1. Chemical Bonds The Formation of Compounds From Atoms

2. O U T L I N E Periodic Trends in Atomic Properties
Lewis Structures of Atoms
The Ionic Bond: Transfer of Electrons From One Atom to Another
Predicting Formulas of Ionic Compounds
The Covalent Bond: Sharing Electrons
Electronegativity
Lewis Structures of Compounds
Complex Lewis Structures
Compounds Containing Polyatomic Ions
Molecular Shape
The Valence Shell Electron Pair (VSEPR) Model

3. Periodic Trends in Atomic Properties

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

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

6. Consider metals and nonmetals:
Chemical property of metals: metals tend to lose electrons and form positive ions.
Chemical property of nonmetals: nonmetals tend to gain electrons and form negative ions.
When metals react with nonmetals, electrons are usually transferred from the metal to the nonmetal.

7. Consider metals and nonmetals:
Physical properties of metals: lustrous, malleable, good conductors of heat, good conductors of electricity
Physical properties of nonmetals: nonlustrous, brittle, poor conductors of heat, poor conductors of electricity
Metalloids have properties that are intermediate between metals and nonmetals.

8. The Metalloids boron silicon germanium arsenic antimony tellurium
polonium

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

10. Atomic radii 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 also added to the nucleus.
This increase in positive nuclear charge pulls all electrons closer to the nucleus.

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

13. The first ionization energy is the amount of energy required to remove the first electron from an atom.
He + first ionization energy → He+ + e-
He + 2,372 kJ/mol → He+ + e-
The second ionization energy is the amount of energy required to remove the second electron from an atom.
He+ + second ionization energy → He2+ + e-
He+ + 5,247 kJ/mol → He2+ + e-

14. As each succeeding electron is removed from an atom, ever higher energies are required.

15. Ionization energies gradually increase from left to right across a period.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.

16. 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.

17. Metals form cations and nonmetals form anions to attain a stable valence electron structure. These rearrangements occur by losing, gaining, or sharing electrons.

18. Fluorine with the electron structure 1s22s22p5 has 7 valence electrons
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

19. Lewis Structures of Atoms

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. With the exception of helium, this structure consists of eight electrons in the outermost shell (highest energy level).

24. The Ionic Bond: Transfer of electrons from one atom to another

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

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

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

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

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

30. In the crystal each chloride ion is surrounded by six sodium ions.
11.5

31. The ratio of Na+ to Cl- is 1:1
There is no molecule of NaCl
11.5

32. A sodium ion is smaller than a sodium atom because:
the sodium atom has lost its outermost electron and
the 10 remaining electrons are now attracted by 11 protons and are drawn closer to the nucleus.

33. A chloride ion is larger than a chlorine atom because:
the chlorine atom has gained an electron and now has 18 electrons and 17 protons.
The nuclear attraction on each electron has decreased, allowing the chlorine to expand.

34. 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.
The positive ion formed by the loss of electrons is much smaller than the metal atom.

35. Nonmetals usually have one, two or three electrons in their outer shells.
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.

36. Predicting Formulas of Ionic Compounds

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

38. Barium and Sulfur Combine.
Metals will lose electrons to attain a noble gas configuration. Nonmetals will gain electrons to attain a noble gas configuration.
Barium and Sulfur Combine.
barium loses two electrons to sulfur and attains a xenon configuration.
sulfur gains two electrons from barium and attains an argon configuration.
Ba → [Xe] + 2e-
S [Ne]3s23p4
Ba [Xe]6s2
S + 2e- → [Ar]
Ba + S → BaS

39. Because of similar electron structures, the elements of a family generally form compounds with the same atomic ratios.

41. The elements of a family have the same outermost electron configuration except that the electrons are in different energy levels.
The group numbers for the representative elements are equal to the total number of outermost electrons in the atoms of the group.

42. The atomic ratio of the alkali metal sodium to chlorine is 1:1 in NaCl.
The atomic ratios of the other alkali metal chlorides can be predicted to also be 1:1.
LiCl, KCl, CsCl

43. Predict the formulas of these compounds
Lithium and Bromine Lithium and Oxygen
LiBr Li2O
Magnesium and Bromine Magnesium and Oxygen
MgBr MgO
Aluminum and Bromine Aluminum and Oxygen
AlBr Al2O3

44. The Covalent Bond: Sharing Electrons

45. 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.

46. 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.

47. 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.

48. Covalent bonding in the hydrogen molecule
The most likely region to find the two electrons is between the two nuclei.
The orbital of the electrons includes both hydrogen nuclei.
Two 1s orbitals from each of two hydrogen atoms overlap.
Two 1s orbitals from each of two hydrogen atoms overlap.
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.

