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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 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. Metals / cations Non- metals/ anions

5 Valence electrons

6 Periodic Table

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

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) + 2H 2 O (l) 2NaOH (aq) + H 2(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 (Na 2 O) instantly and this oxide forms a white coating and protects the underlying metal from any further reaction.  4Na (s)+ O 2 (g) 2Na 2 O (s)

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

12 Reaction with acids Sodium reacts with acid to produce the corresponding salt & hydrogen :  2Na (s)+ 2HCl(aq) 2NaCl (aq)+ H 2 (g)  H 2 SO 4 (aq)+ 2 Na (s) Na 2 SO 4 (aq)+ H 2 (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) + Cl 2(g) 2Na + Cl - (s)  Sodium + bromine = sodium bromide 2Na(s)+ Br 2 (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) + 2H 2 O(l) ---> Ca(OH) 2 (aq ) + H 2 (gas)

17 Reaction with Oxygen  Calcium burn vigorously with a brick-red flame.  Calcium reacts with oxygen to make calcium oxide, CaO.  2Ca (s) + O 2(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) CaCl 2(aq) + H 2(g) when calcium reacts with Sulphuric acid it forms calcium sulphate salts and hydrogen gas as products:  Ca(s)+ H 2 SO 4 (aq) CaSO 4 (s) + H 2

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 (CaBr 2 ). Ca (s) + Br 2 (g) Ca Br 2  Calcium & the halogen fluorine form calcium fluoride  Ca(s) + F 2(g) Ca F 2 (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) + 2H 2 O(l) -> Mg(OH) 2 (aq) + H 2 (g) The reaction with steam occurs readily, however, & the products are magnesium oxide & hydrogen: Mg (s) + H 2 O(g) -> MgO(s) + H 2 (g)

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

23 Reaction with acids Magnesium undergoes vigorous reactions with dilute acids, to produce the corresponding salt (magnesium cloride; magnesium sulphate) & hydrogen Mg (s) + 2HCl (aq) --> MgCl 2 (aq) + H 2(g) Mg (s) + H 2 SO 4 (aq) -> MgSO 4 (aq) + H 2(g)

24 Reaction with halogens  Magnesium reacts readily with chlorine to produce magnesium chloride Mg (s) + Cl 2 (g) -> MgCl 2 (s)  Magnesium reacts readily with bromine to produce magnesium bromide Mg (s) + Br 2 (g) -> MgBr 2 (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 H 2 O(l) 2Al(OH) 3 (s)+ H 2 (g)  Aluminium reacts readily with steam, but aluminium oxide is formed:  2 Al(s) + 3 H 2 O(g) Al 2 O 3 (s) + 3H 2 (g)

27 Reaction with Oxygen  Aluminium burns with a characteristic brilliant white flame & the oxide is formed:  2 Al(s) + 3 O 2 (g) 2Al 2 O 3 (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 (aluminium cloride; aluminium sulphate) salt:  2Al(s) + 6HCl(aq) 2AlCl 3 (aq)+ 3H 2 (g)  2Al(s) + 3H 2 SO 4 (aq) Al 2 (SO 4 ) 3 aq + 3H 2 (g)

29 Reaction with halogens  Aluminium reacts vigorously with all halogens to form the aluminium halide:  2 Al(s) + 3 Cl 2 (g) Al 2 Cl 3 (s)

30 Reaction with alkali  Aluminium dissolves in sodium hydroxide with the evolution of hydrogen gas, H 2, and the formation of aluminates of the type [Al(OH) 4 ] -.  2Al(s) + 2NaOH(aq) + 6H 2 O → 2Na + (aq) + 2[Al(OH) 4 ] - + 3H 2 (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 + 6H 2 O (l) …… …. 2NaAl(OH) 4 (aq)+ 3H 2 (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) + H 2 O (g) → ZnO (s) + H 2 (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) + O 2 (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, H 2.  Zn(s) + H 2 SO 4 (aq) → ZnSO 4 (aq)+H 2 (g)  Zn(s) + 2HCl(aq) → ZnCl 2 (aq)+H 2 (g)

