1. Chemical Bonding
Ch. 6

2. What is a chemical bond?
mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms together
Why don’t noble gases do this?
Already have filled s and p orbitals
stable octet: 8 valence e- (or 2 if, you’re helium)
Atoms that don’t have a stable octet are more reactive

3. Key Point #1: By forming bonds with each other, most atoms reduce their potential energy, becoming more stable.
This is a chemical change! All chemical changes involve energy!

4. What types of bonds can be formed?

5. Metallic Bonding
In a metal, the empty orbitals in the atoms’ outer energy levels overlap
Delocalized Electron: outer electron that does not belong to any one atom but can move freely through the metal’s network of empty atomic orbitals.
sea of electrons: mobile electrons around the metal atoms, which are packed together in a crystal lattice.

6. metallic bonding: chemical bonding that results from the attraction between metal atoms and surrounding sea of electrons
Key Point: In metallic bonding, valence electrons move freely throughout a network of metal atoms.

7. Metallic Bonding

8. Unique Characteristics of Metals
Metals have many unique properties because of their sea of electrons
Malleability: ability of a substance to be hammered or beaten into thin sheets
Ductility: ability of a substance to be pulled into a thin wire
Why? atoms can slide past one another along a plane without breaking bonds

9. Luster: shiny appearance
Why? Absorb a wide range of light frequencies, many orbitals separated by small energy differences
Conductivity
Thermal: ability to conduct heat
Electrical: ability to conduct electricity
Why? Electrons move easily through network of empty orbitals

10. Metallic Bonding Strength
The strength of metallic bonding is determined by the enthalpy of vaporization:
the amount of energy required to vaporize (turn into a gas) 1 mol of a metal
In general, the strength of the metallic bond INCREASES moving left to right across the periodic table.
Soft metals (less dense) metals  harder (more dense) metals toward right

11. Properties of Metals: Malleability and Ductility

12. Properties of Metals: Surface Appearance

13. Properties of Metals: Electrical and Thermal Conductivity

14. Types of Bonds What type of bonds can be formed? Ionic bond
Covalent bond
Nonpolar covalent
Polar covalent
Ionic bonding: bonds that result from electrical attractions between cations and anions
1 atom losses electrons
1 atom gains electrons
Covalent bonding: sharing of electrons between 2 or more atoms

16. Key Point 2: Rarely is bonding between atoms purely ionic or purely covalent.
Instead, it usually falls somewhere between the two extremes. Why?
Key Point 3: The extent of ionic or covalent bonding between two atoms can be estimated by calculating the difference in each elements’ electronegativity.

17. Covalent Bonding
Large difference in E.N.: bond has more ionic character
Small difference in E.N: bond has more covalent character
Types of Covalent Bonds
Non-polar covalent bonding: both electrons equally shared between atoms
Polar covalent bonding: unequal attraction for the shared electrons

18. 6.1 Practice Worksheet Part 1
The property of electronegativity, which is the measure of an atom’s ability to attract electrons, can be used to predict the degree to which the bonding between atoms of two elements is ionic or covalent.
The greater the electronegativity difference, the more ionic the bonding is.

19. If the calculated electronegative difference is…
> 1.7 : ionic bond is formed
> 0.3 , < 1.7 : polar-covalent bond
0 – 0.3 : non-polar covalent bond
Increasing difference in electronegativity
Nonpolar
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-

20. Electronegativity Difference
Elements
Electronegativity
Electronegativity Difference
Bond Type
Element 1
Element 2
Mg to Cl
H to O
C to Cl
N to H
C to S
K to F
Na to Cl
H to H
3.0
1.8
Ionic
1.2
2.1
3.5
1.4
Polar covalent
2.5
3.0 .5
Polar covalent
3.0
2.1 .9
Polar covalent
2.5
2.5
Non-polar covalent .8
4.0
3.2
Ionic .9
3.0
2.1
Ionic
2.1
2.1
Non-polar covalent

