1 / 83

Ch 17. J.C. Rowe

Windsor University School of Medicine. Characteristics of metals & non-metals Accept the challenge, so you may feel the exhilaration of victory. George S. Patton. Ch 17. J.C. Rowe. Periodic Table/ Review.

harris
Download Presentation

Ch 17. J.C. Rowe

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Windsor University School of Medicine Characteristics of metals & non-metalsAccept the challenge, so you may feel the exhilaration of victory.George S. Patton Ch 17. J.C. Rowe

  2. Periodic Table/ Review • The periodic table is a table of the chemical elements in which the elements are arranged by order of atomic number in such a way that the periodic properties (chemical periodicity) of the elements are made clear. • The standard form of the table includes periods (usually horizontal in the periodic table) and groups (usually vertical). Elements in groups have some similar properties to each other. • The periodic table is a masterpiece of organised chemical information. The evolution of chemistry's periodic table into the current form is an astonishing achievement with major contributions from many famous chemists and other eminent scientists

  3. Metals & non-metals in the Periodic Table.

  4. Stable electronic configurations Metals / cations Non- metals/ anions • Metals tend to lose their outershell electrons in order to be stable. • Stable positively charged ions (cations) • The ability to lose electrons makes them good reducing agents. • Non-metals tend to easily gain electrons in order to be stable. • Stable negatively charged ions (anions) • The ability to gain electrons makes them good oxidising agents.

  5. Valence electrons

  6. Periodic Table

  7. Chemical properties of selected metals Sodium (Na) Potassium (K) Calcium (Ca) Magnesium (Mg) Aluminium (Al) Zinc (Zn) Iron (Fe) Copper (Cu) Lead (Pb)

  8. Chemical Properties of Sodium • Chemical properties are all those properties that are visible only when any reaction is taking place between sodium and any other chemical substance. • As per the periodic table, sodium is more reactive as compared to lithium and has less reactive properties than potassium.

  9. Reaction with water • Reaction of sodium with water results in the formation of sodium hydroxide and hydrogen gas. As heat is produced during this reaction, it is called exothermic reaction. This released heat often ignites the hydrogen gas and as a result fire may break out. If large pieces of sodium is put into water it can lead to loud explosions. • 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)

  10. Reaction with Oxygen • Sodium readily reacts with oxygen to form sodium oxide. • When a fresh piece of sodium comes in contact with air, it forms sodium oxide (Na2 O) instantly and this oxide forms a white coating and protects the underlying metal from any further reaction. • 4Na (s)+ O2 (g) 2Na2O (s)

  11. Reaction with Oxygen cont’d. • When sodium is burned in air, it reacts with atmospheric oxygen to form sodium peroxide (Na2O2). • 2Na (s)+ excess O2 (g) Na2O2 (s)

  12. Reaction with acids Sodium reacts with acid to produce the corresponding salt & hydrogen : • 2Na (s)+ 2HCl(aq) 2NaCl (aq)+ H2 (g) • H2SO4 (aq)+ 2 Na (s)   Na2SO4 (aq)+ H2 (g)

  13. Reaction with halogens • The reaction between sodium and chlorine is very vigorous and the product is sodium chloride. The following equation illustrates this reaction. • sodium + chlorine = sodium chloride • 2 Na(s) + Cl2(g) 2Na + Cl-(s) • Sodium + bromine = sodium bromide 2Na(s)+ Br2(g) 2NaBr(s)

  14. Chemical Properties of Potassium • The reactions of potassium & sodium are similar; however the reactions of potassium are more violent.

  15. Chemical properties of Calcium • Calcium is a moderately active element. It reacts readily with oxygen to form calcium oxide (CaO):  • Calcium reacts with the halogens— fluorine, chlorine, bromine, iodine. • Calcium also reacts readily with cold water, most acids, and most nonmetals, such as sulfur and phosphorus. 

  16. Reaction with water • When water combines with water, it forms slaked lime, or calcium hydroxide (Ca(OH) 2 ): Calcium + Water ---> calcium Hydroxide + Hydrogen • Ca(s) + 2H2O(l) ---> Ca(OH)2(aq ) + H2(gas)

  17. Reaction with Oxygen • Calcium burn vigorously with a brick-red flame. • Calcium reacts with oxygen to make calcium oxide, CaO. • 2Ca (s) + O2(g) 2CaO (s)

  18. Reaction with acids • when calcium reacts with hydrochloric acid it forms calcium chloride (an alkali) and hydrogen gas as products : • Ca(s) + 2HCl(aq) CaCl2(aq) + H2(g) when calcium reacts with Sulphuric acid it forms calcium sulphate salts and hydrogen gas as products: • Ca(s)+ H2SO4 (aq) CaSO4 (s) + H2

  19. Reaction with halogens • Calcium burns in the halogens to form the corresponding calcium halide. • calcium (Ca) and the halogen bromine (Br) form the ionic compound calcium bromide (CaBr2). Ca (s) + Br2 (g) Ca Br2 • Calcium & thehalogen fluorine form calcium fluoride • Ca(s) + F2(g) Ca F2 (S)

  20. Chemical Properties of Magnesium • Magnesium is a very reactive metal and does not exist in a free state in nature. • It reacts with a slow pace with cold water and at a very rapid pace with hot water. • The oxidation process of magnesium is very rapid and if kept in open, a layer of oxidized magnesium is formed on the surface of the metal. Magnesium also burns very rapidly, when it is at room temperature.

