Chemistry Chemistry As Level(Form Five) Chemistry Notes Form Five (As Level)


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Is the chemistry of all the elements and their compounds with the exception of most carbon compounds out of which only the oxides, cyanides and carbonates are considered as inorganic compound. Inorganic compound can also be defined as the study of the elements in the periodic table.


Periodic table is the table of all the known elements arranged in order of increasing atomic numbers. The arrangement reflects the electronics configuration of the elements.


The first comprehensive classification of elements was made independently by Mendeleev’s in Russia and Lothar Meyer in German in 1869.They tabulated all the known elements on the basis of relative atomic mass. The arrangement of elements in the early periodic table was according to the ordinary periodic law which states that ‘The properties of elements are periodic function of their relative atomic masses’.
When elements were arranged in order of increasing atomic masses, elements with similar properties recurred at regular intervals .The recurrence or repetition of elements with similar properties of regular intervals in the periodic table is known as PERIODICITY.
Mendeleev’s and Lothar Meyer placed elements in horizontal rows (periods) which caused elements with similar properties to appear in the same vertical column (group). Some element were not yet discovered and hence absent that periodic table e.g. Noble gases, gallium ,germanium etc.
On the basis of relative atomic mass three anomalies appeared in the early periodic table (Ordinary periodic table). The position of Potassium (39.1), Argon (39.9), Cobalt (58.94), Nickel (58.69), Tellurium (127.6) and Iodine (126.9) had to be reversed to bring them into correct placing on chemical grounds. The strict order of relative atomic mass would have separated the mentioned elements from closely related element. For example Potassium could have been separated from other alkali metals. These anomalies showed clearly that the relative atomic mass was not really the true basis of arranging or classifying elements in the periodic table.
The three anomalies stated above were due to ISOTOPE. For example both Argon and Potassium exhibit isotropy .The principal Isotopes of these elements are shown in the table below.
In the case of Argon, the heavier isotope is predominate(has higher abundance) giving an average atomic mass of 39.9,In Potassium the lighter isotope predominates giving an average atomic mass of 39.1, The explanation is similar in the case of cobalt(no isotopes) and nickel(five isotope) and tellurium(8 isotopes) and iodine(isotope)
Moseley discovered that atomic number was the proper criterion for arranging elements in the periodic table. The atomic number of an element determines the number of electronics in the atom and hence the arrangement and properties of the elements in the periodic table.
When atomic numbers are used instead of atomic masses, the anomalies observed between potassium ,Argon ,Cobalt ,Nickel Tellurium and Iodine disappears .The classification of the known elements in the periodic table is now based on the Modern Periodic law which states that “ The properties of elements are periodic function of their atomic numbers”
NB: One can’t use R.A.M or Atomic mass for arranging elements but atomic number is the best


The periodic table is very important in the study of inorganic chemistry. The relationship between the periodic table, atomic number and properties of elements enable us to obtain an overview of the many facts and features in Inorganic chemistry
The periodic table consist of boxes which are filled by elements .Each box contains the symbol, mass number, and atomic number of elements.
However, other data like electronegativity, boiling point, melting point, oxidation states values may be included or added. There are various lay out of the period table. The periodic table may be in the short or long form .Under this level the long form will be considered .The long form consists of all the elements of a periodic except the Lanthanide and Actinides
The elements with similar properties occur in vertical column called groups. The horizontal rows of elements in the periodic table are calling periods. The long form of the periodic table is into four major blocks according to the sub shell in which the respective elements fill their electronics. The four major blocks are S-blocks, P-blocks, d-blocks, and f- blocks.

The S-blocks consist of elements which fills their outermost electrons in the S-sub shell. These elements use their S-sub shell electrons for bonding. The S-block is constituted by the element of Group IA and Group IIA e.g. Li, Na, K, Mg etc


The P-blocks elements consist of elements which fills their outermost electrons in the P-sub shell. This block consist of elements of group IIIA up to VIIA example Carbon, Sulphur ,Phosphorus ,Nitrogen ,Oxygen, Chlorine, Argon ,Neon, Helium is not a P-block elements.
The S-and P-blocks elements together form main group elements .These elements use only electrons of their outermost shell for bonding .Hence there are seven main groups which constitute main group elements (example IA, IIA, III,IVA, VA, VIA and VIIA). The noble gas contains full outermost electrons S and P-sub shells. Since this configuration is very stable, the noble gases are uncreative.
However they form some compounds with strongly electronegative elements like Oxygen and Fluorine .Example of noble gases include Helium, Neon, Argon, Krypton etc
The d-blocks consist of elements with partially or full-filled d-sub shell. They fill their electrons in the d-sub shell of the penultimate shell. They use electrons from S and d-sub shells for bonding. Example Scandium, Manganese, Iron, Radium, Nickel, Cobalt, Copper, Zinc etc

The f-block consists of elements which are called Lanthanides and Actinides. They are also known inner transition elements. These elements have two partially filled sub-shells namely (n-1)d and (n-2)f .They all belong to III B because they are so similar that it is very difficult to separate one from another example e.g. La, U, Np ,Lw, Th etc
I. The atoms of all elements of the same period have the same number of shells which are partially or fully occupied by electrons.
II. The number of the main groups is equal to the number of electrons in the outermost shell.
III. The number of period is equal to the principal quantum number (n) of the outermost shell and to the total number of electrons shell of the elements in a given period.



