1. Unit 3.1: Ionic and Bonds

2. Vocabulary: Valence electrons: electrons in the outermost energy shell that are available to participate in bonding. Electron dot structure: a diagram that represents the valence electrons of an element as dots; also called Lewis dot structures. Octet rule: rule that states that a valence shell is most stable with 8 electrons. Cation: an element that loses one or more electrons and becomes positively-charged. Anion: an element that gains one or more electrons and becomes negatively charged. Ionic bond: bond formed when an anion interacts with a cation; Formula unit: lowest whole-number ratio of ions in an ionic compound. Metallic bond: bond that forms from the attraction between free- floating valence electrons and positively-charged metal ions. 2

3. I. Formation of Ions (7.1) A.All elements within a group have similar properties because they have same # of valence electrons. 1.# of valence electrons = group # a.Ex: Na is in group IA, has one valence electron in 3s 1 orbital; C is in group IVA, has four valence electrons: 2s 2 2p 2 2.Valence electrons can participate in bonding. 3

4. B.Electron dot structures: represent valence e-’s as dots, one dot for each valence e-. 1.All elements within a group have identical dot structures (except for Helium) a.One dot in each “orbital” before adding second dot to same orbital 2.Also called Lewis dot structures 4

5. C.Atoms are most stable when valence shell contains eight electrons; known as the “octet rule” 1.Achieves the electron configuration of nearest noble gas. 5

6. D.Metals tend to lose valence electrons, achieve noble gas configuration of next-lowest energy level 1.But now less e-’s than p+’s, so positively charged a.Called a ca t ion: “t” in ca t ion looks like a “+” charge b.Cation “keeps its name”; is named for its parent element: Ex: sodium atom becomes “sodium ion” 6

7. 2.Group IA and IIA metals, cation charge = group # of the parent atom: a.Group IA elements always have a +1 charge. b.Group IIA elements always have a +2 charge. 3.For transition metals, cation charge may vary: a.Ex: Iron (Fe) can lose either 2 or 3 electrons: Fe = [Ar]4s 2 3d 6 ; if it loses the 4s 2 electrons, becomes Fe 2+ But not as stable b/c still has 3d 6 ; if loses 1 more electron, becomes 3d 5 which has one electron in each of the 5 d sub- orbitals – more stable, becomes Fe 3+ 7

8. 4.Other transition metals achieve a pseudo noble- gas configuration b/c would have to gain or lose too many electrons to be stable. a.Ex: Zn (1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 ) would need to lose 12 electrons or gain 6 electrons to achieve noble-gas configuration, but gain/loss of more than 3 electrons is unstable. b.If Zn loses both of the 4s 2 electrons, then it doesn’t have any electrons in that energy level, achieves stability. c.Becomes Zn 2+ 8

9. E.Most nonmetals gain valence electrons, achieve noble gas configuration of same energy level 1.But now more e-’s than p+’s, so negatively charged a.Called an anion: “n” in “anion”, “n” in “negative” b.Anion drops the ending of parent element, replaces ending with “-ide”: Ex: chlorine atom becomes “chloride” 9

10. 2.Halogens in Group VIIA become “halides” a.Group VII elements have 7 valence e-’s, need only one more electron to achieve octet (8) b.If halogens gain one more e-, carry charge of -1 1)Ex: Fluorine (F) gains one e-, becomes F-, called “fluoride” 2)Ex: Chlorine (Cl) becomes Cl-, called “chloride” 3.Nonmetals in Group VIA (O, S, Se) a.Have 6 valence e-’s, need only 2 more to achieve octet (8) b.If they gain 2 e-’s, carry charge of -2 1)Ex: Oxygen (O) gains 2 e-’s, becomes O 2-, called “oxide” 2)Ex: Sulfur (S) gains 2 e-’s, becomes S 2-, called “sulfide” 10

11. II.Ionic bonds form between metals, nonmetals A.Ionic compounds form when an anion bonds to a cation through electrostatic (+/-) forces. 1.Metal (cation) “gives up” e - ’s to nonmetal (anion) B.Ionic compounds are electrically neutral because total anionic charge cancels total cationic charge. 1.Ex: Na+ bonds with Cl- in a 1:1 ratio, becomes NaCl 2.Ex: Mg 2+ bonds with O 2- in a 1:1 ratio, becomes MgO 3.Ex: Mg 2+ bonds to 2 Cl- in a 1:2 ratio, becomes MgCl 2 11

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13. C.Chemical formula: describes the number of each type of atom participating in the bond: 1.If there is no subscript, then there is just 1 of that type of atom participating: a.Ex: NaCl = 1 sodium ion (Na+), 1 chloride ion (Cl-) b.Ex: Li 2 O = 2 lithium ions (Li+), 1 oxide ion (O 2- ) c.Ex: CaCl 2 = 1 calcium ion (Ca 2+ ), 2 chloride ions (Cl-) 13

14. D.Formula unit: the smallest whole-number ratio of the ions that participate in the bond. 1.Use “crossover rule” to determine number of each type of ion that participates in the bond: The number of (+) charges for cation becomes subscript for anion and the number of (-) charges for anion becomes subscript for cation. Ex: MgCl 2 : write out ions: Mg 2+ ; Cl - Then bring cation charge down as anion subscript and anion charge down as cation subscript Mg 2+, Cl - MgCl 2 If subscripts are equal, then reduce to 1 each. Ex: Mg 2+, O 2- would make Mg 2 O 2, reduces to MgO 14

15. 2.More examples: Aluminum Chloride: Al 3+ ; Cl - :will need 3 Cl - to balance out the Al 3+ So AlCl 3 Aluminum Oxide: Al 3+, O 2- Al 2 O 3 *Charges must balance: 2 molecules of Al 3+ will have 6 (+) charges; 3 molecules of O 2- will have 6 (-) charges, so balances! 15

16. E.Naming ionic compounds: list cation first, followed by anion (in “-ide” form): 1.Ex: CaCl 2 : calcium chloride 2.Ex: BeO: beryllium oxide 3.Your turn: NaS : ______________ _____________ F.If cation can have more than one charge (Fe 2+, Fe 3+ ), then use Roman numeral after the cation name to denote which cation: 1.Ex: FeO: iron (II) oxide 2.Ex: Fe 2 O 3 : iron (III) oxide 3.Your turn: Cu 2 O: ______________ _____________ 4.Your turn: CuO: ______________ ___________ 16

17. III.Properties of Ionic Compounds A.Most are crystalline solids at room temperature. 1.Arranged in repeating 3-D patterns 2.Creates strong attractive forces that make ionic compounds very stable, so… B.Ionic compounds have high melting points C.But are also “brittle” b/c atoms have no flexibility to move 1.Shatter when hit with hammer 17

18. D.Conduct electric current when dissolved in water, but not as solid crystal: 1.Water breaks apart the crystalline structure into its ionic particles 2.The ions have charge and can then conduct electricity a.Solutions of dissolved ions are called “electrolytes” 18

19. IV. Bonds A.Metals are made of closely packed cations surrounded by “sea of electrons” 1.Valence electrons are “mobile” and can drift between cations of the metal. 19

20. B.Properties of metals: 1.Metals conduct electricity b/c electrons can move freely: electrons move into one end of the metal, and an equal number of electrons leave the other end 2.Metals conduct heat well: closely packed atoms allows for easier transfer of heat. 3.Metals are “ductile”: can be drawn into thin wires a.Individual atoms are not bonded to each other, so they can “slide” past each other. 20