1. Chapter Three Chapter 3 Lecture
Fundamentals of General, Organic, and Biological Chemistry
8th Edition
McMurry, Ballantine, Hoeger, Peterson
Chapter Three
Ionic Compounds
Christina A. Johnson
University of California, San Diego
© 2017 Pearson Education, Inc.

2. Outline 3.1 Ions 3.2 Ions and the Octet Rule
3.3 Ions of Some Common Elements
3.4 Periodic Properties and Ion Formation
3.5 Naming Monoatomic Ions
3.6 Polyatomic Ions
3.7 Ionic Bonds
3.8 Formulas of Ionic Compounds
3.9 Naming Ionic Compounds
3.10 Some Properties of Ionic Compounds
3.11 H+ and OH– Ions: An Introduction to Acids and
Bases

3. Concepts to Review The Periodic Table Electron Configuration
Sections 2.4 and 2.5
Electron Configuration
Sections 2.7 and 2.8

4. 3.1 Ions
Ions play critical roles in many cellular processes, including signal transmission between nerve cells.
Learning Objective: Describe ion formation processes and distinguish between anions and cations.

5. 3.1 Ions
Alkali metals (group 1A) form compounds with halogens (group 7A).
Properties of these compounds include the following:
High melting points
Stable, white, crystalline solids
Soluble in water
Conduct electricity when dissolved in water

6. 3.1 Ions
Sodium chloride (table salt) is formed by the reaction of sodium with chlorine.
A solution of sodium chloride in water conducts electricity, allowing the bulb to light.
Electricity can only flow through a medium containing charged particles (ions) that are free to move.

7. 3.1 Ions
An ion is formed when a neutral atom gains or loses electrons.
The loss of one or more electrons from a neutral atom gives a positively charged ion called a cation.
The gain of one or more electrons by a neutral atom gives a negatively charged ion called an anion.

8. 3.1 Ions
Sodium and other alkali metal atoms have a single electron in their valence shell and an electron configuration of ns1.
By losing this electron, an alkali metal is converted to a positively charged cation with a stable noble gas configuration.

9. 3.1 Ions
Chlorine and other halogen atoms have a ns2np5 electron configuration.
By gaining an electron, a halogen is converted to a negatively charged anion with a stable noble gas configuration.

10. 3.1 Ions
The symbol for a cation is written by adding the positive charge as a superscript to the symbol for the element.
The sodium cation would be written as Na+.
The symbol for an anion is written by adding the negative charge as a superscript.
The chlorine anion would be written as Cl–.
If the charge is greater than 1, the number is used, as in Ca2+ and N3–.

11. 3.2 Ions and the Octet Rule
Octet rule: Main group elements tend to undergo reactions that leave them with eight valence electrons.
Learning Objective: Use the octet rule and electron configurations to explain the charge associated with ions.

12. 3.2 Ions and the Octet Rule
When sodium or any other alkali metal reacts with chlorine or any other halogen, the metal transfers an electron from its valence shell to the valence shell of the halogen.

13. 3.2 Ions and the Octet Rule
Main group metals lose electrons to form cations and attain an electron configuration like that of the noble gas just before them in the periodic table.
Main group nonmetals gain electrons to form anions and an electron configuration like that of the noble gas just after them in the periodic table.

14. Worked Example 3.1
Write the electron configuration of magnesium (Z=12). Show how many electrons a magnesium atom must lose to form an ion with a filled shell (eight electrons), and write the configuration of the ion. Explain the reason for the ion’s charge, and write the ion’s symbol. Analysis: Write the electron configuration of magnesium, and count the number of electrons in the valence shell.

