Topic 3.2: Other Trends (Physical and Chemical)

Transcription

1 Topic 3.2: Other Trends (Physical and Chemical) Essential Ideas: 3.2: Elements show trends in their physical and chemical properties across periods and down groups

2 Physical/ Chemical Properties Chemical properties are largely determined by the number of electrons in the outermost energy level Physical Properties Chemical Properties Conductors of Electricity and Heat Density Luster (shininess) Softness Appearance/ Color State of Matter Reactivity with Metals Water Oxygen Bond Type Acid/Base Properties

3 Physical Property: Melting Points Melting Points depend on BOTH the type of bonding and the structure (more discussion in Unit 4) Decrease down Group 1 because the metallic structures are held together by the attractive forces of delocalized outer shell electrons and the positive ions. This attraction decreases with distance.

4 Physical Property: Melting Points Increases down Group 7 because the molecular structures are held together by Van der Waal s forces which increase with the number of electrons in a molecule.

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6 Physical Property: Melting Points Increase across a period until they reach Group 14, then they decrease to reach a minimum at Group 18 This is due to a change in bond type/ structure Example: Period 3 Melting Points Atomic Number Element Melting Point (K) Bond Type 11 Na Mg 922 Metallic 13 Al Si 1683 Giant Covalent 15 P (P 4 ) S (S 8 ) Cl (Cl 2 ) 172 Covalent (simple molecules) 18 Ar 84

7 Practice Why do you think Sulfur has a high melting point than Phosphorus?

8 Practice Why do you think Sulfur has a high melting point than Phosphorus? Sulfur forms S 8 while P forms P 4 heavier molecule = more electrons = greater London Dispersion Forces More Dispersion Forces = Higher MP

9 Practice pg 110 # 32: Use the data below to identify the states of the 4 oxides listed under standard conditions. Explain the difference in melting points by referring to the bonding and the structure of each case. The oxides are added to separate samples of pure water. State whether the resulting liquid is acidic, neutral, or alkaline. Describe all chemical reactions by giving chemical equations. Oxides Melting Point (K) Boiling Point (K) MgO SiO 2 (quartz) P 4 O SO

10 pg 110 #32 Answer Oxides State of Matter Structure and Bonding MgO Solid Giant structure Ionic Bonding, Strong attraction between oppositely charged ions SiO 2 (quartz) Solid Giant Covalent Bonding- strong covalent bonds throughout structure P 4 O 10 Solid Molecular covalent bonding with weak IMFs between molecules (larger molecule than SO 2 therefore it has more London Dispersion Forces than SO 2 (stronger attraction) SO 2 Gas Molecular Covalent Bonding with weak IMF s between molecules

11 Group 18- Noble Gases Physical Properties Chemical Properties Colorless gases Monatomic: Exist as single atoms Least Reactive Elements due to stable octet

12 Group 1- Alkali Metals Physical Properties Good Conductors of Heat and Electricity due to mobile electrons (metallic bonding) Low Density Can cut with knife Silvery color Lustrous Chemical Properties VERY reactive (1 ve -1 ) Form ionic compounds with non-metals Reacts with water to form a base (alkaline solution)

13 Group 1- Alkali Metals- Reactions with Water Alkali Metals react with water to produce hydrogen gas and the metal hydroxide- you will need to be able to reproduce these equations from memory! Reaction becomes more vigorous as IE decreases. Example: Sodium 2 Na (s) + 2H 2 O (l) 2NaOH (aq) + H 2 (g) Observations (only what you see, not what you know) Sodium floats on the surface, melts into a sphere Bubbling due to gas production (not H 2 gas produced) Temperature of solution increases White smoke/vapor is produced

14 Group 17- Halogens Physical Properties Colored Show a gradual change from gases (F 2 & Cl 2 ) to liquids (Br 2 ) and solids (I 2 and At 2 ) Chemical Properties VERY reactive (7 ve -1 ) Reactivity decreases down the group Form Ionic compounds with metals Form Covalent compounds with nonmetals