49. Covalent bonding in the chlorine molecule
The orbital of the electrons includes both chlorine nuclei.
The two nuclei are shielded from each other by the electron pair. This allows the two nuclei to draw close together.
Two 3p orbitals from each of two chlorine atoms overlap.
Two 3p orbitals from each of two chlorine atoms overlap.
The most likely region to find the two electrons is between the two nuclei.
Each unpaired 3p orbital on each chlorine atom contains 1 electron.
Each chlorine now has 8 electrons in its outermost energy level.

50. 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.

51. Electronegativity

52. Electronegativity: The relative attraction that an atom has for a pair of shared electrons in a covalent bond.

53. 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.

54. 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.

55. Polar Covalent Bonding in HCl
Partial positive charge on hydrogen.
Partial negative charge on chlorine.
Polar Covalent Bonding in HCl : H Cl + - :
Chlorine has a greater attraction for the shared electron pair than hydrogen.
The attractive force that an atom of an element has for shared electrons in a molecule or a polyatomic ion is known as its electronegativity.
Shared electron pair.
The shared electron pair is closer to chlorine than to hydrogen.

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

57. Electronegativity generally increases left to right across a period.
Electronegativity decreases down a group for representative elements.

58. The electronegativities of the nonmetals are high.
The electronegativities of the metals are low.

59. The polarity of a bond is determined by the difference in electronegativity values of the atoms forming the bond.

60. If the electronegativity difference between two bonded atoms is greater than , the bond will be more ionic than covalent.
If the electronegativity difference is greater than 2, the bond is strongly ionic.
If the electronegativity difference is less than 1.5, the bond is strongly covalent.

61. If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally. H
Hydrogen Molecule

62. If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally. Cl
Chlorine Molecule

63. If the electronegativities are not the same, the bond is polar covalent and the electrons are shared unequally. H Cl + -
Electronegativity Difference = 0.9
Electronegativity
2.1
3.0
Hydrogen Chloride Molecule

64. If the electronegativities are very different, the bond is ionic and the electrons are transferred to the more electronegative atom.
Electronegativity Difference = 2.1
Sodium Chloride
Na+
Cl-
Electronegativity
0.9
3.0

65. 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
+ -

66. 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

67. 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.

68. Relating Bond Type to Electronegativity Difference.

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

70. 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.

71. Lewis Structures of Compounds

72. Valence Electrons of Group A Elements
Atom
Group
Valence Electrons Cl 7A 7 H 1A 1 C 4A 4 N 5A 5 S 6A 6 P I

73. Step 1. Obtain the total number of valence electrons to be used in the structure by adding the number of valence electrons in all the atoms in the molecule or ion.
If you are writing the structure of an ion, add one electron for each negative charge or subtract one electron for each positive charge on the ion.

74. Write the Lewis structure for H2O.
Step 1. The total number of valence electrons is eight, two from the two hydrogen atoms and six from the oxygen atom.

75. Step 2. Write the skeletal arrangement of the atoms and connect them with a single covalent bond (two dots or one dash).
Hydrogen, which contains only one bonding electron, can form only one covalent bond.
Oxygen atoms normally have a maximum of two covalent bonds (two single bonds, or one double bond).

76. : : H O H H O H Write the Lewis structure for H2O. or
Step 2. The two hydrogen atoms are connected to the oxygen atom. Write the skeletal structure: : :
H O H or
H O H
Place two dots between the hydrogen and oxygen atoms to form the covalent bonds.

77. Step 3. Subtract two electrons for each single bond you used in Step 2 from the total number of electrons calculated in Step 1.
This gives you the net number of electrons available for completing the structure.

78. : H O H Write the Lewis structure for H2O.
Step 3. Subtract the four electrons used in Step 2 from eight to obtain four electrons yet to be used.
H O H :

79. Step 4. Distribute pairs of electrons (pairs of dots) around each atom (except hydrogen) to give each atom a noble gas configuration.

80. : : : : H O H H O H Write the Lewis structure for H2O. or
Step 4. Distribute the four remaining electrons in pairs around the oxygen atom. Hydrogen atoms cannot accommodate any more electrons. : : : :
H O H or
H O H
The shape of the molecule is not shown by the Lewis structure.
These arrangements are Lewis structures because each atom has a noble gas electron structure.

81. Write a Lewis structure for CO2.
Step 1. The total number of valence electrons is 16, four from the C atom and six from each O atom.

