1. Chapter 5 Compounds and Their Bonds
5.1 Octet Rule and Ions

2. Octet Rule An octet is 8 valence electrons
is associated with the stability of the noble gases
does not occur with He; He is stable with 2 valence electrons (duet)
Valence Electrons
He 1s
Ne 1s22s22p
Ar 1s22s22p63s23p
Kr 1s22s22p63s23p64s23d104p

3. Ionic and Covalent Bonds
Atoms form octets
to become more stable
by losing, gaining, or sharing valence electrons
by forming ionic or covalent bonds

4. Metals Form Positive Ions
by a loss of their valence electrons
with the electron configuration of the nearest noble gas
that have fewer electrons than protons
Group 1A(1) metals ion 1+
Group 2A(2) metals ion 2+
Group 3A(3) metals ion 3+

5. Formation of a Sodium Ion, Na+
Sodium achieves an octet by losing its one valence electron.
With the loss of its valence
electron, a sodium ion has a 1+
charge. Sodium atom Sodium ion
11p p+
11e– e–

6. Formation of Magnesium Ion, Mg2+
Magnesium achieves an octet by losing its two valence electrons.
With the loss of two valence
electrons magnesium forms a
positive ion with a 2+ charge Mg atom Mg2+ ion
12p+ 12p+
12e– e–

7. Formation of Negative Ions
In ionic compounds, nonmetals
achieve an octet arrangement
gain electrons
form negatively charged ions with 3–, 2–, or 1– charges

8. Formation of a Chloride Ion, Cl–
Chlorine achieves an octet by adding an electron to
its valence electrons.

9. Ionic Charge from Group Numbers
The charge of a positive ion is equal to its Group number.
Group 1A(1) = 1+
Group 2A(2) = 2+
Group 3A(3) = 3+
The charge of a negative ion is obtained by subtracting 8 or 18 from its Group number .
Group 6A(16) = – 8 = 2–
or 16 – 18 = 2–

10. Group Number and Ionic Charge
Ions
achieve the electron configuration of their
nearest noble gas
of metals in Groups 1A(1), 2A(2), or 3A(13) have positive 1+, 2+, or 3+ charge.
Of nonmetals in Groups 5A(15), 6A(16), or 7A(17) have negative 3–, 2–, or 1– charge.

11. Groups Numbers for Some Positive and Negative Ions

12. Chapter 5 Compounds and Their Bonds
5.2 Ionic Compounds

13. Ionic Compounds Ionic compounds consist of positive and negative ions
have attractions called ionic bonds between positively and negatively charged ions
have high melting and boiling points
are solid at room temperature

14. Salt Is an Ionic Compound
Sodium chloride, or “table salt,” is an example of an ionic compound.

15. Ionic Formulas An ionic formula
consists of positively and negatively charged ions
is neutral
has charge balance
total positive charge = total negative charge
The symbol of the metal is written first, followed by the symbol of the nonmetal.

16. Charge Balance for NaCl, “Salt”
In NaCl,
a Na atom loses its valence electron
a Cl atom gains an electron
the symbol of the metal is written first, followed by the symbol of the nonmetal.

17. Charge Balance In MgCl2 In MgCl2,
a Mg atom loses two valence electrons
two Cl atoms each gain one electron
subscripts indicate the number of ions needed to give charge balance

18. Writing Ionic Formulas from Charges
Charge balance is used to write the formula for
sodium nitride, a compound containing Na+ and N3−.
Na+
3 Na N3− = Na3N
3(1+) (3–) =

19. Chapter 5 Compounds and Their Bonds
5.3
Naming and Writing Ionic Formulas

20. Naming Ionic Compounds with Two Elements
To name a compound
with two elements,
identify the cation and anion
name the cation first, followed by the name of the anion

21. Examples of Ionic Compounds with Two Elements
Formula Ions Name
Cation Anion
NaCl Na+ Cl– sodium chloride
K2S K S2– potassium sulfide
MgO Mg2+ O2– magnesium oxide
CaI Ca2+ I– calcium iodide
Al2O3 Al S2– aluminum sulfide

22. Transition Metals Form Positive Ions
Most transition metals and Group 4(14) metals,
Form 2 or more positive ions
Zn2+, Ag+, and Cd2+ form only one ion.

