頁二第張五第日八十月二年丑発展复
WAH KIU YAT PO
報日橋筆
日期星
1973英文中學會考試題預習專欄、
BY IN AN & APT 3R
化學科 (廿三)
Chemistry (23))
Solution of last week. Part I
Part II
1. (a) The presence of oxygen in
nitric acid is demonstrated by the following experiment, in which a clay pipe is used as shown in the diagram.
Frinc HWD,
SE
The pipe stem is heated to red heat at one point with a burner. By dropping concent- rated nitric acid into the bowl of the pipe, a colourless gas bubbles up at the other end and is collected in the boiling-tube. When the gas is tested with a glowing splint, it bursts into flame. The gas is oxygen. Oxygen is, there- fore, present in nitric acid.
4HNO ► 220 + 4NO2 + 02
(b) By dropping a piece of zinc into the solution of copper nitrate, a reddish mass is observed. The zinc is tapped for some minutes until all disappears. By filtration a reddish mass of copper is left in filter-paper. Copper nitrate, therefore, contains copper.
Zn + Cu(NO2)2 •Zn(NO3)2 + Cu
(a) 25 cc. of potassium hydro- xide solution are transferred to a conical flask. To the solution, a few drops of litmus solution are added. Nitric acid is then run from a burette into the conical flask until the colour of the solution just changes. The amount of acid to bring about this change is the amount need- ed to neutralise the potassium hydroxide solution. The experi -ment is repeated without litmus and with a second lot of 25 cc. of potassium hydro- xide and same volume of nitric acid. The mixture will not now contain any excess nitric acid or any excess potassium hydroxide, but will be a neutral solution of potassium nitrate.
KOH + HNO3
KNO 3
H2O
Pure potassium nitrate crystals can be obtained from the solution by crystallisation. The solution is evaporated for a time to remove some water from it. At intervals, a drop of the solution is taken out
with a glass rod, and is then observed to find if the solut- ion will crystallize on coolin When nearly saturated, the hot solution is set aside to cool and crystallize. The crystals of potassium nitrate which se- parate out on cooling, are filtened off, washed two or three times with a small quantity of cold distilled water and dried between filter- papers.
(b) Lead sulphate can be pre- parated by adding dilute sulph- uric acid to lead nitrate solution.
Pb(NO3)2
2HNO 3
Dilute sulphuric acid is continuously added drop by drop to the solution of lead nitrate until no more white precipitate of lead sulphate comes out. The precipitate can be obtain- ed by filtration, washed with plenty of distilled water to remove any soluble impurities, and then dried in water.
(a) Add a little ferrous sul- phate solution to the solution of sodium nitrate in a test tube, shake to mix them well, tilt the test tube slightly, and carefully pour down along the edge of the test tube con- centrated sulphuric acid. The concentrated sulphuric acid, being much denser, sinks to the bottom, and a brown ring appears at the surface separat- ing the two liquids.
Mixture of Highly and Fase set...
·Cine the seq.
Brown Ring H2SO4
At the surface between the two liquids the conc. sulphuric acid acts on the sodium nitrate to produce nitric acid, NÁNO ̧ + H2SO ——— NaHSO HNO
4
This nitric acid is then reduced by the ferrous sulphate. to nitric oxide
2HNO3 + 3H2SO4 + 6F€2 (SO4)3
4H 0 2NO
The nitric oxide combines. with more ferrous sulphate to form the brown compound FeSO.NO seen in the brown
ring.
NO+ FeSO
FeSO NO
(b) When nitric acid is poured on dry sawdust, it takes fire, When nitric acid is poured on magnesium, the metal is dis- Solved; magnesium nitrate, ammonia and water are formed.
When nitric acid is poured. in solution of potassium iodide then iodine nitrogen peroxide, potassium nitrate and water formed.
When two drops of nitric acid are added to a solution of sulphur dioxide and a solut-
ion of barium chloride is add- ed, the solution becomes milky because the solution of sul phur dioxide is oxidized by nitric acid to form sulphuric acid which reacts with barium chloride forming the white, precipitate of barium sulphate,
12. Rate of reaction and energet ics.
Many chemical reactions go to completion very fast some are slow. Examples of the former type are neutralisation and pre- cipitation reactions, i.e. react iona between ions. There are also many reactions in which the reacting substances are not completely transformed into the products, because the latter react to reform the original reacting substances. Thus when steam is passed over heated iron hydrogen and magnetic iron oxide are formed. On the other hand, when hydrogen is passed over heated iron oxide, steam and iron produced. This is an example of a reversible chemical action.
3Fe + 4H2O
Fe3O4 + 4H
The direction in which the action proceeds depends on the conditions employed. Such as to remove the hydrogen as fast as it is formed, the action preceds fron left to right. If the steam is constantly removed no iron oxide. remains. By having steam and iron in a vessel a condition is reached in which all four substances exist together in equilibrium. The reaction is then a balanced one, The equilibrium, however, is a dynamic one and not a static one- that is, the molecules are still reacting together, but the veloci- ities of the forward and backward
reactions have become equal.
