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Chemical Bonding.

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Presentation on theme: "Chemical Bonding."— Presentation transcript:

1 Chemical Bonding

2 Chemical Bonding is all about the electrons.
in most of our discussions we will concentrate on the outer ‘s’ and ‘p’ electrons. These are called the valence electrons. since there are 4 of these orbitals in any quantum it requires 8 electrons (an octet) to fill them. the tendency of atoms to try to fill out the outer ‘s’ and ‘p’ shells is the octet rule.

3 Chemical Bonding Elements will combine in order to fill their valence shell with electrons, like the noble gases. This can happen in one of two ways: To share electrons - The outer orbitals of 2 atoms overlap so that each atom is in the vicinity of a full set of valence electrons. This type of bonding is called covalent bonding.


5 Chemical bonding Atoms gain or lose electrons to arrive at a full set of valence electrons. When atoms gain or lose electrons they become ions. Ions are attracted to ions of opposite charge and repelled by ions of the same charge. This type of bonding is called ionic bonding


7 How Many Bonds ??? the number of bonds made by atoms in the ‘s’ and ‘p’ blocks of the Periodic Table is determined by how many electrons they are away from an octet:

8 How Many Bonds ??? Group Valence Electrons Covalent Bonds Ionic Bonds
1 1 (1+) 2 21 2 (2+) 13 3 31 3 (3+) 14 4 41 - 15 5 3 (52) 3 (3-) 16 6 2 (62) 2 (2-) 17 7 1 (72) 1 (1-) 18 8 0 (82) 1electron promotion 2valence level expansion

9 most of our discussion will be centred on covalent bonding.

10 Valence level expansion
Some compounds occur which cannot be easily explained: PF5, SF6, ClF7, ArF8 in each case the number of chemical bonds is equal to the number of valence electrons. this can only happen if electrons are promoted to a higher energy level. In this case it is the adjacent ‘d’ orbital.

11 PF5 the normal orbital diagram looks like this:
with valence level expansion it looks like this:

12 Valence level expansion
includes elements of groups 15 to 18, from period 3 down; periods 1 and 2 do not have a ‘d’ orbital to promote to. Please note that valence level expansion is the exception, not the rule.

13 Polar-Covalent Bonding
Covalent bonding implies equal sharing of electrons. If sharing is not equal, the electrons in a bond will spend more time with one atom than the other.

14 Polar-Covalent Bonding
The atom where the electrons spend more time will have a net negative charge, while the atom at the other end of the bond will be positive. This type of bond is polar-covalent.

15 Electronegativity is a measure of how strongly an atom is holding on to its valence electrons. If an atom loses an electron fairly easily it has a low electronegativity (and tends to be a cation). If an atom tends not to lose electrons, but tends to steal them from other atoms (and become an anion) it has a high electronegativity.

16 To determine what type of bond exists between two atoms you subtract their respective electronegativities: if the electronegativity difference is 0.2 or less, the bond is covalent if the electronegativity difference is 1.7 or greater the bond is ionic.

17 Electronegativity If the electronegativity difference between two atoms is between 0.3 and 1.6 the bond is polar-covalent. The greater the electronegativity difference the greater the ionic character of the bond:


19 Assignment Determine the electronegativity difference for each chemical bond. If the bond is polar covalent draw an arrow in the direction of the dipole, from positive to negative: C - H END = | | = 0.4 polar covalent bond N - H B - F S - O P - H Si - Cl Cu - Br N - I Br - Cl O - H C - Cl C - O

20 N - H END = |3.0 – 2.1| = 0.9 polar covalent bond: N - H B - F END = |2.0 – 4.0| = 2.0 ionic bond: [B]3+[F]1- S – O END = |2.5 – 3.5| = 1.0 polar covalent bond: S - O

21 P - H END = |2.1 – 2.1| = 0.0 covalent bond: P – H Si - Cl END = |1.8 – 3.0| = 1.2 polar covalent bond: Si - Cl Cu - Br END = |1.9 – 2.8| = 0.9 polar covalent bond: Cu - Br

