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General Chemical Properties of Non Metals
The non-metals are very reactive and most of them react with other elements to form different compounds.
The following are important chemical properties of non-metals which are connected with their tendency towards electron gain in the course of formation of compounds:
1. The oxide of a non-metal is either acidic or neutral but never basic. The oxide of a non-metal is a covalent compound. Being acidic, it combines with water to form an acid, e.g.
2. A non-metal never replaces hydrogen in an acid to form a salt. This is because replacement of hydrogen in an acid is due to the fact that H+ accepts electrons supplied by a metallic atom.
A non-metal is an electron acceptor and so cannot release electrons to hydrogen ions in solution.
3. Non-metals form covalent chlorides, for example, the behaviour of phosphorus forming its chlorides is well known.
A covalent chloride like this is usually a volatile liquid, a non electrolyte, and rapidly hydrolysed by water.
These properties are characteristic of non metallic chlorides (except CCl4 which is not hydrolysed by water).
4. Non-metals combine with hydrogen to form many hydrides. A covalent compound is formed by equal sharing of electrons between or among the combining atoms. For example, methane ammonia, hydrogen chloride and hydrogen sulphide are the covalent compounds.
5. Non-metals are oxidizing agents As discussed early, non-metals accept electrons from other substances. Therefore, they are called oxidizing agents because, upon accepting the electrons, the substances donating these electrons are oxidized. So they act as the agents for oxidation of other substances.
The Oxidizing Properties of Non-metals

Non-metals react by gaining electrons to become negative ions. A univalent non-metal accepts one electron while a divalent one accepts two electrons. The ion formed carries the corresponding number of negative charges, but they rarely exceed two and never exceed three.

When a substance loses electron(s), it becomes oxidized, i.e. its oxidation number increases. This is called oxidation. Due to the fact that non-metals accept the electrons(s) donated by other substances, particularly metals, they are, therefore, termed as oxidizing agents. This is because by accepting the electrons, they help oxidize the electron donors.
Those substances or metals which donated the electrons are called reducing agents. This is because the electrons they donate reduce the oxidation number of non-metals. This process is called reduction. In this respect, non-metals act as oxidizing agents while metals act as reducing agents.

Strong and weak oxidants
As we have already seen, non-metals ionize by electron gain. In all cases, the extra electron(s) accepted lead to the formation of negative ions. The easiness of formation of negative ions depends on the ability of an element to accept the electrons. The ability of accepting electrons is called electronegativity of an element. Some elements are more electronegative than others.
The order of electronegativity of some non-metals is as follows: Fluorine < Chlorine > Bromine > Iodine > Nitrogen > Carbon
The degree of electronegativity indicates reactivity and hence oxidizing power of the element. Elements with higher electronegativity will displace those elements with lower electronegativity from their compounds.
Referring to the series above, fluorine will displace all the rest of the elements from their compounds as it is more electronegative than any other element in the series. Likewise, chlorine can displace bromine, iodine and nitrogen from their compounds. The displacement reaction occurs in the manner:
Where X is more electronegative than Y
The higher the electronegativity the stronger the oxidant. For example, bromine is a stronger oxidant than iodine, nitrogen and carbon.


The Displacement of Non-metals by another Non-metal from a Compound

Non-metals in the reactivity series
It is useful to placecarbonandhydrogeninto the reactivity series because these elements can be used to extract metals.
Here is the reactivity series including carbon and hydrogen:
Note that zinc and iron can bedisplacedfrom theiroxidesusing carbon but not using hydrogen. However, copper can be extracted using carbon or hydrogen.
Carbon Dioxide
Preparation of a Dry Sample of Carbon Dioxide Gas in the Laboratory
Prepare a dry sample of carbon dioxide gas in the laboratory
Carbon dioxide is one of the oxides of carbon. The gas is present in the air at a level of approximately 0.03% by volume. It is also found dissolved in water. The gas is one of the by-products of all decaying organic matter. Without carbon dioxide there is no life on earth. It is used by all plants in the process of photosynthesis and both plants and animals evolve carbon dioxide in respiration.
Laboratory preparation of carbon dioxide
Carbon dioxide is prepared in the laboratory by the action of dilute hydrochloride acid on marble (calcium carbonate).
When dilute hydrochloric acid is poured on marble chips, effervescence occurs. Dilute hydrochloric acid reacts with the marble chips to give calcium chloride, water and carbon dioxide.
CaCO3(s) + 2HCl(aq) → CaCl2(g) + H2O(l) + CO2(g)
The Properties of Carbon Dioxide
Analyse the properties of carbon dioxide
Physical properties
  1. Carbon dioxide is a colourless and odourless gas.
  2. It is denser than air.
  3. When the gas is cooled to –78°C, it turns straight into the solid (it sublimes). Sublimation is the change of a gas straight into a solid or change of a solid straight into a gas. Solid carbon dioxide is called dry ice. It sublimes when it is heated or exposed to air.
  4. It has a melting point of –199°C and boiling point of –91.5°C.
  5. Carbon dioxide does not support combustion. This is why it is used in fire extinguishers.