21. Polyatomic Ions
It is also possible if a compound contains polyatomic ions, for both types of bonding to be present.
Monatomic Ions: Fe2+ , Na+, Cl-
Polyatomic Ions: PO43-, NH4+ , NO-1
Groups of atoms are bonded covalent together, but because of few or more than expected valence electrons they have an overall charge (so they can also bond ionically with other ions)
Ex: Ca2+ and SO42-  CaSO4 (metal & diff. nonmetals)

22. Classify the following as ionic, covalent, or both
1. CaCl2 = __________
5. BaSO4 = ___________
(metal & nonmetal)
(metal & diff. nonmetals)
2. CO2 = __________
6. H2O = ____________
(nonmetal & nonmetal)
3. MgO = __________
7. SO3 = ___________
4. HCl = ___________
8. AlPO4 = ___________
Covalent
Covalent
Ionic
Covalent
Covalent
Both

23. Section 6.2
Covalent Bonding

24. What is a molecule?
Neutral group of atoms that are held together by covalent bonds.
Chemical formula: indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.

25. Formation of Covalent Bonds
The electrons of one atom and protons of the other atom attract each another.
The two nuclei and two electrons repel each other.
These two forces cancel out to
form a covalent bond at a length
where the potential energy is
at a minimum.

27. Bond Length vs. Bond Energy
Bond length (pm): distance between two bonded atoms at their minimum potential energy
Bond energy (kJ/mol): energy required to break a chemical bond and form neutral isolated atoms.
Breaking bonds: absorbs (requires) energy
Forming bonds: releases energy
Key Point: As you increase the number of bonds between 2 atoms the bond energy increases, while the bond length decreases. This is an inverse relationship.

28. Bond Energies & Bond Lengths
A. How many electrons are shared in a
single bond:
double bond:
triple bond:
B. Which bond is shorter? C – C or C = C
C. Which bond requires more energy to break?
In addition to finding an ideal bond length, atoms also lower their potential energy by achieving a stable octet of 8 valence electrons

29. Bond Energies & Bond Lengths
A. How many electrons are shared in a
single bond: 2 e-
double bond: 4 e-
triple bond: 6e-
B. Which bond is shorter? C – C or C = C
C. Which bond requires more energy to break? =
In addition to finding an ideal bond length, atoms also lower their potential energy by achieving a stable octet of 8 valence electrons

30. Octet Rule Hydrogen: 2e- Boron: 6e-
Octet Rule: chemical compounds tend to form so that each atom has an octet of e-’s in its highest occupied energy level
Exceptions to the octet rule:
Atoms that cannot fit eight electrons
Atoms that can fit more than eight electrons
Hydrogen: 2e-
Boron: 6e-
Phosphorus, Sulfur, & Xenon: expanded valence, more than 8e-

31. How can we represent molecules?
Lewis Structures: formulas in which atomic symbols represent nuclei and inner shells, which are surrounded by dot-pairs/dashes represent valence electrons

32. Chapter 6.5
Molecular Geometry

33. VSEPR THEORY Lewis Structures are 2D but we live in a 3D world!
molecular geometry: the three-dimensional arrangement of a molecule’s atoms
What do those 3D structure/shapes look like??
Follow the Valance Shell Electron Pair Repulsion Theory or VSEPR
Repulsion between the sets of valence electrons surrounding an atom causes them to be oriented as far away from each other as possible

34. Why use VSEPR Theory?
Key Point: VSERP Theory is used to predict the shape of molecules based on the fact that electron pairs strongly repel each other.
Following VSEPR allows us to predict bond polarity:
uneven distribution of electrons

35. AB2 – Linear Central atom
Atoms/group of atoms attached to central atom
Atoms bonded to central atom (B)
Number of Lone Pairs on central atom (E)
Bond Angle 2
180˚

36. Other Linear Geometries
The shape of two atoms bonded together is not given in the chart.
Ex: F2
What is the only possible shape a binary compound can have?
LINEAR!

