Chemical Reactivity

1. 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

2. 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 8 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.

3. 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.

4. 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

5. Section1 Simple Ions Chapter 5 Chemical Reactivity, continued Alkali Metals and Halogens Are the Most Reactive Elements, continued • Chlorine, 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

6. 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.

7. 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.

8. 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.

9. 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.

10. 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.

11. 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.

12. 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.

13. 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.

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

15. 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.

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

17. 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.

18. 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.

19. 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.

20. Visual Concepts Chapter 5 Ionic Bonding

21. 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.

22. 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.

23. Formation of Sodium Chloride Section2 Ionic Bonding and Salts Chapter 5

24. 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.

25. Visual Concepts Chapter 5 Lattice Energy

26. 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.

27. 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.

28. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued Ionic Bonds Are Strong • Repulsive forces in a salt crystal include those between like-charged ions. • Each Na+ ion repels the other Na+ ions. Each Cl ion repels the other Cl ions. • Another repulsive force exists between the electrons of ions that are close together. • Attractive forces include those between the positive nucleus of one ion and electrons of other ions.

29. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued Ionic Bonds Are Strong, continued • Attractive forces exist between oppositely charged ions and involve more than a single cation and anion. • Six Na+ ions surround each Cl ion and vice versa. • As a result, the attractive force between oppositely charged ions is significantly greater in a crystal than it would be if the ions existed only in pairs. • Overall, the attractive forces are much stronger than the repulsive ones, so ionic bonds are strong.

30. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued Ionic Compounds Have Distinctive Properties • Most ionic compounds have high melting and boiling points because of the strong attraction between ions. • To melt, ions cannot be in fixed locations. • Because the bonds between ions are strong, a lot of energy is needed to free them. • Still more energy is needed to move ions out of the liquid state and cause boiling, so ionic compounds are rarely gaseous at room temperature.

31. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued Ionic Compounds Have Distinctive Properties, continued

32. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued Liquid and Dissolved Salts Conduct Electric Current • To conduct an electric current, a substance must satisfy two conditions: • it must contain charged particles • those particles must be free to move • Ionic solids, such as salts, generally are not conductors because the ions cannot move. • When a salt melts or dissolves,the ions can move about and are excellent electrical conductors.

33. Sodium Chloride in Three Phases Section2 Ionic Bonding and Salts Chapter 5

34. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued Salts Are Hard and Brittle • Like NaCl, most ionic compounds are hard and brittle. • Hardmeans that the crystal is able to resist a large force applied to it. • Brittlemeans that when the applied force becomes too strong to resist, the crystal develops a widespread fracture rather than a small dent. • Both properties are due to the patterns in which the cations and anions are arranged in all salt crystals.

35. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued Salts Are Hard and Brittle, continued • The ions in a crystal are arranged in a repeating pattern, forming layers. • Each layer is positioned so that a cation is next to an anion in the next layer. The attractive forces between opposite charges resist motion. • As a result, the ionic compound will be hard. • Also, it will take a lot of energy to break all the bonds between layers of ions.

36. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued Salts Are Hard and Brittle, continued • If a force causes one layer to move, ions of the same charge will be positioned next to each other. • The cations in one layer are now lined up with other cations in a nearby layer. The anions are also. • Because like charges are next to each other, they will repel each other and the layers will split apart. • This is why all salts shatter along a line extending through the crystal known as a cleavage plane.

37. Section2 Ionic Bonding and Salts Chapter 5 Ionic Compounds, continued How to Identify a Compound as Ionic • All ionic compounds are solid at room temperature. • Tap the substance. • Ionic compounds do not break apart easily and they fracture into tiny crystals. • Heat the substance. • Ionic compounds generally have high melting and boiling points. • Use a conductivity device to find if the dissolved or melted substance conducts electricity. • Dissolved and molten ionic compounds conduct electricity.

38. Section2 Ionic Bonding and Salts Chapter 5 Salt Crystals • Despite their differences, the crystals of all salts are made of simple repeating units. • These repeating units are arranged in a salt to form a crystal lattice, the regular pattern in which a crystal is arranged. • These repeating patterns within a salt are the reason for the crystal shape that can be seen in most salts.

39. Section2 Ionic Bonding and Salts Chapter 5 Salt Crystals, continued Crystal Structure Depends on the Sizes and Ratios of Ions • Formulas indicate ratios of ions. • For example, the formula for NaCl indicates there is a 1:1 ratio of sodium cations and chlorine anions. • Within a NaCl crystal, each Na+ ion is surrounded by six Cl ions, and each Cl ion by six Na+ ions. • Because the edges of the crystal do not have this arrangement, they are locations of weak points.