23. Metals with Variable Charge
The names of transition metals with two or more positive ions (cations) use a Roman numeral after the name of the metal to identify ionic charge.

24. Naming Ionic Compounds with Variable Charge Metals

25. Naming FeCl2 STEP 1 Determine the charge of the cation from the anion.
Fe ion Cl– = Fe ion – = 0
Fe ion = 2+ = Fe2+
STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge.
Fe2+ = iron(II)
STEP 3 Write the anion with an ide ending. chloride
STEP 4 Name the cation first, then the anion.
iron(II) chloride

26. Naming Cr2O3
STEP 1 Determine the charge of cation from the anion Cr ions + 3O2– = 2Cr ions + 3(2–)
= 2Cr ions + 6– = 0
2Cr ions = 6+ Cr ion = 3+ = Cr3+
STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge.
Cr3+ = chromium(III)
STEP 3 Write the anion with an ide ending. oxide
STEP 4 Name the cation first, then the anion.
chromium (III) oxide

27. Guide to Writing Formulas from the Name

28. Writing Formulas Write a formula for potassium sulfide.
STEP 1 Identify the cation and anion.
potassium = K+
sulfide = S2−
STEP 2 Balance the charges.
K S2− K+
2(1+) + 1(2–) = 0
STEP 3 Write the cation first.
2K+ and 1S2− = K2S1 = K2S

29. Writing Formulas STEP 1 Identify the cation and anion.
Write a formula for iron(III) chloride.
STEP 1 Identify the cation and anion.
iron (III) = Fe3+ (III = charge of 3+)
chloride = Cl−
STEP 2 Balance the charges.
Fe Cl−
Cl−
1(3+) + 3(1–) = 0
STEP 3 Write the cation first.
1Fe3+ and 3Cl− = FeCl3

30. Chapter 5 Compounds and Their Bonds
5.4 Polyatomic Ions

31. Polyatomic Ions A polyatomic ion is a group of atoms
has an overall ionic charge
Examples:
NH4+ ammonium OH− hydroxide
NO3− nitrate NO2− nitrite
CO32− carbonate PO43− phosphate
HCO3− hydrogen carbonate
(bicarbonate)

32. Some Names of Polyatomic Ions
The names of common polyatomic anions
end in ate
NO3− nitrate PO43− phosphate
with one oxygen less end in ite
NO2− nitrite PO33− phosphite
with hydrogen attached use prefix hydrogen (or bi)
HCO3− hydrogen carbonate (bicarbonate)
HSO3− hydrogen sulfite (bisulfite)

33. Guide to Naming Compounds with Polyatomic Ions

34. Examples of Names of Compounds with Polyatomic Ions
The positive ion is named first followed by the name of the polyatomic ion.
NaNO3 sodium nitrate
K2SO4 potassium sulfate
Fe(HCO3)3 iron(III) bicarbonate
or iron(III) hydrogen carbonate
(NH4)3PO3 ammonium phosphite

35. Writing Formulas with Polyatomic Ions
The formula of an ionic compound
containing a polyatomic ion must have a charge balance that equals zero (0)
Na+ and NO3− NaNO3
with two or more polyatomic ions put the polyatomic ions in parentheses.
Mg2+ and 2NO3− Mg(NO3)2
subscript 2 for charge balance

36. Chapter 5 Compounds and Their Bonds
5.5 Covalent Compounds

37. Forming Octets in Molecules
In a fluorine (F2) molecule, each F atom
shares one electron
acquires an octet

38. Diatomic Elements
These elements share electrons to form diatomic, covalent molecules.

39. Guide to Writing Electron-Dot Formulas

40. Single and Multiple Bonds
In a single bond, one pair of electrons is shared.
In a double bond, two pairs of electrons are shared.
In a triple bond, three pairs of electrons are shared.