37. Atoms bonded to central atom (B) Number of Lone Pairs (E)
AB2E1 – Bent
Atoms bonded to central atom (B)
Number of Lone Pairs (E)
Bond Angle 2 1
<120˚
What happens to the bond angle between atoms as you increase the number of “lone pair electrons” on the central atom?
Bond angles decrease!

38. Atoms bonded to central atom (B) Number of Lone Pairs (E)
AB2E2 – Bent
Atoms bonded to central atom (B)
Number of Lone Pairs (E)
Bond Angle 2
104.5˚

39. Atoms bonded to central atom (B) Number of Lone Pairs (E)
AB3 – Trigonal Planar
Atoms bonded to central atom (B)
Number of Lone Pairs (E)
Bond Angle 3
120˚
Shape is often associated with atoms that break octet rule, but doesn’t have to be

40. AB3E1 – Trigonal Pyramidal
Atoms bonded to central atom (B)
Number of Lone Pairs (E)
Bond Angle 3 1
107˚

41. Atoms bonded to central atom (B) Number of Lone Pairs (E)
AB4 – Tetrahedral
Atoms bonded to central atom (B)
Number of Lone Pairs (E)
Bond Angle 4
109.5˚

42. Predicting Molecular Geometry
Draw Lewis structure for molecule.
VSEPR theory: if any lone pairs of electrons are found on the central atom, these electrons decrease the bond angles of atoms attached to it.
2. Draw a revised Lewis structure to show more accurate geometry S F S O
AB4E
AB2E
tetrahedral
bent

43. Predicting Molecular Polarity
3. To indicate the polarity of the bonds, we use this symbol: __________________ , which always points toward the more electronegative element.
4. When multiple bonds are found in a molecule, we must identify polarity of each bond.
electron rich
region
electron poor
region H F
e- poor
e- rich F H d+ d-

44. 5. Observe the overall polarity of the molecule
5. Observe the overall polarity of the molecule. Think of it as “tug-of-war” for valence electrons between the various atoms.
Non-polar covalent molecules: If the atom is symmetrical and all atoms have an equal pull on electrons
Polar covalent molecules: If the atom is not symmetrical and/or the atoms do not all have an equal chance of winning the tug of war for electrons

45. Intermolecular Forces

46. Intermolecular forces: attractive forces between molecules.
Intramolecular forces: attractive forces within a molecule (the bonds)
Intermolecular Forces
Intramolecular Forces (bond)
Intramolecular Forces
intermolecular forces are much weaker than intramolecular forces

47. Strength of IMF Hydrogen Bond Dipole – Dipole Induced Dipole
London Dispersion Forces
strongest
weakest

48. d+ d- Dipoles What is a dipole? A polar molecule
Uneven sharing of electrons so there is a separation of charge H F
electron poor
region
electron rich
region
e- poor F H d+ d-

49. Dipole-Dipole Forces
Attraction between two polar molecules — — + +

50. Hydrogen Bonding Special type of Dipole – Dipole
Attraction between: Hydrogen & Nitrogen/Oxygen/Fluorine

51. Electrons shift toward positive end of dipole
Induced Dipole
Attraction between one polar and one nonpolar molecule — +
Electrons shift toward positive end of dipole — — + +

52. London Dispersion Forces
Attraction between two nonpolar molecules
Electrons become uneven and form a dipole — — + +

53. What does IMF effect?
Viscosity
Surface Tension
Boiling Point

54. Stronger IMF  Higher Boiling Point
Point at which liquid particles escape the surface of the liquid into the gas phase
Stronger IMF  Higher Boiling Point

55. Stronger IMF  Higher Surface Tension
result of an imbalance of forces at the surface of a liquid.
Stronger IMF  Higher Surface Tension

57. Stronger IMF  Higher Viscosity
Measures a fluid’s resistance to flow
Stronger IMF  Higher Viscosity