15. Worked Example 3.1 Cont.
Solution: Magnesium has the electron configuration 1s22s22p63s2. Because the second shell contains an octet of electrons (2s22p6) and the third shell is only partially filled (3s2), magnesium can achieve a valence-shell octet by losing the two electrons in the 3s2 subshell. The result is the formation of a doubly charged cation, Mg2+, with the neon configuration:

16. Worked Example 3.1 Cont. Solution Continued:
Mg2+: 1s22s22p6 (neon configuration or [Ne])
A neutral magnesium atom has 12 protons and 12 electrons. With the loss of 2 electrons, there is an excess of 2 protons, accounting for the +2 charge of the ion, Mg2+.

17. Worked Example 3.2
How many electrons must a nitrogen atom, Z=7, gain to attain a noble gas configuration? Write the electron dot and ion symbols for the ion formed. Analysis: Write the electron configuration of nitrogen, and identify how many more electrons are needed to reach a noble gas configuration.

18. Worked Example 3.2 Cont. Solution:
Nitrogen, a group 5A element, has the electron configuration 1s22s22p3.
The second shell contains five electrons (2s22p3) and needs three more to reach an octet.
The result is the formation of a triply charged anion, N3–, with eight valence electrons, matching the neon configuration.
N3–: 1s22s22p6 (neon configuration)

19. 3.3 Ions of Some Common Elements
Learning Objective: Use the periodic table to predict the charge associated with ions of main group elements.

20. 3.3 Ions of Some Common Elements
Group 1A: M → M+ + e–
Group 2A: M → M2+ + 2e–
Group 3A: Al → Al3+ + 3e–; no other common ions
Group 4A, 5A: No common ions
Group 6A: X + 2e– → X2–
Group 7A: X + e– → X–
Transition metals form cations, but they can lose one or more d electrons in addition to losing valence s electrons. The octet rule is not followed.

21. 3.3 Ions of Some Common Elements
Metals form cations by losing one or more electrons.
Group 1A and 2A metals form +1 and +2 ions, respectively, to achieve a noble gas configuration.
Transition metals can form cations of more than one charge by losing a combination of valence-shell s electrons and inner-shell d electrons.

22. 3.3 Ions of Some Common Elements
Reactive nonmetals form anions by gaining one or more electrons to achieve a noble gas configuration.
Group 6A nonmetals oxygen and sulfur form the anions O2– and S2–.
Group 7A elements (the halogens) form –1 ions.
Group 8A elements (the noble gases) are unreactive.

23. 3.3 Ions of Some Common Elements
Ionic charges of main group elements can be predicted using the group number and the octet rule.
For 1A and 2A metals, cation charge = group number.
For nonmetals in groups 5A, 6A, and 7A, anion charge = 8 – (group number).

24. Worked Example 3.3
Which of the following ions is likely to form? (a) S3– (b) Si2+ (c) Sr2+ Analysis: Count the number of valence electrons in each ion. For main group elements, only ions with a valence octet of electrons are likely to form.

25. Worked Example 3.3 Cont.
Solution: (a) Sulfur (S) is in group 6A, has six valence electrons, and needs only two more to reach an octet. Gaining two electrons gives an S2– ion with a noble gas configuration, but gaining three electrons does not. The S3– ion is, therefore, unlikely to form.

26. Worked Example 3.3 Cont.
Solution Continued: (b) Silicon (Si) is a nonmetal in group 4A. Like carbon, it does not form ions because it would have to gain or lose too many electrons (four) to reach a noble gas configuration. The Si2+ ion does not have an octet and will not form.

27. Worked Example 3.3 Cont.
Solution Continued: (c) Strontium (Sr) is a metal in group 2A, has only two outer-shell electrons and can lose both to reach a noble gas configuration. The Sr2+ ion has an octet and, therefore, forms easily.

28. 3.4 Periodic Properties and Ion Formation
Metals on the left side of the periodic table tend to lose electrons, whereas nonmetals on the right side of the periodic table tend to gain electrons. How does this trend change as you move across the periodic table? Learning Objective: Explain the formation of anions and cations based on periodic trends.