36 Reaction with halogens  Zinc dibromide, zinc(II) dibromide, ZnBr 2, and zinc diiodide, zinc(II) diiodide, NiI 2, are formed in the reactions of zinc metal and bromine, Br 2, or iodine, I 2.  Zn(s) + Br 2 (g) → ZnBr 2 (s) [white]  Zn(s) + I 2 (g) → ZnI 2 (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) + 2H 2 O (l) K 2 Zn(OH) 4 (aq) + H 2 (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) + 4H 2 O(g) → Fe 3 O 4 (s) + 4H 2 (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, O 2, the result is formation of the iron oxides Fe 2 O 3 and Fe 3 O 4.  4Fe(s) + 3O 2 (g) → 2Fe 2 O 3 (s)  3Fe(s) + 2O 2 (g) → Fe 3 O 4 (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, H 2. In practice, the Fe(II) is present as the complex ion [Fe(OH 2 ) 6 ] 2+.  Fe(s) + H 2 SO 4 (aq) → Fe 2+ (aq) + SO 4 2- (aq) + H 2 (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 F 2, Cl 2, and Br 2, to form ferric, that is, Fe(III), halides.  Iron(III) fluoride; Iron(III) chloride; Iron(III) bromide.  2Fe(s) + 3F 2 (g) → 2FeF 3 (s) (white)  2Fe(s) + 3Cl 2 (g) → 2FeCl 3 (s) (dark brown)  2Fe(s) +3Br 2 (l) → 2FeBr 3 (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) + O 2 (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, H 2.  Cu(s) + H 2 SO 4 (aq) → Cu 2+ (aq) + SO 4 2- (aq) + H 2 (g)

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

48 Reaction with halogens  The reaction between copper metal and the halogens fluorine, F 2, chlorine, Cl 2, or bromine, Br 2, affords the corresponding dihalides copper(II) fluoride, CuF 2, copper(II) chloride, CuCl 2, or copper(II) bromide, CuBr 2 respectively.  Cu(s) + F 2 (g) → CuF 2 (s) [white]  Cu(s) + Cl 2 (g) → CuCl 2 (s) [yellow-brown]  Cu(s) + Br 2 (g) → CuBr 2 (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, HNO 3. In the latter case, nitrogen oxides are formed together with lead(II) nitrate, Pb(NO 3 ) 2.  Halogens  Alkali

51 Reaction with Oxygen  The surface of metallic lead is protected by a thin layer of lead oxide, PbO. Only upon heating lead to 600-800°C does lead react with oxygen in air to from lead (II) oxide, PbO (a yellow solid).  2Pb(s) + O 2 (g) → 2PbO(s)

52  Reactivity series  What makes a metal reactive ?  Displacement of metals The reactivity of metals

53 The Reactivity Series of Metals  Although most metals are usually electropositive in nature and lose electrons in a chemical reaction they do not react with the same vigor or speed.  Metals display different reactions towards different substances. The greater the ease with which an element loses its electrons and acquires a positive charge, the greater is its reactivity. Further, the greater the number of shells and lesser the number of valence electrons, the greater is the reactivity of the metal.  The activity series of metals, arranges all metals in order of their decreasing chemical activity. As we go down the activity series from potassium to gold the ease with which a metal loses electrons, and forms positive ions in solutions, decreases.

54 The Most Reactive and Least Reactive Metals  The most active metal, potassium, is at the top of the list and the least reactive metal, gold, is at the bottom of the list. Although hydrogen is a non-metal it is included in the activity series due to the fact that it behaves like a metal in most chemical reactions i.e., the hydrogen ion has a positive charge [H + ] like other metals.

55 Cont’d.  Two non-metals, carbon and hydrogen are important chemical reference points with regard to the method of metal extraction and reactivity towards acids. Metals above carbon cannot be extracted by carbon reduction and are usually extracted by electrolysis. Metals below hydrogen will not displace hydrogen from acids.

56 The Most Reactive and Least Reactive Metals Cont’d.  The higher the metal in the series, the more reactive it is i.e., its reaction is fast and more exothermic.  This also implies that the reverse reaction becomes more difficult i.e., the more reactive a metal, the more difficult it is to extract from its ore. The metal is also more susceptible to corrosion with oxygen and water.  The reactivity series can be established by observation of the reaction of metals with water, oxygen or acids.

57 Within the general reactivity or activity series there are some periodic table trends:  Down Group 1 (I) the "Alkali Metals" the activity increases Cs > Rb > K > Na > Li. Down Group 2 (II) the activity increases e.g., Ca > Mg.  On the same period, the Group 1 metal is more reactive than the group II metal, and the group II metal is more reactive than the Group III metal, and all three are more reactive than the "Transition Metals".  e.g., Na > Mg > Al (on Period 3) and K > Ca > Ga > Fe/Cu/Zn etc. (on Period 4)

58  The action of heat on the hydroxides; on the nitrates; on the carbonates  The action of water on the oxides; on the hydroxides; on the carbonates  Reduction of metal oxides  Reaction of metal oxides & hydroxides with dilute acids. The chemical properties of compounds of metals.

