Chapter 6. Valence Electrons

Transcription

1 Chapter 6 Ionic Bonds & Some Main-Group Chemistry Chapter 6 1 Valence Electrons When an atom undergoes a chemical reaction, only the outermost electrons are involved. These electrons are of the highest energy and are furthest away from the nucleus. These are the valence electrons. For the main group elements, the valence electrons are the s and p electrons beyond the noble gas core. Chapter 6 2 1

2 Predicting Valence Electrons For the main group elements, the Group number indicates the number of valence electrons. Chapter 6 3 Ion Electron Configurations When we write the electron configuration of a cation, we remove one electron for each positive charge: Na Na + 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 When we write the electron configuration of an anion, we add one electron for each negative charge: O O 2-1s 2 2s 2 2p 4 1s 2 2s 2 2p 6 Chapter 6 4 2

3 Ion Electron Configuration: Transition Metals Transition metals also lose valence electrons when forming cations However, these elements lose their valence-shell s- electrons before losing their d-electrons. Electrons with the highest n-quantum number are lost first. Fe Fe 2+ [Ar] 4s 2 3d 6 [Ar] 3d 6 Never [Ar] 4s 2 3d 6 [Ar] 4s 2 3d 4 Chapter 6 5 Isoelectronic Ions Isoelectronic atoms and ions are those which have the same number of electrons For example: A Sodium ion has 10 electrons: Na Na + 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 A Fluorine ion also has 10 electrons: F F - 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6 Chapter 6 6 3

4 Ions and Ionic Radii Cations are smaller than their parent atom. The size decreases because of a decrease in valence shell size, causing an increase in Z eff Anions are larger than their parent atom. The size increase is due to an increase the Z eff and in electronelectron repulsion Chapter 6 7 Ionization Energy The ionization energy (E I ) of an atom is the amount of energy required to remove an electron in the gaseous state (units of kj/mol). The closer the electron is to the nucleus, the harder it is to remove. Core electrons shield valence electrons effectively Electrons in the same shell shield one another less effectively Chapter 6 8 4

5 There are some irregularities in the E I trend Why do you think that Be has a higher E I than Boron? Ionization Energy There are additional Ionization Energies, depending on the number of the electron you are trying to remove. The values for these ionization energies depend greatly on the location of the electron in the orbitals, so use your configurations! Chapter 6 9 Electron Affinity The electron affinity (E ea ) of an atom is the amount of energy released when an electron is added to an atom in the gaseous state (units of kj/mol). Because the energy is released, electron affinities have negative values. The more negative the value, the higher the electron affinity! The value of an atom s E ea is due to a balance between: 1) the attraction between the new electron and the nucleus 2) the increase in electronelectron repulsions Chapter

6 Ionic Bonds and Solids An ionic bond is formed when an element with a small E I value comes in contact with an element with a negative E ea value. The oppositely charged ions that are formed by this interaction are then attracted to one another by electrostatic forces and are joined by an ionic bond. This electrostatic force is similar to the attraction between opposite poles on two magnets. Ionic bonds are very strong and result in the formation of a rigid, crystalline structure called an ionic solid. Chapter 6 11 The bonds in an ionic compounds are very strong which is why most ionic compounds are solids at room temperature These compounds have very high melting points In their solid state, these compounds are not good conductors of electricity However, when the substance is melted or dissolved, the crystal is destroyed, allowing the ions to roam free and conduct electricity. Ionic Compounds Chapter

7 Lattice Energy (U) The lattice energy (U) is the sum of the electrostatic interaction energies between the ions in a crystal. The LE is listed as the amount of energy needed to breakup a crystal and has a positive value If you want to know the amount of energy given off when a lattice forms, just change the sign to negative! The lattice energy is dependent on two factors: The charges on the ions (z 1 and z 2 ) The distance between the ions (d) U = k z 1 z 2 d Chapter 6 13 Lattice Energy (U) One of the below pictures represents NaCl and the other is MgO. Which is which? Which of the two crystals do you think has the higher lattice energy? Chapter

8 Energetics of Ionic Solid Formation For a reaction to proceed, the overall net energy change ( E) must be negative Merely looking at the E I and E ea of two particular elements is not enough. Too simple. For Example: NaCl E I for Na = kj/mo E ea for Cl = kj/mol E = kj/mol (No Go!) The total energetics of ionic reactions can be viewed on a Born Haber Cycle which shows how each step contributes to the overall reaction energy. Chapter 6 15 Energetics of Ionic Solid Formation E I E ea Lattice Energy (U) Chapter

9 Energetics of Ionic Solid Formation Chapter 6 17 Energetics of Ionic Solid Formation Calculate the overall energy change (in kj/mol) for the formation of CaCl from its elements and the formation of CaCl 2 from its elements. E a for Cl = kj/mol E I1 for Ca = kj/mol E I2 for Ca = kj/mol Heat of Sublimation for Ca = kj/mol Bond Dissociation Energy for Cl 2 = +243 kj/mol Lattice Energy for CaCl 2 = kj/mol Lattice Energy for CaCl = +717 kj/mol Which is more likely to form, CaCl or CaCl 2? Chapter