29. 3.4 Periodic Properties and Ion Formation
Ionization energy is the energy required to remove one electron from a single atom in the gaseous state.
Small values indicate ease of losing electrons to form cations.
Electron affinity is the energy released on adding an electron to a single atom in the gaseous state.
Halogens have the largest values and gain electrons most easily.

30. 3.4 Periodic Properties and Ion Formation
Halogens gain electrons most easily.
Alkali metals lose electrons most easily.
Noble gases neither lose nor gain electrons.
Elements near the middle of the periodic table do not form ions easily.

31. Worked Example 3.4
Predict where the ionization energy of rubidium is likely to fall on the chart below?
Analysis: Identify the group number of rubidium (group 1A), and find where other members of the group appear.

32. Worked Example 3.4 Cont. Solution:
Rubidium (Rb) is the alkali metal below potassium (K) in the periodic table. Since the alkali metals Li, Na, and K all have ionization energies near the bottom of the chart, the ionization energy of rubidium is probably similar.

33. Worked Example 3.5
Which element is likely to lose an electron more easily, Mg or S? Analysis: Identify the group numbers of the elements, and find where members of those groups appear below.

34. Worked Example 3.5 Cont.
Solution: Magnesium, a group 2A element on the left side of the periodic table, has a relatively low ionization energy and loses an electron easily. Sulfur, a group 6A element on the right side of the table, has a higher ionization energy and loses an electron less easily.

35. 3.5 Naming Monoatomic Ions
Learning Objective: Name common monoatomic anions and cations.
Main group metal cations are named by identifying the metal, followed by the word ion.
K Mg Al3+
Potassium ion Magnesium ion Aluminum ion

36. 3.5 Naming Monoatomic Ions
Transition metals can form more than one type of cation. Two naming systems are used.
Old: The ion with the smaller charge is given the ending -ous, and the ion with the larger charge is given the ending -ic.
New: The charge on the ion is given as a roman numeral in parentheses right after the metal name.
Cr Cr3+
Old name: Chromous ion Chromic ion
New name: Chromium (II) ion Chromium (III) ion

37. 3.5 Naming Monoatomic Ions

38. 3.5 Naming Monoatomic Ions
Anions are named by replacing the ending of the element name with -ide, followed by the word ion.

39. 3.6 Polyatomic Ions
Polyatomic ion: An ion that is composed of more than one atom Learning Objective: Identify the name, formula, and charge of common polyatomic ions.

40. 3.6 Polyatomic Ions
The atoms in a polyatomic ion are held together by covalent bonds.
A polyatomic ion is charged because it contains a total number of electrons that is different from the total number of protons in the combined atoms.
These ions are encountered so frequently that it is essential to memorize their names and formulas.

41. 3.6 Polyatomic Ions

42. Chemistry in Action

43. 3.7 Ionic Bonds The arrangement of Na+ and Cl– ions in a sodium
chloride crystal.
The crystal is held together
by ionic bonds.
Learning Objective: Explain the nature of the ionic bonds holding ions together in ionic compounds.

44. 3.7 Ionic Bonds
Ion-transfer reactions of metals and nonmetals form products unlike either element.

45. 3.7 Ionic Bonds
Because opposite electrical charges attract each other, the positive ion and negative ion are said to be held together by an ionic bond.
There are many ions attracted by ionic bonds to their nearest neighbors. These crystals are ionic solids.
Compounds of this type are referred to as ionic compounds.

46. 3.8 Formulas of Ionic Compounds
Since all chemical compounds are neutral, it is easy to figure out the formulas of ionic compounds. Learning Objective: Determine the formula of ionic compounds based on the formulas and charges of the cations and anions.

47. 3.8 Formulas of Ionic Compounds
All chemical compounds are neutral.
Once the ions are identified, decide how many ions of each type give a total charge of zero.
The chemical formula of an ionic compound tells the ratio of anions and cations.