59 The action of heat on the hydroxides  Potassium & sodium are stable; there is no decomposition.  The metal hydroxide of Calcium, Magnesium, Aluminium, Zinc, Iron, Lead, Copper, Silver decomposes to form the corresponding oxide & water.  Mg(OH) 2 (s) -> MgO (s) + H 2 O(g)  Cu(OH) 2 (s) -> CuO (s) + 2H 2 O(g) heat

60 The action of heat on the nitrates  Potassium nitrate & sodium nitrate decompose to form the nitrite & oxygen gas:  2KNO 3 (s) -> 2KNO 2 (s) + O 2 (g)  2NaNO 3 (s) -> 2NaNO 2 (s) + O 2 (g) heat

61 The action of heat on the nitrates Cont’d  Nitrates of calcium, magnesium, aluminium, zinc, iron, lead, copper decompose to form the oxide, nitrogen dioxide (poisonous gas) & oxygen.  2Zn(NO 3 ) 2 (s) -> 2ZnO (s) +4NO 2 (g) +O 2 (g)  2Pb(NO 3 ) 2 (s) -> 2PbO (s)+4NO 2 (g) +O 2 (g)  Silver nitrate decomposes to give silver metal  2AgNO 3 (s) -> 2Ag (s) + 2NO 2 (g) + O 2 (g) heat

62 The action of heat on the carbonates  Potassium carbonate & sodium carbonate are stable. There is no decomposition.  Carbonates of Ca, Mg, Al, Zn, Fe, Pb, Cu, Ag decompose to form the metal oxide & carbon dioxide.  CaCO 3 (s) -> CaO (s) + CO 2 (g)  ZnCO 3 (s) -> ZnO (s) + CO 2 (g)  CuCO 3 (s) -> CuO (s) + CO 2 (g) heat

63 The action of water on the oxides  Oxides of potassium, sodium, calcium, magnesium are soluble in water forming a solution of corresponding hydroxide (an alkali).  K 2 O (s) + H 2 O(l) -> 2KOH (aq)  CaO (s) + H 2 O(l) -> Ca(OH) 2 (s)  Oxides of Al, Zn, Fe, Pb, Cu, Ag do not react with water.

64 The action of water on the hydroxides  Potassium, sodium :  Very soluble in water  Calcium, magnesium :  slightly soluble in water  Aluminium, zinc, iron, lead, copper, silver:  Insoluble in water

65 The action of water on the carbonates  Potassium, sodium :  Soluble in water  Calcium, magnesium, aluminum, zinc, iron, lead, copper, silver :  Insoluble in water

66 Reduction of metal oxides  Metal oxides can be reduced by hydrogen or carbon to the corresponding metal; however, this reaction occurs particularly with the oxides of the less reactive metals.  CuO (s) + H 2 (g) -> Cu (s) + H 2 O (l)

67 Reduction of metal oxides cont’d.  Carbon can be used to reduce the oxides of zinc & those metals that are lower in the reactivity series, while hydrogen wil reduce the oxides of iron, lead, copper & silver to the corresponding metal.  Fe 2 O 3 (s) + 3C (s) -> 2Fe (s) + 3CO(g)

68 Reaction of metal oxides & hydroxides with dilute acids  Metal oxides & hydroxides are classified as bases, so they will react with acids to produce a salt & water only. This is called a neutrali- sation reactions.  NaOH(aq) +HCl (aq) NaCl (aq) + H 2 O (l)  Ca(OH) 2 (s) + H 2 SO 4 (aq) -> CaSO 4 (s) + 2H 2 O(l)

69 Reaction of metal oxides & hydroxides with dilute aci ds cont’d.  CuO(s) + H 2 SO 4 (aq) -> CuSO 4 (aq) + H 2 O(l)  MgO(s) + 2HCl (aq) -> MgCl 2 (aq) + H 2 O(l)

70  Hydrogen  Carbon  Nitrogen  Oxygen  Silicon  Sulfur  Chlorine Chemical reactions of some non-metals

71 Chemical properties of H 2 Reaction of hydrogen with air  Hydrogen is a colourless gas, H 2, that is lighter than air. Mixtures of hydrogen gas and air do not react unless ignited with a flame or spark, in which case the result is a fire or explosion with a characteristic reddish flame whose only products are water, H 2 O.  2H 2 (g) + O 2 (g) → 2H 2 O(l)

