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1 Copyright© by Houghton Mifflin Company. All rights reserved.
Chemical Bonding Copyright© by Houghton Mifflin Company. All rights reserved.

2 Intramolecular Bonding Intermolecular Bonding
Within molecules Intermolecular Bonding Between molecule Copyright© by Houghton Mifflin Company. All rights reserved.

3 Octet Rule Atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons A very important rule An Octet consist of eight electrons This refers to the outermost eight electrons Compounds form to satisfy the Octet Rule Ionic or Covalent Bonds

4 Types of IntramolecularBonds
Ionic Formed by transfer of electrons Held together by electrostatic attraction Covalent Formed by sharing of electrons Held together by shared electrons Metallic Formed when electron(s) become detached from metal atoms Sea of electrons among + ions metal of atoms


6 Electronegativity values for selected elements.

7 Electronegativity Difference and Bond Type
(Approximate) Bond Type Example 0-0.4 Nonpolar Covalent H-H Moderately Polar Covalent H-Cl 1.0-2 Very Polar Covalent H-F ≥2.0 Ionic NaCl

8 Ions as packed spheres.

9 Na+1 + Cl-1  NaCl (table salt)
Ionic Bonds Ionic bonding occurs when one atom transfers an electron to another atom Both atoms involved become charged One is negatively charged One is positively charged This occurs when a metal reacts with a nonmetal Na+1 + Cl-1  NaCl (table salt)


11 Ions and Charges Cations Positive Ions Anions Negative Ions
Group 1, 2, 3 Form ions with charges equal to their group number Group 5, 6, 7 Form ions with charges equal to group number minus eight

12 Ionic Compounds and Formulas
The formula of a compound describes what elements are in the compound and in what proportions. Compounds that are held together by ionic bonds are called ionic compounds.

13 Common polyatomic ion names
Formula Name NH4+ Ammonium ion CO32- Carbonate ion PO43- Phosphate ion SO42- Sulfate ion OH- Hydroxide ion NO3- Nitrate ion

14 Formation of Ionic Compounds
Al SO42- Al2(SO4)3 Ionic compounds forms such that total of ionic charges is zero. For above case: 2 x (+3) = 6 3 x (-2) = -6 Total

15 O The Lewis dot structure for Oxygen
Oxygen is in group VIA so it has 6 valence electrons

16 Cl The Lewis dot structure for Chlorine
chlorine is in group VIIA so it has 7 valence electrons

17 Ca The Lewis dot structure for calcium
calcium is in group IIA so it has 2 valence electrons

18 Making calcium chloride
+ Ca Cl CaCl2

19 Covalent Bonds A covalent bond is a chemical bond that is formed when two atoms share a pair of electrons. H. + H.  H:H Covalent Compounds and Formulas In the above example, each hydrogen has a filled valence shell simulating the electron configuration of helium. Compounds that are held together by covalent bonds are called covalent compounds. Covalent compounds form from atoms on the right side of the periodic table

20 Multiple Bonds. In electron dot notations, a pair of electrons can be represented by a pair of dots : . This can be a bonding pair or a lone pair (non-bonding pair). Bonding pairs can also be represented by lines connecting atoms. H:H = H-H When one pair of electrons is shared, it is called a single bond. H-H

21 When two pairs of electrons are shared it is called a double bond.
When three pairs of electrons are shared it is called a triple bond.

22 Bond Length Bond Length varies for covalent bonds between different elements See p. 187 Two important trends As one moves down a group, bond length between atoms increases F—F nm Cl—Cl nm Br—Br nm Multiple bonds are shorter than single bonds Example: two carbon atoms bonded together Single covalent bond nm Double covalent bond nm Triple covalent bond nm


24 Exceptions to the Octet Rule
Exceptions to the octet rule include those for atoms that cannot fit eight electrons, and for those that can fit more than eight electrons, into their outermost orbital. Hydrogen forms bonds in which it is surrounded by only two electrons. Boron has just three valence electrons, so it tends to form bonds in which it is surrounded by six electrons. Main-group elements in Periods 3 and up can form bonds with expanded valence, involving more than eight electrons, e.g. PF5 and SF6

25 Electron-Dot Notation
To keep track of valence electrons, it is helpful to use electron-dot notation. Electron-dot notation is an electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol. The inner-shell electrons are not shown.

26 Lewis Structures Chapter 6
Electron-dot notation can also be used to represent molecules. The pair of dots between the two symbols represents the shared electron pair of the hydrogen-hydrogen covalent bond. For a molecule of fluorine, F2, the electron-dot notations of two fluorine atoms are combined.

27 Chapter 6 Lewis Structures
The pair of dots between the two symbols represents the shared pair of a covalent bond. In addition, each fluorine atom is surrounded by three pairs of electrons that are not shared in bonds. An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom. See Example, p. 185

28 Multiple Covalent Bonds
Chapter 6 Multiple Covalent Bonds Double and triple bonds are referred to as multiple bonds, or multiple covalent bonds. (See Table 2, p. 187) In general, double bonds have greater bond energies and are shorter than single bonds. Triple bonds are even stronger and shorter than double bonds. When writing Lewis structures for molecules that contain carbon, nitrogen, or oxygen, remember that multiple bonds between pairs of these atoms are possible.

