= Solution
\b[“Carbonic oxide” means <carbon monoxide>, $\mathrm{CO}$: a combustible, reducing, colourless gas distinct from <carbon dioxide>.] Several preparations remove part of the oxygen from a carbon–oxygen compound or prevent complete oxidation of carbon.
Passing <carbon dioxide> over sufficiently hot <carbon> gives
$$
\boxed{\mathrm{CO}_2+\mathrm C\rightleftharpoons2\mathrm{CO}.}
$$
High temperature favours substantial conversion. The reaction is important in furnaces, but gas emerging from such a process need not be pure <carbon monoxide>; residual <carbon dioxide> and other gases must be considered. Incomplete <combustion> of carbonaceous fuel is another source, with the ideal net balance $2\mathrm C+\mathrm O_2\to2\mathrm{CO}$. It is a source principle rather than a selective method guaranteeing a pure sample.
For a historical laboratory preparation, <formic acid> can be dehydrated by a suitable strongly dehydrating medium:
$$
\boxed{\mathrm{HCOOH}\longrightarrow\mathrm{CO}+\mathrm H_2\mathrm O.}
$$
<Oxalic acid> under an appropriate dehydrating treatment yields a mixture:
$$
\boxed{\mathrm H_2\mathrm C_2\mathrm O_4\longrightarrow\mathrm{CO}+\mathrm{CO}_2+\mathrm H_2\mathrm O.}
$$
An alkaline absorption stage removes much of the <carbon dioxide>, while <carbon monoxide> passes on. A generator, droplet trap, purifier and collecting vessel illustrate why preparation and purification are separate operations. The gas is only slightly soluble in <water>, so collection over water is possible in principle. Concentrated <sulfuric acid> acting on carbonaceous material can also give mixed carbon oxides and <sulfur dioxide>; that route requires additional purification and should not be presented as clean CO evolution without by-products.
The <relative density> of <carbon monoxide> against dry <air> is approximately $28.01/28.97\simeq0.967$. It has no reliable warning odour. It burns with a blue flame when supplied with suitable <oxygen>:
$$
\boxed{2\mathrm{CO}+\mathrm O_2\longrightarrow2\mathrm{CO}_2.}
$$
A CO sample itself does not ordinarily support combustion of a taper. Nor does it immediately give the usual limewater test of <carbon dioxide>; oxidation of the CO first gives a gas that does. This distinguishes fuel behaviour from the property of supporting combustion of another material.
Its reducing power is illustrated by
$$
\mathrm{CuO}+\mathrm{CO}\longrightarrow\mathrm{Cu}+\mathrm{CO}_2
$$
and by reduction of suitable iron oxides in metallurgy. The products reflect transfer of <oxygen> from the oxide to <carbon monoxide>. Its ability to bind metal centres is also chemically important. In biological exposure, binding to <haemoglobin> and other effects disrupt oxygen delivery and use; it is a serious poison, not merely an innocuous replacement of atmospheric oxygen. This modern explanation must not turn an historical gas-preparation answer into a proposed exposure experiment.<Carbon monoxide poisoning basics by the Centers for Disease Control and Prevention>[Sources and the colourless, odourless character of carbon monoxide]{ref}<Clinical Guidance for Carbon Monoxide Poisoning by CDC>[How carbon monoxide affects the body]{ref}
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