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2. Resources Chapter Presentation Bellringer Transparencies Sample Problems Visual Concepts Standardized Test Prep

3. Ions and Ionic Compounds Chapter 5 Table of Contents Section 1Simple Ions Section 2Ionic Bonding and Salts Section 3Names and Formulas of Ionic Compounds

5. Section1 Simple Ions Chapter 5 Bellringer • Review the properties and atomic structures of the noble gas (Group 18) elements. • What do the electron configurations of these elements have in common?

6. Section1 Simple Ions Chapter 5 Objectives • Relate the electron configuration of an atom to its chemical reactivity. • Determine an atom’s number of valence electrons, and use the octet rule to predict what stable ions the atom is likely to form. • Explain why the properties of ions differ from those of their parent atoms.

7. Section1 Simple Ions Chapter 5 Chemical Reactivity • How much an element reacts depends on the electron configuration of its atoms. • For example, oxygen will react with magnesium. In the electron configuration for oxygen, the 2p orbitals, which can hold six electrons, have only four: • [O] = 1s22s22p4 • Neon has no reactivity. Its 2p orbitals are full: • [Ne] = 1s22s22p6

8. Section1 Simple Ions Chapter 5 Chemical Reactivity, continued Noble Gases Are the Least Reactive Elements • Neon is a noble gas. • The noble gases, which are found in Group 18 of the periodic table, show almost no chemical reactivity. • The noble gases have filled outer energy levels. • This electron configuration can be written as ns2np6 where n represents the outer energy level.

9. Section1 Simple Ions Chapter 5 Chemical Reactivity, continued Noble Gases Are the Least Reactive Elements, continued • The eight electrons in the outer energy level fill the s and p orbitals, making these noble gases stable. • In most chemical reactions, atoms tend to match the s and p electron configurations of the noble gases. • This tendency to have either empty outer energy levels or full outer energy levelsof eight electrons is called the octet rule.

10. Visual Concepts Chapter 5 The Octet Rule

11. Section1 Simple Ions Chapter 5 Chemical Reactivity, continued Alkali Metals and Halogens Are the Most Reactive Elements • An atom whose outer s and p orbitals do not match the electron configurations of a noble gas will react to lose or gain electrons so the outer orbitals will be full. • When added to water, an atom of potassium (an alkali metal) gives up one electron in its outer energy level. • Then, it has the s and p configuration of a noble gas. • 1s22s22p63s23p64s1 1s22s22p63s23p6

12. Section1 Simple Ions Chapter 5 Chemical Reactivity, continued Alkali Metals and Halogens Are the Most Reactive Elements, continued • Chlolrine, a halogen, is also very reactive. • An atom of chlorine has seven electrons in its outer energy level. • By gaining just one electron, it will have the s and p configuration of a noble gas. • 1s22s22p63s23p5 1s22s22p63s23p6

13. Section1 Simple Ions Chapter 5 Valence Electrons • Potassium after it loses one electron has the same electron configurationas chlorine after it gains one. • Both are the same as that of the noble gas argon. • [Ar] = 1s22s22p63s23p6 • The atoms of many elements become stable by achieving the electron configuration of a noble gas. • The electrons in the outer energy level are known as valence electrons.

14. Visual Concepts Chapter 5 Valence Electrons

15. Section1 Simple Ions Chapter 5 Valence Electrons, continued Periodic Table Reveals an Atom’s Number of Valence Electrons • To find out how many valence electrons an atom has, check the periodic table. • For example, the element magnesium, Mg, has the following electron configuration: • [Mg] = [Ne]3s2 • This configuration shows that a magnesium atom has two valence electrons in the 3s orbital.

16. Section1 Simple Ions Chapter 5 Valence Electrons, continued Periodic Table Reveals an Atom’s Number of Valence Electrons, continued • The electron configuration of phosphorus, P, is [Ne]3s23p3. • Each P atom has five valence electrons: two in the 3s orbital and three in the 3p orbital.

17. Section1 Simple Ions Chapter 5 Valence Electrons, continued Atoms Gain Or Lose Electrons to Form Stable Ions • All atoms are uncharged because they have equal numbers of protons and electrons. • For example, a potassium atom has 19 protons and 19 electrons. • After giving up one electron, potassium still has 19 protons but only 18 electrons. • Because the numbers are not the same, there is a net electrical charge.