72 Reaction of hydrogen with the halogens  Hydrogen gas, H 2, reacts with fluorine, F 2, in the dark to form hydrogen fluoride.  H 2 (g) + F 2 (g) → 2HF(g)  Hydrogen gas reacts with chlorine, Cl 2, to form hydrogen chloride.  H 2 (g) + Cl 2 (g) → 2HCl(g)

73 Reaction of hydrogen with nitrogen & sulphur  Hydrogen reacts with nitrogen & sulphur to give respectively ammonia & hydrogen sulfide  H 2 (g) + N 2 (g) → 2NH 3 (g)  H 2 (g) + S(s) → H 2 S (g)  Hydrogen acts as reducing agent since it removes oxygen from the oxides of metal:  H 2 (g) + CuO (s) → Cu (s) + H 2 O (l)

74 Chemical properties of C  Carbon, as graphite, burns to form gaseous carbon dioxide, CO 2. Diamond is a form of carbon and also burns in air when heated to 600-800°C - an expensive way to make carbon dioxide!  C(s) + O 2 (g) → CO 2 (g)  When the air or oxygen supply is restricted, incomplete combustion to carbon monoxide, CO, occurs.  2C(s) + O 2 (g) → 2CO(g)

75 Chemical properties of C cont’d.  Carbon is a good reducing agent since it removes the oxygen from metal oxides to form the metal & carbon dioxide:  2PbO (s) + C (s)……..2Pb (s) + CO 2 (g)  Carbon does not react with dilute acids & alkalis

76 Chemical properties of N 2  Nitrogen combine with oxygen & hydrogen to form nitrogen monoxide & ammonia, respectively  N 2 (g) + O 2 (g) → 2NO (g)  N 2 (g) + 3H 2 (g) → 2NH 3 (g)

77 Chemical properties of O 2  Oxygen is needed for combustion & therefore reacts with metals & non-metals to produce oxides.  O 2 (g) + Cu(s) → CuO 2 (g)  O 2 (g) + 2Ca(s) → 2CaO(s)

78 Chemical properties of Si  Silicon reacts vigorously with all the halogens to form silicon tetrahalides. So, it reacts with fluorine, F 2, chlorine, Cl 2, bromine, I 2, and iodine, I 2, to form respectively silicon(IV) fluoride, SiF 4, silicon(IV) chloride, SiCl 4, silicon(IV) bromide, SiBr 4, and silicon(IV) iodide, SiI 4. The reaction with fluorine takes palce at room temperature but the others requiring warming over 300°C.  Si(s) + 2F 2 (l) → SiF 4 (g)  Si(s) + 2Cl 2 (l) → SiCl 4 (g)  Si(s) + 2Br 2 (l) → SiBr 4 (l)  Si(s) + 2I 2 (l) → SiI 4 (s)

79 Reaction of silicon with air  The surface of lumps of silicon is protected by a very thin layer of silicon dioxide, SiO 2. This renders silicon more or less inert to further oxidation by air even up to about 900°C. After this, reaction with oxygen in the air gives silicon dioxide. At temperatures above about 1400°C, silicon reacts with nitrogen, N 2, in the air as well as oxygen, to form the silicon nitrides SiN and Si 3 N 4.  Si(s) + O 2 (g) → SiO 2 (s)  2Si(s) + N 2 (g) → 2SiN(s)  3Si(s) + 2N 2 (g) → Si 3 N 4 (s)

80 Reaction of silicon with bases  Silicon is attacked by bases such as aqueous sodium hydroxide to give silicates, highly complex species containing the anion [SiO 3 ] 2-.  Si(s) + 2NaOH(aq) +HO (l)→ Na 2 SiO 3 (aq) + 2H 2 (g)

81 Chemical properties of S  Sulphur burns in air to form the gaseous sulphur dioxide, SO 2.  S(s) + O 2 (g) → SO 2 (g)  Sulphur is also oxidised to sulphur dioxide in the presence of oxidising agents such as sulfuric acid.

82 Chemical properties of S Cont’d.  Sulphur combines with most metals to form the metal sulphide :  S(s) + Fe (s)………. FeS (s)  Hydrogen reacts with sulphur to produce hydrogen sulpide. Hydrogen sulphide has a characteristic odour loke that of rotten eggs.  Sulphur does not react with either dilute acids or alkalis.

83 The primary duty of the University to a student is to provide him with such instructors as will make him realize that the responsibility for progress is his own and no one else's. S.E. Whitnall, 1933


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