# 19.1 Occurrence, preparation, and properties of transition metals  (Page 9/27)

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$\text{CoO}\left(s\right)+{\text{2HNO}}_{3}\left(aq\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}\text{Co}{\left({\text{NO}}_{3}\right)}_{2}\left(aq\right)+{\text{H}}_{2}\text{O}\left(l\right)$
${\text{Sc}}_{2}{\text{O}}_{3}\left(s\right)+\text{6HCl}\left(aq\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{2ScCl}}_{3}\left(aq\right)+{\text{3H}}_{2}\text{O}\left(l\right)$

The oxides of metals with oxidation states of 4+ are amphoteric, and most are not soluble in either acids or bases. Vanadium(V) oxide, chromium(VI) oxide, and manganese(VII) oxide are acidic. They react with solutions of hydroxides to form salts of the oxyanions ${\text{VO}}_{4}{}^{3-},$ ${\text{CrO}}_{4}{}^{2-},$ and ${\text{MnO}}_{4}{}^{-}.$ For example, the complete ionic equation for the reaction of chromium(VI) oxide with a strong base is given by:

${\text{CrO}}_{3}\left(s\right)+{\text{2Na}}^{\text{+}}\left(aq\right)+{\text{2OH}}^{\text{−}}\left(aq\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{2Na}}^{\text{+}}\left(aq\right)+{\text{CrO}}_{4}{}^{2-}\left(aq\right)+{\text{H}}_{2}\text{O}\left(l\right)$

Chromium(VI) oxide and manganese(VII) oxide react with water to form the acids H 2 CrO 4 and HMnO 4 , respectively.

## Hydroxides

When a soluble hydroxide is added to an aqueous solution of a salt of a transition metal of the first transition series, a gelatinous precipitate forms. For example, adding a solution of sodium hydroxide to a solution of cobalt sulfate produces a gelatinous pink or blue precipitate of cobalt(II) hydroxide. The net ionic equation is:

${\text{Co}}^{2+}\left(aq\right)+{\text{2OH}}^{\text{−}}\left(aq\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}\text{Co}{\left(\text{OH}\right)}_{2}\left(s\right)$

In this and many other cases, these precipitates are hydroxides containing the transition metal ion, hydroxide ions, and water coordinated to the transition metal. In other cases, the precipitates are hydrated oxides composed of the metal ion, oxide ions, and water of hydration:

${\text{4Fe}}^{3+}\left(aq\right)+{\text{6OH}}^{\text{−}}\left(aq\right)+\text{n}\phantom{\rule{0.2em}{0ex}}{\text{H}}_{2}\text{O}\left(l\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}2{\text{Fe}}_{2}{\text{O}}_{3}\text{·}\left(\text{n}+3\right){\text{H}}_{2}\text{O}\left(s\right)$

These substances do not contain hydroxide ions. However, both the hydroxides and the hydrated oxides react with acids to form salts and water. When precipitating a metal from solution, it is necessary to avoid an excess of hydroxide ion, as this may lead to complex ion formation as discussed later in this chapter. The precipitated metal hydroxides can be separated for further processing or for waste disposal.

## Carbonates

Many of the elements of the first transition series form insoluble carbonates. It is possible to prepare these carbonates by the addition of a soluble carbonate salt to a solution of a transition metal salt. For example, nickel carbonate can be prepared from solutions of nickel nitrate and sodium carbonate according to the following net ionic equation:

${\text{Ni}}^{2+}\left(aq\right)+{\text{CO}}_{3}{}^{2-}\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{NiCO}}_{3}\left(s\right)$

The reactions of the transition metal carbonates are similar to those of the active metal carbonates. They react with acids to form metals salts, carbon dioxide, and water. Upon heating, they decompose, forming the transition metal oxides.

## Other salts

In many respects, the chemical behavior of the elements of the first transition series is very similar to that of the main group metals. In particular, the same types of reactions that are used to prepare salts of the main group metals can be used to prepare simple ionic salts of these elements.

A variety of salts can be prepared from metals that are more active than hydrogen by reaction with the corresponding acids: Scandium metal reacts with hydrobromic acid to form a solution of scandium bromide:

$\text{2Sc}\left(s\right)+\text{6HBr}\left(aq\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}2{\text{ScBr}}_{3}\left(aq\right)+{\text{3H}}_{2}\left(g\right)$

The common compounds that we have just discussed can also be used to prepare salts. The reactions involved include the reactions of oxides, hydroxides, or carbonates with acids. For example:

I don't use to see the messages
how can you determine the electronegativity of a compound or in molecules
when u move from left to right in a periodic table the negativity increases
reeza
Are you trying to say that the elctronegativity increases down the group and decreases across the period?
Ohanaka
yes and also increases across the period
reeza
for instance when you look at one group of elements in a periodic table electronegativity decreases when you go across the table electronegativity increases. hydrogen is more electronegative than sodium, potassium of that group. oxygen is more electronegative than carbon.
reeza
i hope we all know that organic compounds have carbon as their back bone
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Osakue
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hanna
what is the oxidation number of nitrogen, oxygen and sulphur
Osakue
5, -2 & -2
hanna
What is an atom?
is a smallest particle of a chemical element that can exist
Osakue
Osakue
it is a substance that cannot be broken down into simpler units by any chemical reaction
An atom is the smallest part of an element dat can take part in chemical reaction.
Idris
an atom is the smallest part of an element that can take part in a chemical reaction nd still retain it chemical properties
Precious
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Find the number of calcium atoms present in a sample weighing 2.0*10 raise to the power of -3g
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Is a branch of science that deals with matter in relation to energy
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acid, base and salt
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it could be prepared by extraction with water.
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If you extract it
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it is a test to confirm the presence of nitrate ion in solutions
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