18. Section1 Simple Ions Chapter 5 Valence Electrons, continued Atoms Gain Or Lose Electrons to Form Stable Ions, continued • An ion is an atom, radical, or molecule that has gained or lost one or more electrons and has a negative or positive charge. • The following equation shows how a potassium atom forms an ion with a 1+ charge. • K  K+ + e • An ion with a positive charge is called a cation.

19. Visual Concepts Chapter 5 Ion

20. Section1 Simple Ions Chapter 5 Valence Electrons, continued Atoms Gain Or Lose Electrons to Form Stable Ions, continued • In the case of chlorine, far less energy is required for an atom to gain one electron rather than give up its seven valence electrons to be more stable. • The following equation shows how a chlorine atom forms an ion with a 1− charge. • Cl + e → Cl • An ion with a negative charge is called an anion.

21. Section1 Simple Ions Chapter 5 Valence Electrons, continued Characteristics of Stable Ions • Both an atom and its ion have the same number of protons and neutrons, so the nuclei are the same. • The chemical properties of an atom depend on the number and configuration of its electrons. • Therefore, an atom and its ion have different chemical properties.

22. Section1 Simple Ions Chapter 5 Valence Electrons, continued Many Stable Ions Have Noble-Gas Configurations • Many atoms can form stable ions with a full octet. For example, Ca, forms a stable ion. • The electron configuration of a calcium atom is: • [Ca] = 1s22s22p63s23p64s2 • By giving up its two valence electrons in the 4s orbital, it forms a stable cation with a 2+ charge: • [Ca2+] = 1s22s22p63s23p6 • This electron configuration is like that of argon.

23. Some Ions with Noble-Gas Configurations Section1 Simple Ions Chapter 5

24. Section1 Simple Ions Chapter 5 Valence Electrons, continued Some Stable Ions Do Not Have Noble-Gas Configurations • Not all stable ions have an electron configuration like those of noble gases. Transition metals often form ions without complete octets. • With the lone exception of rhenium, Re, the stable transition metal ions are all cations. • Also, some elements, mostly transition metals, form stable ions with more than one charge.

25. Stable Ions Formed by the Transition Elements and Some Other Metals Section1 Simple Ions Chapter 5

26. Section1 Simple Ions Chapter 5 Atoms and Ions • Having identical electron configurations does not mean that a sodium cation is a neon atom. • They still have different numbers of protons and neutrons. Ions and Their Parent Atoms Have Different Properties • Both sodium and chlorine are very reactive. • When they are mixed, a violent reaction takes place, producing a white solid—table salt (sodium chloride). • It is made from sodium cations and chloride anions.

27. Section1 Simple Ions Chapter 5 Atoms and Ions, continued Atoms of Metals and Nonmetal Elements Form Ions Differently • Nearly all metals form cations. For example, magnesium metal, Mg, has the electron configuration: • [Mg] = 1s22s22p63s2 • To have a noble-gas configuration, the atom must either gain six electrons or lose two. • Losing two electrons requires less energy than gaining six.

28. Section1 Simple Ions Chapter 5 Atoms and Ions, continued Atoms of Metals and Nonmetal Elements Form Ions Differently, continued • The atoms of all nonmetal elements form anions. For example, oxygen, O, has the electron configuration: • [O] = 1s22s22p4 • To have a noble-gas configuration, an oxygen atom must either gain two electrons or lose six. • Acquiring two electrons requires less energy than losing six.

29. 10. How could each of the following atoms react to achieve a noble-gas configuration? a. iodine b. strontium c. nitrogen d. krypton

30. 11. Write the electron configuration for each of the following ions. a. Al3+ b. Se2− c. Sc3+ d. As3−

31. HOMEWORK SECTION 5.1 REVIEW Pg.165 Q 1 13

32. 1. Explain why the noble gases tend not to react. 2. Where are the valence electrons located in an atom? 3. How does a cation differ from an anion?

33. 4. State the octet rule. 5. Why do the properties of an ion differ from those of its parent atom? 6. Explain why alkali metals are extremely reactive.

34. 7. How can you determine the number of valence electrons an atom has? 8. Explain why almost all metals tend to form cations. 9. Explain why, as a pure element, oxygen is usually found in nature as O2.

35. 12. In what way is an ion the same as its parent atom? 13. To achieve a noble-gas configuration, aphosphorus atom will form a P3− anion rather than forming a P5+ cation.Why?