  21. Reaction with water • Magnesium reacts slowly with cold water to form a solution of magnesium hydroxide & hydrogen: Mg (s) + 2H2O(l) -> Mg(OH)2(aq) + H2(g) The reaction with steam occurs readily, however, & the products are magnesium oxide & hydrogen:Mg (s) + H2O(g) -> MgO(s) + H2(g)

  22. Reaction with Oxygen • The magnesium is burned and reacts with the oxygen in the air. Magnesium + Oxygen gas → Magnesium oxide 2Mg (s)+ O2 (g)  → 2MgO (s)

  23. Reaction with acids Magnesium undergoes vigorous reactions with dilute acids, to produce the corresponding salt (magnesium cloride; magnesium sulphate) & hydrogenMg (s) + 2HCl(aq) --> MgCl2 (aq) + H2(g) Mg (s) + H2SO4 (aq) ->  MgSO4 (aq) + H2(g)

  24. Reaction with halogens • Magnesium reacts readily with chlorine to produce magnesium chloride Mg (s) + Cl2(g) -> MgCl2(s) • Magnesium reacts readily with bromine to produce magnesium bromide Mg (s) + Br2(g) -> MgBr2(s)

  25. Chemical Properties of Aluminium • The surface of aluminum metal is covered with a thin layer of oxide that helps protect the metal from attack by air. • So, normally, aluminum metal does not react with air. If the oxide layer is damaged, the aluminum metal is exposed to attack.

  26. Reaction with water • The protective oxide layer prevents the reaction with water. However, if the oxide layer is removed then aluminium reacts wiyh cold water to form aluminium hydroxide: • 2 Al(s) + 6 H2O(l) 2Al(OH)3 (s)+ H2 (g) • Aluminium reacts readily with steam, but aluminium oxide is formed: • 2 Al(s) + 3 H2O(g) Al2 O3(s) + 3H2(g)

  27. Reaction with Oxygen • Aluminium burns with a characteristic brilliant white flame & the oxide is formed: • 2 Al(s) + 3 O2(g) 2Al2 O3(s)

  28. Reaction with acids • Aluminium reacts slowly @ first & then speeds up after the oxide layer has been removed & the metal is exposed. • Aluminium reacts with dilute acids to produce the corresponding (aluminiumcloride; aluminiumsulphate) salt: • 2Al(s) + 6HCl(aq) 2AlCl3(aq)+ 3H2(g) • 2Al(s) + 3H2SO4(aq) Al2(SO4)3 aq+ 3H2(g)

  29. Reaction with halogens • Aluminium reacts vigorously with all halogens to form the aluminium halide: • 2 Al(s) + 3 Cl2 (g) Al2 Cl3(s)

  30. Reaction with alkali • Aluminium dissolves in sodium hydroxide with the evolution of hydrogen gas, H2, and the formation of aluminates of the type [Al(OH)4]-. • 2Al(s) + 2NaOH(aq) + 6H2O → 2Na+(aq) + 2[Al(OH)4]- + 3H2(g)

  31. Reaction with alkali cont’d. • Aluminium reacts with sodium (or potassium) hydroxide to form a complex salt (an aluminate: [Al(OH)4]- ) & hydrogen: • 2 Al(s) + 2NaOH + 6H2 O (l) …… …. 2NaAl(OH)4 (aq)+ 3H2 (g)

  32. Chemical properties of Zinc • Reacts with water • Reacts with dilute acids • Reacts with oxygen • Reacts with halogens • Reacts with alkali

  33. Reaction with water • Zinc does not react with water; however, red-hot zinc will react with steam to produce zinc oxide & hydrogen: • 2Zn(s) + H2 O(g) → ZnO (s) + H2 (g)

  34. Reaction with Oxygen • Zinc metal tarnishes in moist air. Zinc metal burns in air to form the white zinc(II) oxide, a material that turns yellow on prolonged heating. • 2Zn(s) + O2(g) → 2ZnO(s) [white]

  35. Reaction with acids • Zinc metal dissolves slowly in dilute acids to form solutions containing the aquated Zn(II) ion to produce the corresponding salt and hydrogen gas, H2. • Zn(s) + H2SO4(aq) → ZnSO4(aq)+H2(g) • Zn(s) + 2HCl(aq) → ZnCl2(aq)+H2(g)