41. Electron-Dot Formula of CS2
Write the electron-dot formula for CS2.
STEP 1 Determine the atom arrangement. The C atom is the central atom.
S C S
STEP 2 Determine the total number of valence electrons for 1C and 2S.
1 C(4e–) S(6e–) = 16e–
STEP 3 Attach each S atom to the central C atom using one electron pair.
S : C : S
16e– – 4e– = 12e– remaining
STEP 4 Attach 12 electrons as 6 lone pairs.
: S : C : S :

42. Electron-Dot Formula of CS2 (continued)
To complete octets, form one or more multiple bonds.
Convert two lone pairs to bonding pairs between C and
S atoms to make two double bonds.

43. A Nitrogen Molecule has A Triple Bond
In a nitrogen molecule, N2,
each N atom shares 3 electrons
each N atom attains an octet
the sharing of 3 sets of electrons is a multiple bond called a triple bond

44. Resonance Structures Resonance structures are
two or more electron-dot formulas for the same arrangement of atoms
related by a double-headed arrow ( )
written by changing the location of a double bond between the central atom and a different attached atom

45. Writing Resonance Structures
Sulfur dioxide has two resonance structures.
STEP 1 Write the arrangement of atoms.
O S O
STEP 2 Determine the total number of valence electrons.
1 S(6e−) + 2 O(6e−) = 18e−
STEP 3 Connect bonded atoms by single electron pairs.
O : S : O e− used
18e− – 4e− = 14e− remaining

46. Writing Resonance Structures (continued)
STEP 4 Add 14 remaining electrons as 7 lone pairs.
STEP 5 Form a double bond to complete octets. Two resonance structures are possible.

47. Formula from Ionic Charges
Write the ionic formula of the compound containing Ba2+ and Cl.
Write the symbols of the ions.
Ba2+ Cl
Balance the charges.
Ba2+ Cl two Cl needed Cl
Write the ionic formula using a subscript 2 for two chloride ions that give charge balance.
BaCl2

48. Chapter 5 Compounds and Their Bonds
5.6
Naming and Writing Covalent Formulas
NO nitrogen oxide
NO2 nitrogen dioxide
N2O4 dinitrogen tetroxide

49. Names of Covalent Compounds
Prefixes are used
in the names of covalent compounds
because two nonmetals can form two or more different compounds
Examples of compounds of N and O:
NO nitrogen oxide
NO2 nitrogen dioxide
N2O dinitrogen oxide
N2O4 dinitrogen tetroxide
N2O5 dinitrogen pentoxide

50. Naming Covalent Compounds
STEP 1 Name the first nonmetal by its element name.
STEP 2 Name the second nonmetal with an ide ending.
STEP 3 prefixes to indicate the number (from subscripts) of atoms of each nonmetal. Mono is usually omitted.

51. Naming Covalent Compounds (Continued)
What is the name of SO3?
STEP 1 The first nonmetal is S sulfur.
STEP 2 The second nonmetal is O, named oxide.
STEP 3 The subscript 3 of O is shown as the prefix tri.
SO3 → sulfur trioxide
The subscript 1(for S) or mono is understood.

52. Naming Covalent Compounds (Continued)
Name P4S3
STEP 1 The first nonmetal P is phosphorus.
STEP 2 The second nonmetal S is sulfide.
STEP 3 The subscript 4 of P is shown as tetra.
The subscript 3 of O is shown as tri.
P4S3 → tetraphosphorus trisulfide

53. Guide to Writing Formulas for Covalent Compounds

54. Writing Formulas of Covalent Compounds
Write the formula for carbon disulfide.
STEP 1 Elements are C and S
STEP 2 No prefix for carbon means 1 C
Prefix di = 2
Formula: CS2

55. Chapter 5 Compounds and Their Bonds
5.7 Electronegativity and Bond Polarity

56. Electronegativity The electronegativity value
indicates the attraction of an atom for shared electrons
increases from left to right going across a period on the periodic table
decreases going down a group on the periodic table
is high for the nonmetals, with fluorine as the highest
is low for the metals

57. Some Electronegativity Values
High values
Low values

58. Nonpolar Covalent Bonds
A nonpolar covalent bond
occurs between nonmetals
has an equal or almost equal sharing of electrons
has almost no electronegativity difference (0.0 to 0.4)
Examples:
Atoms Electronegativity Type of Bond Difference _______________
N–N – 3.0 = Nonpolar covalent
Cl–Br 3.0 – 2.8 = Nonpolar covalent
H–Si – 1.8 = Nonpolar covalent