82. : O C O Write a Lewis structure for CO2.
Step 2. The two O atoms are bonded to a central C atom. Write the skeletal structure and place two electrons between the C and each oxygen. :
O C O

83. : O C O Write a Lewis structure for CO2.
Step 3. Subtract the four electrons used in Step 2 from 16 (the total number of valence electrons) to obtain 12 electrons yet to be used.
O C O :

84. : : : : : : O C O O C O O C O Write a Lewis structure for CO2.
Step 4. Distribute the 12 electrons (6 pairs) around the carbon and oxygen atoms. Three possibilities exist. :
O C O : : : :
O C O :
O C O
Many of the atoms in these structures do not have eight electrons around them.

85. : : : O C O O C O O C O Write a Lewis structure for CO2.
Step 5. Remove one pair of unbonded electrons from each O atom in structure I and place one pair between each O and the C atom forming two double bonds.
O C O :
O C O :
O C O :
Carbon is sharing 4 electron pairs.
double bond
Each atom now has 8 electrons around it.

86. Complex Lewis Structures

87. There are some molecules and polyatomic ions for which no single Lewis structure consistent with all characteristics and bonding information can be written.

88. Write a Lewis structure for NO31-
Step 1. The total number of valence electrons is 24, 5 from the nitrogen atom and 6 from each O atom, and 1 from the –1 charge.

89. : O O N O Write a Lewis structure for NO31-
Step 2. The three O atoms are bonded to a central N atom. Write the skeletal structure and place two electrons between each pair of atoms.
O N O : O

90. : O O N O Write a Lewis structure for NO31-
Step 3. Subtract the 6 electrons used in Step 2 from 24, the total number of valence electrons, to obtain 18 electrons yet to be placed.
O N O : O

91. : : O O N O Write a Lewis structure for NO31-
Step 4. Distribute the 18 electrons around the N and O atoms. :
O N O O
electron deficient :

92. : : - O O N O Write a Lewis structure for NO31-
Step 4. Since the extra electron present results in nitrate having a –1 charge, the ion is enclosed in brackets with a – charge. - : O :
O N O

93. : : : : : : - O O O N Write a Lewis structure for NO31-
Step 5. One of the oxygen atoms has only 6 electrons. It does not have a noble gas structure. Move the unbonded pair of electrons from the N atom and place it between the N and the electron-deficient O atom, making a double bond. - O : O : O : N : : :
electron deficient

94. : : : N O - N O - N O - Write a Lewis structure for NO31-
Step 5. There are three possible Lewis structures.
A molecule or ion that shows multiple correct Lewis structures exhibits resonance. Each Lewis structure is called a resonance structure. N O : - N O : - : N O -

95. Compounds Containing Polyatomic Ions

96. A polyatomic ion is a stable group of atoms that has either a positive or negative charge and behaves as a single unit in many chemical reactions.

97. Molecular Shape

98. The 3-dimensional arrangement of the atoms within a molecule is a significant feature in understanding molecular interactions.

100. The Valence Shell Electron Pair Repulsion Model (VSEPR)

101. To accomplish this minimization, the electron pairs will be arranged as far apart as possible around a central atom.
The VSEPR model is based on the idea that electron pairs will repel each other electrically and will seek to minimize this repulsion.

102. BeCl2 is a molecule with only two pairs of electrons around beryllium, its central atom.
Its electrons are arranged 180o apart for maximum separation.

103. This arrangement of atoms is called trigonal planar.
BF3 is a molecule with three pairs of electrons around boron, its central atom.
This arrangement of atoms is called trigonal planar.
Its electrons are arranged 120o apart for maximum separation.

104. CH4 is a molecule with four pairs of electrons around carbon, its central atom.
An obvious choice for its atomic arrangement is a 90o angle between its atoms with all of its atoms in a single plane.
However, since the molecule is 3-dimensional, the molecular structure is tetrahedral with a bond angle of 109.5o.

105. Ball and stick models of methane, CH4, and carbon tetrachloride, CCl4.

106. Ammonia, NH3, has four electron pairs around nitrogen.
The arrangement of the electron pairs is tetrahedral.

107. One of its electron pairs is a nonbonded (lone) pair.
The shape of the NH3 molecule is pyramidal.

108. Water has four electron pairs around oxygen.
The arrangement of electron pairs around oxygen is tetrahedral.

109. Two of its electron pairs are nonbonded (lone) pairs.
The H2O molecule is bent.

110. Structure Determination Using VSEPR
Draw the Lewis structure for the molecule.
Count the electron pairs and arrange them to minimize repulsions.
Determine the positions of the atoms.
Name the molecular structure from the position of the atoms.