for keeping track of these “chemically important electrons.” David P. White Prentice Hall © 2003
  22. 22. The Octet Rule “In most compounds, the representative elements in the compounds achieve noble gas configurations” • This statement is usually called the Octet Rule, because the noble gas configurations have 8 e in their outermost shells (except for He, which has 2 e). • Many Lewis formulas are based on this idea
  23. 23. How to write Lewis structures for molecules & polyatomic ions? • Step 1 : Count the total number of valence electrons in the molecule  Doesn‟t matter which element it comes from, just the total of all elements. • Step 2 : Set up a skeleton structure  least electronegative element, or carbon in center.  H is always at the outskirts.  Most of the time (but not always) halogen will also be at the outskirts.
  24. 24. • Step 3 : Place the valence electrons until every atom has a stable octet by: i) drawing bond (either „ ‟ , „ ‟ or „ ‟) and/or ii) drawing lone pair (non-bonding pair of electrons) (:) **Make sure every atom has a stable octet, which means, every atom needs to have 8 valence electrons around it. Except for H of course, which satisfy its stability with 2 electrons. **Don‟t forget to show all lone pair electrons.
  25. 25. Now, let’s practise writing Lewis structures!!! Draw Lewis structures for: 1) CCl2F2 2) NF3 3) N2 4) CO32-
  26. 26. Exceptions to the Octet Rule 1) Electron deficient structures   Gaseous molecules containing either beryllium or boron as central atom ~ they have fewer than eight electrons around Be or B atom. Lewis structure for BeCl2 and BF3  Multiple bonds to the central atoms give unlikely structures:  Halogens are much more electronegative than Be or B.
  27. 27. Exceptions to the Octet Rule 2) Odd – electrons molecules   A few molecules contain a central atom with odd number of valence electrons ~ free radicals, contain a lone (unpaired) electron. Resonance for nitrogen oxide, NO2  But, the form with the lone electron on N (left) must be impotant because of the way NO2 react.  When two NO2 molecules collide, the lone electron pair up to form N–N bond in dinitrogen tetraoxide (N2O4) and each N attains an octet. David P. White Prentice Hall © 2003
  28. 28. Exceptions to the Octet Rule 3) Expanded valence shells/octet    Molecules and ions have more than eight valence electrons around the central atom. Expanded valence shells occur only with a central nonmetal atom in which d orbitals are available, one from period 3 or higher. Example: Sulfur hexafluoride, SF6. The central sulfur is surrounded by six single bonds, one to each fluorine for a total of 12 electrons. David P. White Prentice Hall © 2003
  29. 29. Resonance structures • A molecule or polyatomic ion for which two or more Lewis formulas with the same arrangements of atoms can be drawn to describe the bonding is said to exhibit resonance. • The three structures below are resonance structures of the carbonate ion: • The relationship among them is indicated by the doubleheaded arrows.The double-headed arrow indicates that the structures shown are resonance structures.
  30. 30. Now, let’s practise!!! • Draw resonance structures for CO32- and NO3– . David P. White Prentice Hall © 2003
  31. 31. Formal Charge • Formal charge is the hypothetical charge on an atom in a molecule or polyatomic ion. • The concept of formal charges helps us to write correct Lewis formulas in most cases. • The most energetically favorable formula for a molecule is usually one in which the formal charge on each atom is zero or as near zero as possible.
  32. 32. Formal Charge Formal charge of atom totalnumber of valence e in free atom totalnumber of nonbonding e 1 ( totalnumber of bonding e) 2 or number of valence e in free atom totalnumber of unshared valence e David P. White Prentice Hall © 2003 1 ( number of shared valence e) 2
  33. 33. David P. White Prentice Hall © 2003
  34. 34. Now, let’s practise!!! • Write formal charges for the carbonate ion, CO32-. • Write formal charges for the ammonia molecule, NH3 and to the ammonium ion, NH4+. David P. White Prentice Hall © 2003
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