48. 3.8 Formulas of Ionic Compounds
If the ions have the same charge, one of each is needed.
K+ + F– → KF
If the ions have different charges, unequal numbers of anions and cations must combine to have a net charge of zero.
2 K+ + O2– → K2O
Ca Cl– → CaCl2

49. 3.8 Formulas of Ionic Compounds
When the two ions have different charges, the number of one ion is equal to the charge on the other ion.

50. 3.8 Formulas of Ionic Compounds
The formula of an ionic compound shows the lowest possible ratio of atoms and is known as a simplest formula.
There is no such thing as a single neutral particle of an ionic compound.

51. 3.8 Formulas of Ionic Compounds
Formula unit: The formula that identifies the smallest neutral unit of an ionic compound
For NaCl, the formula unit is one Na+ ion and one Cl– ion.
For CaF2, the formula unit is one Ca2+ ion and two F– ions.

52. 3.8 Formulas of Ionic Compounds

53. 3.8 Formulas of Ionic Compounds
Once the numbers and kinds of ions in a compound are known, the formula is written using the following rules:
List the cation first and the anion second.
Do not write the charges of the ions.
Use parentheses around a polyatomic ion formula if it has a subscript.

54. Worked Example 3.6
Write the formula for the compound formed by calcium ions and nitrate ions. Analysis: Knowing the formula and charges on the cation and anion, we determine how many of each are needed to yield a neutral formula for the ionic compound.

55. Worked Example 3.6 Cont.
Solution: The two ions are Ca2+ and NO3–. Two nitrate ions, each with a –1 charge, will balance the +2 charge of the calcium ion. Ca2+ Charge = 1 × (+2) = +2 NO3– Charge = 2 × (–1) = –2 Because there are two ions, the nitrate formula must be enclosed in parentheses. Ca(NO3)2 Calcium nitrate

56. 3.9 Naming Ionic Compounds
Ionic compounds are named by citing first the cation and then the anion, with a space between words. Learning Objective: Based on the cations and anions in the chemical formula, write the name of ionic compounds.

57. 3.9 Naming Ionic Compounds
There are two kinds of ionic compounds:
Type I ionic compounds contain cations of main group elements.
Type II ionic compounds contain metals that can exhibit more than one charge.
These require different naming conventions.

58. 3.9 Naming Ionic Compounds
Type I ionic compounds contain cations of main group elements.
The charges on these cations do not vary.
Do not specify the charge on the cation.
NaCl is sodium chloride.
MgCO3 is magnesium carbonate.

59. 3.9 Naming Ionic Compounds
Type II ionic compounds contain metals that can exhibit more than one charge.
Specify the charge on the cation in these compounds with either the old (-ous, -ic) or the new (roman numerals) system.
FeCl2 is iron(II) chloride or ferrous chloride.
FeCl3 is iron(III) chloride or ferric chloride.

60. 3.9 Naming Ionic Compounds
Type II ionic compounds continued:
Do not name FeCl2 iron dichloride or FeCl3 iron trichloride.
Once the charge on the metal is known, the number of anions needed to yield a neutral compound is also known.
Charges do not need to be included as part of the compound name.

61. 3.9 Naming Ionic Compounds

62. Worked Example 3.7
Sodium and calcium both form a wide variety of ionic compounds. Write formulas for the following compounds.
Sodium bromide and calcium bromide
Sodium sulfide and calcium sulfide
Sodium sulfate and calcium sulfate
Sodium phosphate and calcium phosphate

63. Worked Example 3.7 Cont.
Analysis: Using the formulas and charges for the cations and anions, we determine how many of each cation and anion are needed to yield a formula that is neutral.
Solution:
The cations are sodium (Na) and calcium (Ca). Sodium is a 1A metal and would have a +1 charge (Na+); calcium, as a 2A metal, would have a +2 charge (Ca2+).

64. Worked Example 3.7 Cont. Solution Continued:
The anion is in group 7A and would have a –1 charge (Br–). For the charges to be neutral, the respective compounds are NaBr and CaBr2.
The anion is in group 6A and would have a –2 charge (S2–). The neutral formulas would be Na2S and CaS.

