Unit 4 Review

1. Unit 4 Review

2. Section 1: Forming Bonds 1) Why are noble gases unreactive? 2) Why do ions form? 3) What is the octet rule? How is it used? 4) Which is more likely to form a compound – Mg or Ne? Why? They have a full octet To achieve a stable octet (8 valence electrons) Atoms will gain/lose or share electrons to get 8 electrons. Used to predict ion formation. Mg, because it does not have 8 valence electrons.

3. Section 1: Forming Bonds 5) Why are the halogens and the alkali metals most likely to form ions? 6) How would the following elements achieve stability? Mg _________ O___________ N _________ Na __________ I __________ Xe _______ Because they only need one electron to be stable. Lose 2 electrons gain 2 electrons Gain 3 electrons Lose 1 electron Gain 1 electron Nothing – already stable

4. “+” 7) What is the charge on cations? ________ Which is more likely to form a cation – a metal or a nonmetal? Do cations form by gaining or losing electrons? Would they be smaller or larger than their neutral atom?

5. “-” 8)What is the charge on anions? __________ Which is more likely to form a anion – a metal or a nonmetal? Do anions form by gaining or losing electrons? Would they be smaller or larger than their neutral atom?

6. Attraction between a cation and anion 9) What is an ionic bond? 10) Describe how an ionic bond forms between sodium and chlorine. 11) Why don’t we talk about ionic compounds as individual molecules? 12) What is another name for ionic compounds? ___________________ Sodium loses one electron to chlorine; Na+ + Cl- = NaCl Because we never have individuals – they repeat 1000’s of times in a crystal lattice salts

7. Metal Nonmetal 13) An ionic bond always forms between a __________________ & a _________________. 14) What are the properties of salts? Know what causes each property. 1. 2. 3. 4. Hard and dense brittle High melting point Only conduct when dissolved.

8. Al (OH)3 Aluminum hydroxide ____________ Iron (III) oxide __________________ Magnesium sulfide _____________ Strontium phosphide ____________ Lithium sulfide _________________ Rubidium phosphate ____________ Fe2O3 MgS Sr3P2 Li2S Rb3PO4

9. Rb2Se BaI2 Rubidium selenide _____________ Barium iodide _________________ Copper (II) hydroxide____________ Manganese (IV) chloride _________ Strontium acetate ______________ Calcium sulfate _________________ Mercury (I) nitride _______________ Cesium nitride _________________ Cu(OH)2 MnCl4 Sr(C2H3O2)2 CaSO4 Hg3N Cs3N

10. calcium chloride silver chloride CaCl2 ____________________________ AgCl ____________________________ K2S _____________________________ CuO ____________________________ Ba3N2 ____________________________ FeN _____________________________ MgBr2 ___________________________ Fe3N2 ____________________________ potassium sulfide copper (II) oxide barium nitride iron (III) nitride magnesium bromide iron (II) nitride

11. ammonium chloride aluminum phosphate NH4Cl ____________________________ AlPO4 ____________________________ LiNO2 ___________________________ Cs2SO3 __________________________ FeCl3 ____________________________ Cu2O____________________________ Fe(NO2)2__________________________ lithium nitrite cesium sulfite iron (III) chloride copper (I) oxide iron (II) nitrite

12. Metals are not ionic but DO share many properties with ionic compounds. 15. Name one similarity and one difference between ionic and metallic bonding. 16. Describe the electron sea model of metallic bonding & draw a picture. Similarity: attraction between positive and negative ions Difference: metals are malleable and ionics are brittle Positive metal cations have lost electrons to the “sea”

13. 17. What physical properties does the electron sea model give metallic bonds? 18. What characteristic of metallic bonds enable them to conduct electricity? Malleable and ductile Conductors Lustrous Delocalized electrons are free to move which is how electricity is transferred