22 N – I END = |3.0 – 2.5| = 0.5 polar covalent bond: N - I Br – Cl END = |2.8 – 3.0| = 0.2 covalent bond: Br – Cl O – H END = |3.5 – 2.1| = 1.4 polar covalent bond: O – H

23 C – Cl END = |2.5 – 3.0| = 0.5 polar covalent bond: C – Cl C – O END = |2.5 – 3.5| = 1.0 polar covalent bond: C – O

24 Covalently-Bonded Structures
we now have to consider molecules made of several atoms. most of the following discussion will concern itself with molecules made with covalent or polar-covalent bonds. ionic bonds (and others) will return later in the unit.

25 Lewis Structures Lewis structures allow us to predict how atoms will come together to make molecules. Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

26 Writing Lewis Structures
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge. PCl3 (7) = 26

27 Writing Lewis Structures
The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26  6 = 20

28 Writing Lewis Structures
Fill the octets of the outer atoms. Keep track of the electrons: 26  6 = 20  18 = 2

29 Writing Lewis Structures
Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

30 Writing Lewis Structures
If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.

31 Exceptions to Octet Rule
Electron Promotion group 2 central atom will have 4 electrons group 13 central atom will have 6 electrons Valence Level Expansion group 15 central atom will have 10 electrons group 16 central atom will have 12 electrons group 17 central atom will have 14 electrons group 18 central atom will have 16 electrons

32 Writing Lewis Structures
Write Lewis Structures for the following molecules: F2 SiF4 PCl3 C3H8 SeH2 C3H6 MgH2 C3H4

33 Coordinate Covalent bonds
is defined as a covalent bond where one atom provides both of the electrons:

34 Polyatomic Ions we can use coordinate covalent bond theory to explain most ions:

35 Resonance This is the Lewis structure we would draw for ozone, O3. + -

36 Resonance But this is at odds with the true, observed structure of ozone, in which… …both O—O bonds are the same length. …both outer oxygens have a charge of 1/2.

37 Resonance One Lewis structure cannot accurately depict a molecule such as ozone. We use multiple structures, resonance structures, to describe the molecule.

38 Resonance Just as green is a synthesis of blue and yellow…
…ozone is a synthesis of these two resonance structures.

39 Resonance In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. They are not localized, but rather are delocalized.

40 Resonance The organic compound benzene, C6H6, has two resonance structures. It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

41 Molecular Shapes The shape of a molecule plays an important role in its reactivity. By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.

42 What Determines the Shape of a Molecule?
Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

43 Valence Shell Electron Pair Repulsion Theory (VSEPR)
“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

44 Electron-Domain Geometries
All one must do is count the number of electron domains in the Lewis structure. The geometry will be that which corresponds to that number of electron domains.

45 Molecular Geometries The electron-domain geometry is often not the shape of the molecule, however. The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

46 Molecular Geometries Within each electron domain, then, there might be more than one molecular geometry.

47 Group 2 Geometries Central Atom Bonding Electrons Lone Pair Electrons Bond Type Shape Example Magnesium 2 all single linear MgI2 one double MgO In this domain, there is only one molecular geometry: linear. NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

48 Group 13 Geometries Central Atom Bonding Electrons Lone Pair Electrons Bond Type Shape Example Boron 3 all single trigonal planar BI3 one double linear BIO one triple BN because there are no lone pair electrons the molecular is planar (flat).