Chemical properties
Reaction of carbon dioxide with lime water (Test for carbon dioxide)
When a little carbon dioxide gas is bubbled into lime water (calcium hydroxide solution), the solution turns milky. This is due to the formation of a white precipitate of insoluble calcium carbonate.
Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)
This is a confirmatory test for the presence of carbon dioxide. The test serves to distinguish carbon dioxide from any other gas.
When excess carbon dioxide is bubbled into the lime water, the white perceptible dissolves to form a clear solution of soluble calcium hydrogen carbonate: CaCO3(s) + H2O(l) + CO2(g) → Ca(HCO3)2(aq)
Barium hydroxide can also be used to test for carbon dioxide as it forms a precipitate of barium carbonate: Ba(OH)2(aq) + CO2(g) → BaCO3(s) + H2O(l)

Reaction of carbon dioxide with magnesium

When a burning magnesium ribbon is lowered into a gas jar containing carbon dioxide gas, it continues to burn for a short time with a spluttering flame. A white ash of magnesium oxide and black specks of carbon are formed. The black specks of carbon can be seen on the sides of the gas jar.
2Mg(s) + CO2(g) → 2MgO(s) + C(s)
This clearly shows that carbon dioxide contains carbon and oxygen.
Reaction of carbon dioxide with water
Carbon dioxide reacts with water to form a weak carbonic acid. When carbon dioxide is bubbled into water, it dissolves to form a weakly acidic solution of carbonic acid:
H2O(l) + CO2(g) ⇔ H2CO3(aq)
The solution turns a blue litmus paper pink. This indicates that the solution is slightly acidic and hence too weak to turn the blue litmus paper to red (as strong acids do). The solution has no effect on red litmus paper.

The Uses of Carbon Dioxide

Explain the uses of carbon dioxide
Uses of Carbon Dioxide include:
  1. Fire extinguisher: Carbon dioxide is inert (i.e. it does not burn). It is dense than air and does not support combustion. Hence, it is a very useful fire-fighting chemical. When applied to fire, it forms a blanket over the burning material. Thus, it prevents air (oxygen) from reaching the burning material and therefore, extinguishing the flames.
  2. Manufacture of aerated (fizzy) drinks: Soda water and mineral water contain carbon dioxide dissolved under pressure. Because the gas is only slightly soluble, it is bubbled into these drinks under pressure to make more of it dissolve. When the bottles are opened, the gas escapes and it causes the “fizzy”.Dissolved carbon dioxide is responsible for the pleasant taste of soft drinks such as lemonade, Coca cola, Pepsi cola and other aerated drinks and mineral waters. Other beverages like beers are also bottled together with carbon dioxide.
  3. Refrigeration: Carbon dioxide is used for refrigeration purposes (i.e. in the deep-freezing of foods). The gas liquefies at ordinary pressure to form dry ice which sublimes at -78°C. Dry ice is a good refrigerant because it leaves no liquid after sublimation as is the case with ordinary ice.
  4. Manufacture of sodium carbonate by the Solvary Process:Carbon dioxide is used in the manufacture of anhydrous sodium carbonate in the Solvary Process. The sodium carbonate produced is used in the manufacture of glass.
  5. Manufacture of baking soda: Carbon dioxide is used in making baking soda (sodium bicarbonate). Baking soda is prepared by passing carbon dioxide into cold concentrated sodium hydroxide solution: CO2(g) + 2NaOH(aq) → Na2CO3(aq) + H2O(l).Further addition of carbon dioxide produces sodium bicarbonate which, at sufficiently high concentration, will precipitate out of the solution as a solid: Na2CO3(aq) + CO2(g) + H2O(l) → 2NaHCO3(s) Yeast and sodium bicarbonate (hydrogencarbonate) are important in the baking industry. Thus in baking of bread, yeast is added to flour, sugar and water (forming the dough). In the making of cakes, baking powder (a mixture of bicarbonate and an acid) is used instead of yeast.
  6. Rain making: When pieces of dry ice (solid carbon dioxide) are dropped into clouds, the temperature of the clouds is lowered to such an extent that rain precipitates out.
  7. Photosynthesis: Plants use carbon dioxide from the air to manufacture their own food through the process of photosynthesis.


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