10 Chemical Families Elements within a group (a column) have similar chemical properties Several of these groups form chemical families Trends among these families are most obvious for the main group elements. Chapter 6 19 Group 1A: Alkali Metals With the exception of hydrogen, all the elements in group 1A are known as the Alkali metals These elements all have one valence s electron, and like to form +1 ions This makes them excellent reducing agents As you move down the group, the reactivity of the metal increases Alkali metals do not look like what we think of as a metal (shiny) because they react readily with oxygen which coats their outer layer. These elements are good conductors of electricity and can be readily formed into foils and wires. Chapter

11 Group 1A: Alkali Metals Reaction with Halogens 2 M(s) + X 2 2 MX(s) Reaction with Oxygen Forms oxide (Li 2 O), peroxide (Na 2 O 2 ), or superoxide (KO 2 ) Reaction with Hydrogen 2 M(s) + H 2 2 MH(s) Reaction with Nitrogen 6 Li(s) + N 2 2 Li 3 N(s) Chapter 6 21 Group 2A: Alkali Metals The elements in group 2A are known as the Alkaline Earth metals These elements all have two valence s electrons, and like to form +2 ions This makes them excellent reducing agents As you move down the group, the reactivity of the metal increases These elements are good conductors of electricity and can be readily formed into foils and wires. Compared to Group 1A metals, these metals are smaller, have higher melting and boiling points and are less reactive. Chapter

12 Group 2A: Alkali Metals Reaction with Halogens M(s) + X 2 MX 2 (s) Reaction with Oxygen M(s) + O 2 MO(s) Reaction with Hydrogen 2 Ca(s) + H 2 2 CaH 2 (s) Reaction with Water (only Ba is vigorous) Ba(s) + H 2 O Ba 2+ (aq) + 2 OH (aq) + H 2 (s) Chapter 6 23 Group 3A: Boron Elements in Group 3A contain one semimetal that acts as a nonmetal, and four that are primarily metallic. Boron is acts as a nonmetal and forms covalent bonds. It has many similarities to carbon and silicon. Its first I.E. is too high (801 kj/mol) for electron loss so it will not form an ion Boron forms covalent compounds not ionic. Always trivalent and never monovalent. Boron can have an incomplete octet (does not follow the octet rule!) Elemental boron is a good semiconductor. Chapter

13 Group 3A: Aluminum Alumnium is the most abundant metal in the earth s crust at 8.3% Similar to Group 1A and 2A metals, Al is a reducing agent that undergoes redox reactions by losing all three of its valence electrons ( 1 p and 2 s electrons) to yield Al 3+ ions. Aluminum is less reactive than the Group 1A and 2A metals Will react vigorously with O 2 : Main oxide is Al 2 O 3 Will react vigorously with halogens: Main halides is AlX 3 but can be more complex Chapter 6 25 Group 7A: The Halogens The elements in group 7A are known as the Halogens These elements are reactive and toxic non-metals The Halogens have seven valence electrons, and like to form -1 ions This makes them excellent oxidizing agents As you move down the group, the reactivity of the non-metal decreases These elements exist as solids (I), liquids (Br) or gases (F and Cl) in their native state. The density, melting point and boiling point all increase as you move down the group. Chapter

14 Group 7A: The Halogens Metal Halides: Halogens react with every metal in the periodic table to form metal halides. 2 M + n X 2 2 MX n Hydrogen Halides: They are formed from salt and sulfuric acid: 2 NaCl(s) + H 2 SO 4 (aq) 2 HCl(g) + Na 2 SO 4 (aq) HBr is usually formed from PBr 3 and H 2 O: PBr 3 (l) + 3 H 2 O(l) 3 HBr(g) + H 3 PO 3 (aq) Reactivity decreases down the group! Chapter 6 27 Group 8A: The Nobel Gases The elements in group 8A are known as the Nobel Gases These elements are non-metals and have eight valence electrons Because they have eight valence electrons, these elements do not like to undergo redox reactions at all. These elements exist as gases in their native state. The density, melting point and boiling point all increase as you move down the group. Reactivity He, Ne and Ar undergo no chemical reactions and form no known compounds. Kr and Xe react only with Fluorine. For example, XeF 2, XeF 4 and XeF 6 are all very powerful oxidizing agents. Chapter

15 Octet Rule The octet rule states that main group elements tend to undergo reactions that leave them with eight outer shell electrons. Chapter 6 29 Octet Rule Exceptions There are some exceptions to the octet rule. Group 3A non-metals (Boron) can be electron deficient Elements in Period 3 or higher and Group 3A or higher can be electron rich Chapter