65. Worked Example 3.7 Cont. Solution Continued:
(c) The anion is a polyatomic ion with a –2 charge (SO42–). The neutral formulas would be Na2SO4 and CaSO4.
(d) The anion is a polyatomic ion with a –3 charge (PO43–). The neutral formulas would be Na3PO4 and Ca3(PO4)2.

66. Worked Example 3.8
Name the following compounds using roman numerals to indicate the charges on the cations where necessary. (a) KF (b) MgCl2 (c) AuCl3 (d) Fe2O3 Analysis: For main group metals, the charge is determined from the group number, and no roman numerals are necessary. For transition metals, the charge on the metal can be determined from the total charge(s) on the anion(s).

67. Worked Example 3.8 Cont. Solution: Potassium fluoride
Magnesium chloride
Gold(III) chloride
Iron(III) oxide. Because the three oxide anions (O2–) have a total negative charge of –6, the two iron cations must have a total charge of +6. Thus, each is Fe3+.

68. 3.10 Some Properties of Ionic Compounds
The melting point of sodium chloride is 801 °C.
Learning Objective: Identify the properties of ionic compounds, and explain how the nature of the ionic bond is reflected in the physical properties of ionic compounds.

69. 3.10 Some Properties of Ionic Compounds
Ions in each compound settle into a pattern that efficiently fills space and maximizes ionic bonding.
Ions in an ionic solid are held rigidly in place by attraction to their neighbors.
Once an ionic solid is dissolved in water, the ions can move freely, which accounts for the electrical conductivity of these compounds in solution.

70. 3.10 Some Properties of Ionic Compounds
Ionic compounds have very high melting and boiling points. Sodium chloride melts at 801 °C and boils at 1413 °C.
Ionic solids shatter if struck sharply.
Ionic compounds dissolve in water if the attraction between water and the ions overcomes the attraction of the ions for one another. Not all ionic compounds are water soluble.

71. They are fundamental to the concepts of acids and bases.
3.11 H+ and OH– Ions: An Introduction to Acids and Bases
Two of the most important ions are the hydrogen cation (H+) and the hydroxide anion (OH–).
They are fundamental to the concepts of acids and bases.
Learning Objective: Identify the ions associated with acids and bases.

72. Acid: A substance that provides H+ ions in water
3.11 H+ and OH– Ions: An Introduction to Acids and Bases
Acid: A substance that provides H+ ions in water
A hydrogen cation is simply a proton.
When an acid dissolves in water, the proton attaches to a molecule of water to form a hydronium ion.
H+ + H2O → H3O+
Chemists use hydrogen and hydronium ions interchangeably.

73. Hydrochloric acid, HCl, provides one H+ ion per acid molecule.
3.11 H+ and OH– Ions: An Introduction to Acids and Bases
Different acids can provide different numbers of H+ ions per acid molecule.
Hydrochloric acid, HCl, provides one H+ ion per acid molecule.
Sulfuric acid, H2SO4, can provide two H+ ions per acid molecule.
Phosphoric acid, H3PO4, can provide three H+ ions per acid molecule.
© 2017 Pearson Education, Inc.

74. Base: A substance that provides OH– ions in water
3.11 H+ and OH– Ions: An Introduction to Acids and Bases
Base: A substance that provides OH– ions in water
A hydroxide anion is a polyatomic ion in which an oxygen atom is covalently bonded to a hydrogen atom.

75. Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are bases.
3.11 H+ and OH– Ions: An Introduction to Acids and Bases
Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are bases.
When these compounds dissolve, OH– anions go into solution along with the metal cation.
Different bases can provide different numbers of OH– ions per formula unit.
Sodium hydroxide provides one OH– ion per formula unit.
Barium hydroxide, Ba(OH)2, can provide two OH– ions per formula unit.

76. 3.11 H+ and OH– Ions: An Introduction to Acids and Bases

77. Concept Map