15 Reactions with Group 1 and Group 17 Halogens react with Group 1 metals to form ionic halides Group 1 loses an electron while Group 17 gains an electron = stable octet for both Example: 2Na (s) + Cl 2 2NaCl(s) After the electron is transferred, the ions are held together by the strong electrostatic interaction. More Vigorous = elements that are farthest apart (Fr and F)

16 Displacement Reactions The more reactive halogen displaces the ions of the less reactive halogen from its compound Examples: 2KBr (aq) + Cl (aq) 2KCl (aq) + Br (aq) 2 2 Chlorine is more reactive than Bromine (higher Electronegativity Cl replaces Br

17 Displacement Reactions When Br 2 is displaced by chlorine, the solution turns from clear to orange When I 2 is displaced by bromine, the color darkens even more. It turns purple when mixed with a hydrocarbon Halogens also form insoluble salts with silver, which can help identify the halide

18 Practice

19 Practice Tube 1: Cl + KBr KCl + Br (turns orangish) 2 2 Tube 2: Cl + KI KCl + I (turns dark) 2 2 Tube 4: Br + KI KBr + I (becomes darker) 2 2 What would happen to tubes 2 and 4 if we added a hydrocarbon? Iodine turns violet

20 Bonding of the Period 3 Oxides When elements bond with oxygen, oxides are formed Ionic compounds are formed between a metal and a nonmetal Oxides of Na, Mg, and Al have giant ionic structures (high mp)

21 Bonding of the Period 3 Oxides Covalent compounds are formed between nonmetals so the oxides of P, S, and Cl are molecular covalent (low mp) The oxide of the metalloid silicon has a giant covalent structure (high mp)

22 Ionic Character Depends on the difference in electronegativity Greater difference = more ionic Oxygen = 3.5 EN Oxides Ionic Character decreases from L to R on the periodic table Oxide Ionic Character Increases down a group Electrical conductivity is a measure of Ionic character Can only be measured in molten state (electrons are free to move)

23 Acid- Base Character of Period 3 Oxides Amphoteric: Acts as an acid and a base You will be asked to write the equations from memory! Formula of Oxide Na 2 O MgO Al 2 O 3 SiO 2 P 4 O 10 or P 4 O 6 SO 3 or SO 2 Cl 2 O 7 or Cl 2 O Acid/Base Character Basic Amphoteric Acidic

24 Basic Oxides Dissolve in water to form alkaline solutions (presence of hydroxide ions) Na 2 O + H 2 O 2NaOH A basic oxide reacts with an acid to form salt and water Li 2 O + 2HCl 2LiCl + H 2 O

25 Acidic Oxides Non-metallic oxides react readily with water to produce acidic solutions Phosphorus (V) Oxide produces phosphoric acid P 4 O H 2 O 4H 3 PO 4 Phosphorus (III) oxide produces phosphorous acid P 4 O 6 + 6H 2 O 4H 3 PO 3

26 Acidic Oxides Acidic Oxide Reacts with Water to Form P 4 O 10 Phosphoric Acid H 3 PO 4 P 4 O 6 Phosphorous Acid H 3 PO 3 SO 3 Sulfuric Acid H 2 SO 4 SO 2 Sulfurous Acid H 2 SO 3 Cl 2 O 7 Perchloric Acid HClO 4 Cl 2 O Hypochlorous Acid HClO SiO 2 Insoluble in water

27 Amphoteric Oxides Aluminum oxide does not affect the ph when added to water because it is basically insoluble When reacted with base, it acts as an acid Al O + 3H SO Al (SO ) + 3H O When reacted with acid, it acts as a base Al O + 3H O + 2OH. - 2Al(OH)

28 Practice 1. Identify the oxide that forms an acidic solution when added to water. a. Na O 2 b. MgO c. SiO 2 d. SO 3