49 Group 14 Geometries Central Atom Bonding Electrons Lone Pair Electrons
Bond Type Shape Example Carbon 4 all single tetrahedral CH4 one double trigonal planar COH2 one triple, or two double linear HCN CO2

50 Summary for Groups 2, 13 & 14 because there are no lone pair electrons: central atom bonded to 1 atom: linear central atom bonded to 2 atoms: linear central atom bonded to 3 atoms: trigonal planar central atom bonded to 4 atoms: tetrahedral

51 Group 15 Geometries Central Atom Bonding Electrons Lone Pair Electrons
Bond Type Shape Example Nitrogen 3 1 all single trigonal pyramidal NH3 one double angular NOH one triple linear N2 5 trigonal bipyramidal NCl5

52 Group 16 Geometries Central Atom Bonding Electrons Lone Pair Electrons
Bond Type Shape Example Oxygen 2 all single angular H2O one double linear O2 6 octahedral OF6

53 Group 17 & 18 Geometries group 17 group 18
normally makes 1 chemical bond (linear) with valence level expansion can make 7 bonds (ClF7). group 18 normally makes no bonds with valence level expansion can make 8 bonds (ArF8).

54 Larger Molecules In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole. This molecule is tetrahedral about the first carbon, planar trigonal about the second and angular about the single-bonded oxygen.

55 Larger Molecules This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule.

56 Using VSEPR Theory Determine the shape of each molecule, or atom within the molecule: F2 SiF4 PCl3 C3H8 SeH2 C3H6 MgH2 C3H4

57 Polarity previously we discussed bond dipoles.
But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.

58 Polarity By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule. In other words we can see if the dipoles reinforce each other, or cancel out.

59 Polarity

60 Assignment Perform the following activities for the molecules at the bottom of the slide: Draw a lewis structure Draw a structural diagram and indicate shape. Determine the electronegativity difference for each bond type and determine whether the bond is covalent, polar covalent or ionic Determine if they will be polar, or non-polar. F2 PCl3 SeH2 MgH2 SiF4 C3H8 C3H6 C3H4

61 Complete “Fun With Balls & Sticks”

62 Properties of Molecules
Intermolecular Forces

63 Properties of Molecules
are determined by how the molecules interact with each other. How they interact is determined by the forces of attraction between molecules. These are called intermolecular forces; the forces which act between molecules, to draw them together, forming the various phases of matter.

64 How do we measure force ? the best way is by temperature.
temperature is a measure of kinetic energy; the lowest temperature represents zero kinetic energy. for a substance to melt or boil the kinetic energy must overcome the intermolecular force. the higher the melting temperature or boiling temperature, the greater the intermolecular force.

65 Van der Waals Forces occur between covalently bonded molecules.
three kinds: London disersion dipole-dipole attraction hydrogen bonding

66 London Dispersion Forces
are the dominant forces between covalently bonded, non-polar molecules based on the formation of instantaneous dipoles. the more electrons in a molecule, the stronger the force.


68 Relationship of Boiling Point to Number of Electrons and Molar Mass
Melting Point Boiling Point # e °C °C CH C2H C3H C4H C8H

69 London dispersion forces
are influenced by shape: Normal Pentane (C5H12) m.p. -130C, b.p. 36°C Neopentane (C5H12) m.p. -20°C, b.p. 9°C




73 Dipole-Dipole attraction
is a force which acts between polar molecules (ex. H2S). results from the attraction of the opposite poles of the permanent molecular dipoles. These substances generally have higher melting and boiling points than non-polar molecules with similar molecular weights (or numbers of electrons).


75 Ion-Dipole Interactions
A fourth type of force, ion-dipole interactions are an important force in solutions of ions. The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.


77 Dipole-Dipole Interactions
The more polar the molecule, the higher is its boiling point.

78 How Do We Explain This? The nonpolar series (SnH4 to CH4) follow the expected trend. The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.

79 Hydrogen Bonding a specialized form of dipole-dipole attraction.
It occurs as when O, N, and F are bonded to H, owing to the large electronegativity difference: O - H = 1.4 N - H = 1.0 F - H = 2.0 This is a stronger force than standard dipole-dipole attraction. Molecules with hydrogen bonding will have boiling points and melting points quite a bit higher than molecules that have only dipole-dipole or London dispersion forces. Hydrogen bonding is responsible for many of the unusual properties of water.



