1. Chemeketa Community College
Chemical Bonds
Chapter 11
Larry Emme
Chemeketa Community College

2. Periodic Trends in Atomic Properties

3. Characteristic properties and trends of the elements are the basis of the periodic table’s design.

4. These trends allow us to use the periodic table to accurately predict properties and reactions of a wide variety of substances.

5. Metals and Nonmetals

6. Chemical Properties of Metals
Chemical Properties of Nonmetals
metals tend to lose electrons and form positive ions (cations).
nonmetals tend to gain electrons and form negative ions (anions).
When metals react with nonmetals electrons are usually transferred from the metal to the nonmetal.

7. Metals are found to the left of the metalloids
Nonmetals are found to the right of the metalloids.

8. Atomic Radius

9. Which one is bigger?
2 pounds
100 pounds

10. Radii of atoms increase down a group.
For each step down a group, electrons enter the next higher energy level.

11. Radii of atoms tend to decrease from left to right across a period.
For representative elements within the same period the energy level remains constant as electrons are added.
Each time an electron is added a proton is a added to the nucleus. This increase in positive nuclear charge pulls all electrons closer to the nucleus.

12. Ionization Energy

13. Na + ionization energy → Na+ + e-
The ionization energy of an atom is the energy required to remove the outermost electron from an atom.
Na + ionization energy → Na+ + e-

14. Ionization energies gradually increase from left to right across a period.
Noble Gases 1 2
VIIA VA IA
VIA 3
IVA 4
IIA
IIIA
Periodic relationship of the first ionization energy for representative elements in the first four periods.

15. nonmetals have higher ionization potentials than metals
Ionization energies of Group A elements decrease from top to bottom in a group.
Noble Gases
nonmetals have higher ionization potentials than metals
VIIA VA IA
VIA
IVA
Distance of Outer Shell Electrons From Nucleus
IIA
IIIA
nonmetals
metals
Periodic relationship of the first ionization energy for representative elements in the first four periods.

16. Lewis “Dot” Structures of Atoms

17. Metals form cations and nonmetals form anions to attain a stable valence electron structure.

18. This stable structure often consists of two s and six p electrons.
These rearrangements occur by losing, gaining, or sharing electrons.
This stable structure often consists of two s and six p electrons.

19. The Lewis structure of an atom is a representation that shows the valence electrons for that atom.
Na with the electron structure 1s22s22p63s1 has 1 valence electron.
Fluorine with the electron structure 1s22s22p5 has 7 valence electrons

20. The Lewis structure of an atom uses dots to show the valence electrons of atoms.
Unpaired electron B
Paired electrons
Symbol of the element
2s22p1
The number of dots equals the number of s and p electrons in the atom’s outermost shell.

21. The Lewis structure of an atom uses dots to show the valence electrons of atoms.
3s23p4
The number of dots equals the number of s and p electrons in the atom’s outermost shell.

22. Lewis Structures of the first 20 elements.

23. The chemistry of many elements, especially the representative ones, is to attain the same outer electron structure as one of the noble gases.

24. With the exception of helium, this structure consists of eight electrons in the outermost energy level.

25. The Covalent Bond: Sharing Electrons

26. A covalent bond consists of a pair of electrons shared between two atoms.
In the millions of chemical compounds that exist, the covalent bond is the predominant chemical bond.

27. Substances which covalently bond exist as molecules.
Carbon dioxide bonds covalently. It exists as individually bonded covalent molecules containing one carbon and two oxygen atoms.

28. The term molecule is not used when referring to ionic substances.
Sodium chloride bonds ionically. It consists of a large aggregate of positive and negative ions. No molecules of NaCl exist.

29. Covalent bonding in the hydrogen molecule
The most likely region to find the two electrons is between the two nuclei.
Two 1s orbitals from each of two hydrogen atoms overlap.
The orbital of the electrons includes both hydrogen nuclei.
Each 1s orbital contains 1 electron.
The two nuclei are shielded from each other by the electron pair. This allows the two nuclei to draw close together.