36. Section2 Ionic Bonding and Salts Chapter 5 Bellringer • Review the concepts of ionization energy, electron affinity, anion formation, cation formation, and vaporization. • Write down whether each of these processes involves the gain or loss of energy.

37. Section2 Ionic Bonding and Salts Chapter 5 Objectives • Describe the process of forming an ionic bond. • Explain how the properties of ionic compounds depend on the nature of ionic bonds. • Describe the structure of salt crystals.

38. Section2 Ionic Bonding and Salts Chapter 5 Ionic Bonding • Pyrite is a mineral that is shiny like gold, but it is made of iron cations and sulfur anions. • Because opposite charges attract, cations and anions attract one another and an ionic bond is formed. • The iron cations and sulfur anions of pyrite attract one another to form an ionic compound.

39. Visual Concepts Chapter 5 Ionic Bonding

40. Section2 Ionic Bonding and Salts Chapter 5 Ionic Bonding, continued Ionic Bonds Form Between Ions of Opposite Charge • When sodium and chlorine react to form sodium chloride, sodium forms a stable Na+ cation and chlorine forms a stable Cl anion. • The force of attraction between the 1+ charge on the sodium cation and the 1 charge on the chloride anion creates the ionic bond in sodium chloride. • Sodium chloride is a salt, the scientific name given to many different ionic compounds.

41. Visual Concepts Chapter 5 Salt

42. Section2 Ionic Bonding and Salts Chapter 5 Ionic Bonding, continued Ionic Bonds Form Between Ions of Opposite Charge, continued • All salts are electrically neutralionic compounds that are made up of cations and anions held together by ionic bonds in a simple, whole-number ratio. • However, the attractions between the ions in a salt do not stop with a single cation and a single anion. • One cation attracts several anions, and one anion attracts several cations. • They are all pulled together into a tightly packed crystal structure.

43. Visual Concepts Chapter 5 Characteristics of Ion Bonding in a Crystal Lattice

44. Section2 Ionic Bonding and Salts Chapter 5 Ionic Bonding, continued Transferring Electrons Involves Energy Changes • Ionization energy is the energy that it takes to remove the outermost electron from an atom. • The equation below shows this process for sodium. • Na + energy  Na+ + e • With some elements, such as chlorine, energy is released when an electron is added. • Cl + e Cl + energy

45. Section2 Ionic Bonding and Salts Chapter 5 Ionic Bonding, continued Transferring Electrons Involves Energy Changes, continued • The energy released when chlorine accepts and electron is less than the energy required to remove an electron from a sodium atom. • Adding and removing electrons is only part of forming an ionic bond. • The rest of the process of forming a salt supplies enough energy to make up the difference so that the overall process releases energy.

46. Section2 Ionic Bonding and Salts Chapter 5 Ionic Bonding, continued Salt Formation Also Involves Exothermic Steps, continued • The energy released when ionic bonds are formed is called the lattice energy. • This energy is released when the crystal structure of a salt is formed as the separated ions bond. • Without this energy, there would not be enough energy to make the overall process spontaneous.

47. Visual Concepts Chapter 5 Lattice Energy

48. Section2 Ionic Bonding and Salts Chapter 5 Ionic Bonding, continued Salt Formation Also Involves Exothermic Steps, continued • If energy is released when ionic bonds are formed, then energy must be supplied to break these bonds. • As sodium chloride dissolves in water, water supplies energy for the Na+ and Cl ions to separate. • Because of its much higher lattice energy, magnesium oxide does not dissolve well in water. • There is not enough energy to separate the Mg2+and O2 ions from one another.

49. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds • The ratio of cations to anions is always such that an ionic compound has no overall charge. Ionic Compounds Do Not Consist of Molecules • Water is a molecular compound, so individual water molecules are each made of two hydrogen atoms and one oxygen atom. • Sodium chloride is an ionic compound, so it is made up of many Na+ and Cl ions all bonded together to form a crystal. There are no NaCl molecules.

50. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued Ionic Compounds Do Not Consist of Molecules, continued • Metals and nonmetals tend to form ionic compounds and not molecular compounds. • The formula CaO likely indicates an ionic compound because Ca is a metal and O is a nonmetal. • In contrast, the formula ICl likely indicates a molecular compound because both I and Cl are nonmetals. • Lab tests are used to confirm such indications.