82 Hydrogen bonding is responsible for folding, final structure and function of proteins holds the DNA strands together, but allows them to be unzipped for copying and gene expression.

83 Relationship Between Polarity, Force and Boiling Point
# e- B.P. (°C) Polarity Force C2H Non-polar London Dispersion H2S Polar Dipole Dipole H2O Very polar Hydrogen Bonds

84 Ionic Bonding generally occurs in compounds of metals and non-metals (salts). It is the result of the attraction of oppositely charged ions. The structures formed are very orderly and are given the name crystal lattice. Ionic solids are called crystals. No sharing of electrons occurs between the ions in the crystal lattice. As a result, ionic solids are brittle. Ionic solids conduct electricity only in the molten state, and not very well. Ionic solids are characterized by very high melting and boiling points.


86 Metallic Bonding is the bonding which occurs between metals in the Periodic Table. It is characterized by close packing of the atoms, with the electrons delocalized; that is, they are free to jump from atom to atom, filling unoccupied orbitals. This free sharing of electrons allows metals to conduct electricity freely (copper conducts electricity times better than molten NaCl). The free electrons also act as a lubricant, allowing metal atoms to slide over one another without affecting the integrity of the material. Thus metals are malleable and ductile. This bond is strong, giving most metals high melting and boiling points. This bond is also variable.



89 Network Covalent Bonding
This is the traditional covalent bond, expanded to 2 or 3 dimensions in a network which is theoretically infinite (much like an ionic crystal lattice). Network solids include diamonds, graphite, quartz, and most rocks. Because the covalent bond is stronger than any other bond, the network solid is very hard (diamonds are the hardest substance known). Because electrons are held tightly in their bonds, network solids are brittle, and they do not conduct electricity. Because the orientation of the atoms is very specific to the bonding orientation (tetrahedral, planar trigonal), network solids form distinct crystals. Because of the strength of the bonds, network solids have very high melting and boiling points



92 Summarizing Intermolecular Forces

93 Intermolecular Forces Affect Many Physical Properties
The strength of the attractions between particles can greatly affect the properties of a substance or solution.

94 Viscosity Resistance of a liquid to flow is called viscosity.
It is related to the ease with which molecules can move past each other. Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

95 Surface Tension Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

96 Vapor Pressure At any temperature, some molecules in a liquid have enough energy to escape. As the temperature rises, the fraction of molecules that have enough energy to escape increases.

97 Vapor Pressure As more molecules escape the liquid, the pressure they exert increases.

98 Vapor Pressure The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

99 Vapor Pressure The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. The normal boiling point is the temperature at which its vapor pressure is 760 torr.

100 Intermolecular Forces in Summary

101 Strength of Attraction
Molecular (London dispersion) 1 to 50 (dipole-dipole) 5 (hydrogen bonding) 15 Ionic 50 Metallic 20 to 80 Network Covalent 100

102 Melting and Boiling Point
Molecular (London dispersion) very low (nitrogen boils at C) VARIABLE (dipole-dipole) low (H2S boils at -61 C) (hydrogen bonding) medium (H2O boils at +100 C) Ionic high (NaCl boils at C) Metallic variable +357 C, C) Network Covalent very high (SiO2 boils at 2600 C)

103 Properties of Solids Molecular (London dispersion) soft, waxy
(dipole-dipole) more rigid (hydrogen bonding) crystalline, brittle Ionic long-range crystalline, hard, brittle Metallic short-range crystalline, ductile, malleable Network Covalent

104 Conductance of Heat and Electricity
Molecular (London dispersion) non-conductor (dipole-dipole) (hydrogen bonding) Ionic conductor in liquid or dissolved phase Metallic conductor in solid or liquid phase Network Covalent non-conductor in 3d form, some conductance in 2d

105 Solubility in H2O Molecular (London dispersion) Not soluble
(dipole-dipole) Soluble (hydrogen bonding) Ionic Metallic Network Covalent

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