29 Molecular Shape Arises because electrons repulse one another
Called VSEPR Valence Shell Electron Pair Repulsion VSEPR states that in a small molecule, the pairs of valence electrons are arranged as far apart from each other as possible

30 Summary of Molecular Shapes
Type Bond Angle Unshared Pairs Example Linear 180° Balanced (2 each side) CO2 Trigonal Planar 120° none BCl3 Tetrahedral 109.5° CH4 Pyramidal 107° one NH3 Bent 105° two H2O

31 Bent molecular structure of the water molecule.

32 A closer look at a water molecule (bent)

33 Molecular structure of methane (Tetrahedral).

34 Pyramidal

35 8.3 Molecular Orbitals When two atoms combine, the molecular orbital model assumes that their atomic orbitals overlap to produce molecular orbitals, or orbitals that apply to the entire molecule. Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole. A molecular orbital that can be occupied by two electrons of a covalent bond is called a bonding orbital.

36 Molecular Orbitals 8.3 Sigma Bonds
When two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei, a sigma bond is formed.

37 8.3 Molecular Orbitals When two fluorine atoms combine, the p orbitals overlap to produce a bonding molecular orbital. The F—F bond is a sigma bond. Two p atomic orbitals can combine to form a sigma-bonding molecular orbital, as in the case of fluorine (F2). Notice that the sigma bond is symmetrical around the bond axis connecting the nuclei.

38 Molecular Orbitals 8.3 Pi Bonds
In a pi bond (symbolized by the Greek letter ), the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms.

39 Polarity Bonds can be polar or nonpolar
Depends on electronegativity difference Molecules can be polar or nonpolar Polar molecules are called dipoles One end of a polar molecule has “+” charge; other end has “-” charge Will align in electric field Will be attracted or deflected by magnetic field Polarity of a molecule is determined by shape and the type of bonds between its atoms

40 The chlorine atom attracts the electron cloud more than the hydrogen atom does. Why?

41 Polar Molecules 8.4 A hydrogen chloride molecule is a dipole.
When polar molecules, such as HCl, are placed in an electric field, the slightly negative ends of the molecules become oriented toward the positively charged plate and the slightly positive ends of the molecules become oriented toward the negatively charged plate. Predicting What would happen if, instead, carbon dioxide molecules were placed between the plates? Why?

42 Intermolecular Forces
Attraction between molecules Holds groups of molecules together Intermolecular forces are known as van der Waals forces Dipole-Dipole attraction Dispersion Forces Hydrogen Bonds

43 van der Waals Forces Dipole-Dipole attraction Dispersion Forces
Electrostatic attraction between polar molecules Molecules line up like magnets Dispersion Forces Polarity arises due to electron imbalance Weak electrostatic forces arise Hydrogen Bonds Strong dipoles arising from covalent bonds of hydrogen atom to a very electronegative atom H electronegativity = 2.1 F=4.0 Results in liquids with high boiling points

44 Probability representations of the electron sharing in HF.

45 Charge distribution in the water molecule.

46 Water molecule behaves as if it had a positive and negative end.

47 Polarity Formaldehyde CH2O Carbon Dioxide CO2 Water H2O
Forms dipole due to high electronegativity of O Carbon Dioxide CO2 Has polar bond but they cancel out CO2 is nonpolar Water H2O Water is a bent molecule due to unshared pairs unlike CO2 bond polarities do not cancel out Water is polar and forms dipole Why is water liquid and Carbon Dioxide a gas?

48 Large Molecules Examples: protein, Subunits linked together in a chain
Often bend and twist to form 3D shape Chains, rings, balls

49 The three states of water:


51 Table 12.3

52 Relative sizes of some ions and their parent atoms.

53 The three possible types of bonds.

54 Figure 12.1: The formation of a bond between two atoms.

55 Tetrahedral arrangement of electron pairs.

56 Tetrahedral arrangement of four electron pairs around oxygen.


58 Table 12.1

59 Figure 12.6: Polar water molecules are strongly attracted to negative ions by their positive ends.

60 Figure 12.6: Polar water molecules are strongly attracted to positive ions by their negative ends.

61 The NH3 molecule pyramidal structure.

62 Bond Energies and Bond Lengths for Single Bonds
Section 2 Covalent Bonding and Molecular Compounds Chapter 6 Bond Energies and Bond Lengths for Single Bonds

63 Hybrid Orbitals When atoms bond, their outer orbitals become distorted
Orbitals do not look like those of unbonded atom Hybrid orbitals are formed which are a cross between the bonding orbitals sp sp2 sp3 Orbital hybridization contributes to shape of molecule

64 Hybrid Orbitals

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