30. A dash may replace a pair of dots.
Covalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element.
chlorine
iodine
hydrogen
nitrogen
A dash may replace a pair of dots.
H-H

31. H H-H H2 Cl Cl-Cl Cl2 O O=O O2 N2 N N
Covalent Bonds H
H-H H2 Cl
Cl-Cl Cl2 O
O=O O2 N2 N
N N

32. Electronegativity
Linus Pauling

33. electronegativity The relative attraction that an atom has for a pair of shared electrons in a covalent bond.

34. If the two atoms that constitute a covalent bond are identical then there is equal sharing of electrons.
This is called nonpolar covalent bonding.
Ionic bonding and nonpolar covalent bonding represent two extremes.

36. This is called polar covalent bonding.
If the two atoms that constitute a covalent bond are not identical then there is unequal sharing of electrons.
This is called polar covalent bonding.
One atom assumes a partial positive charge and the other atom assumes a partial negative charge.
This charge difference is a result of the unequal attractions the atoms have for their shared electron pair.

38. Bond Polarity, ElectronegativityCl H
This unequal sharing results in polar bonds.  –
H Cl
Slight positive side
Smaller electronegativity
Slight negative
Larger electronegativity 67 67

39. A scale of relative electronegativities was developed by Linus Pauling.

40. Electronegativity generally increases left to right across a period.

41. Electronegativity increases up a group for representative elements.

42. A dipole can be written as
A dipole is a molecule that is electrically asymmetrical, causing it to be oppositely charged at two points.
A dipole can be written as
+ -

43. An arrow can be used to indicate a dipole.
The arrow points to the negative end of the dipole.
Molecules of HCl, HBr and H2O are polar . H Cl Br O

44. A molecule containing different kinds of atoms may or may not be polar depending on its shape.
The carbon dioxide molecule is nonpolar because its carbon-oxygen dipoles cancel each other by acting in opposite directions.

45. Lewis Structures of Compounds

46. In writing Lewis structures, the most important consideration for forming a stable compound is that the atoms attain a noble gas configuration.

47. The most difficult part of writing Lewis structures is determining the arrangement of the atoms in a molecule or an ion.
In simple molecules with more than two atoms, one atom will be the central atom surrounded by the other atoms.

48. Cl2O has two possible arrangements.
The two chlorines can be bonded to each other.
Cl-Cl-O
The two chlorines can be bonded to oxygen.
Cl-O-Cl
Usually the single atom will be the central atom.

49. Practice Writing Lewis Structures

50. Valence Electrons of Group A Elements
Atom
Group
Valence Electrons Cl
VIIA 7 H IA 1 C
IVA 4 N VA 5 S
VIA 6 P I

51. 3-Dimensional Shapes Linear 180 Bent 105 Trigonal Planar 120
Tetrahedral
109.5
Trigonal
Pyramidal
107

52. Covalent Bonding Structures
Molecular
Formula
Lewis “dot”
Structure
3-D
Name
Bond
Angle
Polar or
Non-polar
H2O
CO2
PH3
NO3–
CH4

53. Covalent Bonding Structures
Molecular
Formula
Lewis “dot”
Structure
3-D
Name
Bond
Angle
Polar or
Non-polar
H2O
CO2
PH3
NO3–
CH4
Bent
105
Polar
Linear
180
Non-Polar
Trigonal
Pyramidal
107
Polar
Trigonal
Planar
120
Non-Polar
Tetrahedral
109.5
Non-Polar

54. H H O H H

55. C O O O O

56. H H P H H H H

57. Electron from cation O
(—) O N O O O O

58. O O O O O O N N N ↔ ↔ O O O (—) (—) (—) Resonance Structures
In chemistry, resonance is a way of describing delocalized electrons where the bonding cannot be expressed by one single Lewis structure.
The molecule or ion with such delocalized electrons is represented by several contributing structures.

59. H H H H C H H H H

60. The Ionic Bond: Transfer of Electrons From One Atom to Another

62. After sodium loses its 3s electron it has attained the same electronic structure as neon.

63. After chlorine gains a 3p electron it has attained the same electronic structure as argon.

64. Formation of NaCl

65. A sodium ion (Na+) and a chloride ion (Cl-) are formed.
The 3s electron of sodium transfers to the half-filled 3p orbital of chlorine.
The force holding Na+ and Cl- together is an ionic bond.
Lewis representation of sodium chloride formation.