  36. Reaction with halogens • Zinc dibromide, zinc(II) dibromide, ZnBr2, and zinc diiodide, zinc(II) diiodide, NiI2, are formed in the reactions of zinc metal and bromine, Br2, or iodine, I2. • Zn(s) + Br2(g) → ZnBr2(s) [white] • Zn(s) + I2(g) → ZnI2(s) [white]

  37. Reaction with alkali • Zinc metal dissolves in aqueous alkalis such as potassium hydroxide, KOH, to form zincates such as [Zn(OH)4]2-. The complex salts formed is called potassium zincate. Zn(s) + 2KOH(aq) + 2H2 O(l) K2 Zn(OH)4 (aq) + H2 (g)

  38. Chemical properties of Iron • Reacts with steam water • Reacts with dilute acids • Reacts with oxygen • React with halogens • Does not react with alkali

  39. Reaction with water • In the absence of air, cold water doesn’t react with iron. In the presence of air , however, rusting occurs. • Hot iron react with steam : • 3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)

  40. Reaction with Oxygen • Iron metal reacts in moist air by oxidation to give a hydrated iron oxide. This does not protect the iron surface to further reaction since it flakes off, exposing more iron metal to oxidation. This process is called rusting and is familiar to any car owner. • On heating with oxygen, O2, the result is formation of the iron oxides Fe2O3 and Fe3O4. • 4Fe(s) + 3O2(g) → 2Fe2O3(s) • 3Fe(s) + 2O2(g) → Fe3O4(s)

  41. Reaction with acids • ron metal dissolves readily in dilute sulphuric acid in the absence of oxygen to form solutions containing the aquated Fe(II) ion together with hydrogen gas, H2. In practice, the Fe(II) is present as the complex ion [Fe(OH2)6]2+. • Fe(s) + H2SO4(aq) → Fe2+(aq) + SO42- (aq) + H2(g) • If oxygen is present, some of the Fe(II) oxidizes to Fe(III).

  42. Reaction with halogens • Iron reacts with excess of the halogens F2, Cl2, and Br2, to form ferric, that is, Fe(III), halides. • Iron(III) fluoride; Iron(III) chloride; Iron(III) bromide. • 2Fe(s) + 3F2(g) → 2FeF3(s) (white) • 2Fe(s) + 3Cl2(g) → 2FeCl3(s) (dark brown) • 2Fe(s) +3Br2(l) → 2FeBr3(s) (reddish brown)

  43. Chemical properties of Copper • Does not react with water • Does not react with alkali • Reacts with dilute acids • Reacts with oxygen • Reacts with halogens

  44. Reaction with water/alkali • Copper doesn’t react with cold water or with steam. • Copper doesn’t react with alkali

  45. Reaction with Oxygen • Copper metal is stable in air under normal conditions. Copper does not burn in oxyen; however, a black oxide layer (copper (II)oxide) is formed on surface of the metal: • 2Cu(s) + O2(g) → 2CuO(s)

  46. Reaction with acids • Copper metal dissolves in hot concentrated sulphuric acid to form solutions containing the aquated Cu(II) ion together with hydrogen gas, H2. • Cu(s) + H2SO4(aq) → Cu2+(aq) + SO42-(aq) + H2(g)

  47. Reaction with acids cont’d • Copper metal also dissolves in dilute or concentrated nitric acid, HNO3. • 3Cu(s) + 8HNO3(aq) → 3Cu(NO3)2 (aq) + 2NO (g) +4 H2O(l)

  48. Reaction with halogens • The reaction between copper metal and the halogens fluorine, F2, chlorine, Cl2, or bromine, Br2, affords the corresponding dihalides copper(II) fluoride, CuF2, copper(II) chloride, CuCl2, or copper(II) bromide, CuBr2 respectively. • Cu(s) + F2(g) → CuF2(s) [white] • Cu(s) + Cl2(g) → CuCl2(s) [yellow-brown] • Cu(s) + Br2(g) → CuBr2(s) [black]

  49. Chemical Properties of Lead • Doesn’t react with water • Doesn’t react with dilute acids • Doesn’t react with alkali • Doesn’t react with halogen • Reacts with oxygen

  50. No Reactions with: • Water • The surface of metallic lead is protected by a thin layer of lead oxide, PbO. It does not react with water under normal conditions. • Dilute acids • The surface of metallic lead is protected by a thin layer of lead oxide, PbO. This renders the lead essentially insoluble in sulphuric acid, and so, in the past, a useful container of this acid. Lead reacts slowly with hydrochloric acid and nitric acid, HNO3. In the latter case, nitrogen oxides are formed together with lead(II) nitrate, Pb(NO3)2. • Halogens • Alkali

More Related