59. Polar Covalent Bonds A polar covalent bond
occurs between nonmetal atoms
has an unequal sharing of electrons
has a moderate electronegativity difference
(0.5 to 1.7)
Examples:
Atoms Electronegativity Type of Bond
Difference
_________________________________________
O–Cl – 3.0 = Polar covalent
Cl–C 3.0 – 2.5 = Polar covalent
O–S – 2.5 = Polar covalent

60. Comparing Nonpolar and Polar Covalent Bonds

61. Chapter 5 Compounds and Their Bonds
5.8 Shapes and Polarity of Molecules

62. VSEPR In the valence-shell electron-pair repulsion theory
(VSEPR), the electron groups around a central atom
are arranged as far apart from each other as possible
have the least amount of repulsion of the negatively charged electrons
have a geometry around the central atom that determines molecular shape

63. Shapes of Molecules The three-dimensional shape of a molecule
is the result of bonded groups and lone pairs of electrons around the central atom
is predicted using the VSEPR theory (valence-shell-electron-pair repulsion)

64. Guide to Predicting Molecular Shape (VSEPR Theory)

65. Two Electron Groups In a molecule of BeCl2,
there are two electron groups bonded to the central atom, Be (Be is an exception to the octet rule)
: Cl : Be : Cl :
to minimize repulsion, the arrangement of two electron groups is 180°, or opposite each other
the shape of the molecule is linear

66. Two Electron Groups with Double Bonds
In a molecule of CO2,
there are two electron groups bonded to C (electrons in each double bond are counted as one group)
repulsion is minimized with the double bonds opposite each other at 180°
the shape of the molecule is linear

67. Three Electron Groups In a molecule of BF3,
three electron groups are bonded to the central atom B (B is an exception to the octet rule) ..
: F:
: F : B : F :
repulsion is minimized with 3 electron groups at angles of 120°
the shape is trigonal planar

68. Two Electron Groups and One Lone Pair
In a molecule of SO2,
S has 3 electron groups; 2 electron groups bonded to O atoms and one lone pair
:O :: S : O : ..
repulsion is minimized with the electron groups at angles of 120°, a trigonal planar arrangement
the shape is bent (120°), with two O atoms bonded to S
● ●

69. Four Electron Groups In a molecule of CH4,
there are four electron groups around C
repulsion is minimized by placing four electron groups at angles of 109°, which is a tetrahedral arrangement
the four bonded atoms form a tetrahedral shape

70. Three Bonding Atoms and One Lone Pair
In a molecule of NH3,
three electron groups bond to H atoms, and the fourth one is a lone (nonbonding) pair
repulsion is minimized with 4 electron groups in a tetrahedral arrangement
the three bonded atoms form a pyramidal (~109°) shape

71. Two Bonding Atoms and Two Lone Pairs
In a molecule of H2O,
two electron groups are bonded to H atoms and two are lone pairs (4 electron groups)
four electron groups minimize repulsion in a tetrahedral arrangement
the shape with two bonded atoms is bent (~109°)

72. Determining Molecular Polarity
Determine the polarity of the H2O molecule. Solution: The four electron groups of oxygen are bonded to two H atoms. Thus the H2O molecule has a net dipole, which makes it a polar molecule.

73. Chapter 5 Compounds and Their Bonds
5.9 Attractive Forces in
Compounds

74. Dipole–Dipole Attractions
In covalent compounds, polar molecules
exert attractive forces called dipole-dipole attractions
form strong dipole attractions called hydrogen bonds between hydrogen atoms bonded to F, O, or N, and other atoms that are very electronegative

75. Dipole–Dipole Attractions (continued)

76. Dispersion Forces Dispersion forces are
weak attractions between nonpolar molecules
caused by temporary dipoles that develop when electrons are not distributed equally

77. Comparison of Bonding and Attractive Forces

78. Melting Points and Attractive Forces
Ionic compounds require large amounts of energy to break apart ionic bonds. Thus, they have high melting points.
Hydrogen bonds are the strongest type of dipole–dipole attractions. They require more energy to break than do other dipole attractions.
Dispersion forces are weak interactions, and very little energy is needed to change state.