66. Formation of MgCl2

67. The forces holding Mg2+ and two Cl- together are ionic bonds.
A magnesium ion (Mg2+) and two chloride ions (Cl-) are formed.
Two 3s electrons of magnesium transfer to the half-filled 3p orbitals of two chlorine atoms.

68. In the crystal each sodium ion is surrounded by six chloride ions.
NaCl is made up of cubic crystals.

69. In the crystal each chloride ion is surrounded by six sodium ions.

70. The ratio of Na+ to Cl- is 1:1
There is no molecule of NaCl

71. Metals usually have one, two or three electrons in their outer shells.
When a metal reacts it:
usually loses one two or three electrons
attains the electron structure of a noble gas
becomes a positive ion.
Courtesy:
Diane Schmitz
The positive ion formed by the loss of electrons is much smaller than the metal atom.

72. When a nonmetal reacts it:
Nonmetals are usually only a few electrons short of having a noble gas structure.
When a nonmetal reacts it:
usually gains one two or three electrons
attains the electron structure of a noble gas
becomes a negative ion.
The negative ion formed by the gain of electrons is much larger than the nonmetal atom.

73. Predicting Formulas of Ionic Compounds

74. In almost all stable chemical compounds of representative elements, each atom attains a noble gas electron configuration.

75. Ions are always formed by adding or removing electrons from an atom.

76. Most often ions are formed when metals combine with nonmetals.
Metals will lose electrons to attain
a noble gas configuration.
Nonmetals will gain electrons to attain
a noble gas configuration.

77. The charge on an ion can be predicted from its position in the periodic table.

78. elements of Group IA have a +1 charge

79. elements of Group IIA have a +2 charge

80. elements of Group VA have a -3 charge

81. elements of Group VIA have a -2 charge

82. elements of Group VIIA have a -1 charge

84. Writing Formulas From Names of Compounds

85. A chemical compound must have a net charge of zero.

86. If the compound contains ions, then the charges on all of the ions must add to zero.

87. Write the formula of calcium chloride.
Step 1. Write down the formulas of the ions by placing the cation first and the anion second.
Ca2+ Cl-1
Step 2. Place the number of the charge of the cation as a subscript of the anion and the number of the charge of the anion as a subscript of the cation.

88. Ca2+ Cl-1
The correct formula is
CaCl2
Charges are removed!

89. Write the formula of barium phosphide.
Step 1. Write down the formulas of the ions by placing the cation first and the anion second.
Ba2+ P-3
Step 2. Place the number of the charge of the cation as a subscript of the anion and the number of the charge of the anion as a subscript of the cation.

90. Ba2+ P-3
The correct formula is
Ba3P2
Charges are removed!

91. Write the formula of magnesium oxide.
Step 1. Write down the formulas of the ions by placing the cation first and the anion second.
Mg2+ O-2
Step 2. Place the number of the charge of the cation as a subscript of the anion and the number of the charge of the anion as a subscript of the cation.

92. Mg2+ O-2
The correct formula is
Mg2O2
MgO
Reduce to lowest ratio!

93. Thanks to Lauren Blaco Fall 2014

94. Write the Formula of Sodium Peroxide
Na2O2
NaO
gives or

95. Write the Formula of Sodium Peroxide
Don’t mess with the subscripts
of polyatomic ions!!
Na2O2
NaO
gives
not
NaO does not contain the peroxide anion

96. Combine to Give Compounds (Do Not Name!)
ions
Br–
O–2
NO3–
PO4–3
CO3–2
NH4+
Sn+2
Al+3 H+

97. Combine to Give Compounds (Do Not Name!)
ions
Br–
O–2
NO3–
PO4–3
CO3–2
NH4+
Sn+2
Al+3 H+
NH4Br
(NH4)2O
NH4NO3
(NH4)3PO4
(NH4)2CO3
SnBr2
SnO
Sn(NO3)2
Sn3(PO4)2
SnCO3
AlBr3
Al2O3
Al(NO3)3
AlPO4
Al2(CO3)3
HBr
H2O
HNO3
H3PO4
H2CO3

98. The End