·月四年三七九一屡公年二十六國民華中 育教僑華
Factors affecting chemical change 1. physical atate of the reactants
The velocity of a chemical action is frequently affected by whether the substances are in solution or in solid form; if in solution, by the solvent employed, and, if solid, by tha state of aggregation of the particles. For examples: a) Solid silver nitrate is
mixed with sodium chloride. no action occurs, while in solution reaction is immedi ate.
b) Ethyl chloride when mixed
with an aqueous solution of silver nitrate gives a precipitate of silver nitra- te precipitation is rapid, Lead burns when sprinkled into oxygen in the form of minute particles.
Concentration of the reactants
The concentration of sub- stances affects both the velo- city of the reaction and the position of equilibrium when the reaction is reversible.
The law of mass action states The rate at which any chemical reaction is proceeding at any instant is proportional to the active masses of the reacting substances at that instant.
A + B [conc. of c). [cone of A]
While K is the equilibriu constant
3. Temperature
a) Effect on Reaction Velocity
The rate of a chemical reaction is always increased by employing a higher temperature,
b) Effect on position of
equilibrium
Le Chatelier's principle: When a system in equilibrium is subjected to a constraint that change takes place which tends to remove the constraint.
1) An endothermic reaction
is favoured by rise of temperature, e.g.
N2
2NO
02
302
CS2
C. +25
In all these cases, the LR reaction is favour
-ed by high temperature. ii) An exothermic reaction is
favoured by lowering of temperature. Examples of exothermic reactions are the following
250
3H
2503 2NH3
4. Pressure
a) Effect on Reaction Velocity
An increase of pressure on a gaseous system is : equivalent to an increase in the effective concentrat- ion of the reacting substan- ces.
Effect on position of Equilibrium
The proporations of result -ants and reactants in a balanced action may or may not be altered by a change in pressure. This depends on whether the reaction is accompainied by a change in volume or not. For Examples?
1) Forward action favoured
by increased pressure:
N2+ 3H2 2502 02
2NH3
2503
ii) Forward action favoured
by decreased pressure:
PC + 12 2N02
PC15
No204
iii) Equilibrium not affected
by pressure:
H2+
N2
02
2HI
2NO
5. Catalyst
A catalyst is an agent which alters the rate of a chemical reaction but remains unchanged,
chemically and in mass, at the end of the reaction. Examples of catalytic reactions
are:
1) Nickel is used as catalyst
in the hydrogenation of oil. ii) Iron is used as catalyst in
the synthesis of ammonia (Haber process).
iii) Platinum or Vanadium
pentoxide (V20) may be used to catalyse the combinatior of sulphur dioxide and oxygen to form sulphur. trioxide.
Notice the following features catalysis:
1) If the rate of reaction is decreased, the catalyst is called a negative catalyst if the rate of reaction is increased, called positive catalyst.
ii) The catalyst may be changed?
physically, or may undergo intermediate chemical change, during catalysis.
iii) It is generally understood);
that a catalyst is operative even if its proportion to the reagents is very minute. iv) Catalytic power is more or, less specific in character.} A substance which improves the efficiency of a catalyst is called a promoter. A substance which gradually suppresses the activity of the catalyst is said to poison the catalyst. Chemical Energy
In a chemical change, energy. is evolved or absorbed in many forms.
A Heat
Nearly all chemical reactions are accompanied by the evolution or absorption of heat. In an exothermic reaction heat is given out to the environment, in an. endothermic reaction heat is taken in from the environment. This may be indicated by adding to the right-hand side of the equation the number of calories evolved or absorbed when the reaction occurs between formula-weights in grammes of the substances as represented by the equation. This quantity is known as the heat of reaction. A positive sign(+) indicates that heat is evolved and a negative sign (-) that heat is absorbed.
In an exothermic reacti
i on the system loses energy. The change in energy in energy of the system is represented by H this being negative when heat isl given out. Thus,
·CO2 + 94,000 cal
or 0 + 0 CO2 AH = -94,000cal
In an endothermic reaction/ heat is taken in from the environ- ment, but this heat is gained by the system. Thus,
CS-25,400 cal
0 + 25
or C + 25 CS2 AH=+25,400 cal B. Electricity
Electricity can be produced
as the result of chemical react. ions. All cells produce their electricity in just this way. For example, the Daniell cell's reaction are as follows: anode: Zn(s)→→ Zn2+(aq.)
2+
H = +72 K cal.
2+
cathode: Cu(s) Cu
(aq) + 2e
H-123 K cal.
overall: Zn(s)
(aq.).
Zn2+(aq.) +
This means that tu(s):
H = -51 K cal.
is an
evolution of 51 K cal (which appears chiefly as electrical energy) for each gramme-aton of zinc dissolved and each gramme- atom of copper deposited.
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