I. I. The atomic size(Atomic Radius)
The atomic radius of an uncombined atom cannot be defined strictly because of the uncertain boundary of electron clouds .The distance between the nuclei of chemically or covalently combined atoms can be measured accurately by x-ray diffraction method
Therefore, the atomic radius is defined as half the distance between the nuclei of two similar/ identical atoms joined by a single covalent or metallic bond. There is significant regular increase in atomic radii among elements down the group .This trend is due to increase of number of electrons and number and number of shells down the group.
The newly added electrons or shell must be at greater distance from the nucleus than that of the proceeding element (i.e a noble gas). Also the added electron is shielded by the inner electrons. Therefore the added shells and electrons reduces the attraction force between the nucleus and the outer electrons leading to increase in atomic size down the group .The following is the trend down group IA.

II. Ionic radii

The ionic radius of the atom of an element is measured in the same way as the atomic radius. The ionic radii like atomic radii increases down the group. The reasons for this are exactly the same as those quoted for atomic radii. The radius of a cation is shorter than the radius of the parents atom (neutral atom) because electron or electrons have been removed and hence the force of attraction from the nucleus has increased e.g. Na=1.57, 0.97
In contrast, all anions are larger than the corresponding neutral atom because electrons have been added to complete the noble gas structure. The added electrons reduces the force of attraction from the nucleus and as the result the size of an anion increase example Cl=0.99 and 1.81. The repulsive force of inner electrons also contribute to increase in ionic size.
The table below shows the trend in ionic radii down group VIIA elements





NB: Ionic radius is half the distance between the nuclei of two ions in an ionic crystal

III. Ionization energy

This is the energy required to remove an electron from a gaseous atom or ion. Hence the first ionization energy of an element, M is the energy required to remove the first electron from it.
The second ionization energy is the energy required to remove the second electron from a gaseous ion.
The successive ionization energies are defined accordingly .The higher this energy the tighter the electron is bound to the atom or ion. The ionization energy decreases down the group because OR due to the increase in atomic size of the element down the group, the increase in radii downs a group correlate with the decreasing ionization energies.
The nuclear charge increases down the group but it is cancelled by shielding effect and screening effect of the electrons of the inner shells. The two effects increases down the group. Since the nuclear charge is affected by the increase of electrons, the atoms become larger down the group and therefore the outer electron(s) is easily removed as it is loosely bound. Thus for the alkali metals Li to Cs, the outer most electron is most readily removed for Cs which is the largest element in group IA. Fluorine having the smallest atomic size in group VIIA has the highest ionization energy.

IV. Electron affinity (E.A)

The electrons affinity is the energy change which occur when an electron is added to a gaseous atom or ion. The non-metallic element readily gain an electron to give a negative ion.
The energy associated with this process is termed as the electron affinity. The process may either absorb or evolve energy (example endothermic and exothermic respectively). Some elements evolve large quantities of energy i.e their electron affinities are large and negative. The more negative the electron affinity, the more the electron is attracted by the nucleus. Thus the first electrons affinity is the energy change which involve the addition of the first electron to a gaseous atom
Negative electron affinity (in kJ)
The most negative electron affinity is found in Ground VIIA elements (I.e halogens) halide ions are most easily formed as they possess high electrons affinity. There is little variation in electron affinity down the group. Also there is no simple trend of electron affinity as shown in group VIIA. However there is general decrease in electron affinity down the group
NB: From the table above Chlorine has higher electron affinity than Fluorine. This is probability due to the relative small atomic size of Fluorine compared to that of Chlorine. The higher electron cloud in small Fluorine atom exerts a great repulsive force to the incoming electron giving rise to a smaller value of electron affinity. Chlorine with larger atomic size has smaller repulsive force and as result an electron can easily to be added to it and hence high electron affinity. All the electron affinity valves quoted for univalent ions is negative .To add an electrons to a univalent anion may require a large amount of energy in order to overcome electrostatic repulsion between the second electron and the charge on the anion. For example the second electron affinity of oxygen is positive.
The first electron is easily accepted
The2nd electron is not easily accepted
More energy required to shift with to higher energy level to create space for incoming electron. Thus energy released smaller than the energy gained
IV. Electronegativity
This is the measure of the attraction which an atom exerts on the electron pairs of a covalent bond. It can also be defined as the power of an atom in a molecule to attract electron to itself. The two atoms assumed to be involved in forming a single covalent bond where shared electrons are not equally attracted by the two atoms

Electronegativity is directly proportional to effective nuclear charge and inversely proportional to atomic radius. Small atoms with relatively large effective nuclear charge tend to have large values of electronegativity. There is general decrease in electronegativity as one goes down a group. Therefore in the first element in each group (main group element) is much more electronegative than the rest. This account for much of the differences between its properties and that of other elements e .g Li is much more electronegative compared to other numbers of the group.
The following is the trend in electronegativity down group IA
The atomic radii decrease steadily across each period. There is step rise in atomic size from halogens to alkali metals.
Consider period 2 elements
From the table above, it is observed that there is decrease in atomic radius from Li to F followed by an increase to Neon. This is probably due to the fact noble gases are mono atomic and their inter-nuclear distance is the distance between non bonded atoms held together by weak Van Der Waal’s forces. The decrease in atomic size gets smaller and smaller as shown in the table above. The same trends are observed in other periods
Generally atomic size (atomic radius) continue to decrease in passing along the period, Hence the alkali metal atom has the largest atomic radius than all the element in each period.
The halogen atom has the smallest atom radius of the elements in each period except iodine which has slightly larger radius than some of the transition metals in the middle of the period.

On moving horizontally along the periods we find that the first ionization energy increased. This due to the increase effective nuclear charge and decrease in atomic size for a small atom the outer most electrons are firmly held by the nucleus and hence much energy is required to remove them.
Across period 2 and 3, a contractions in radius correlates with the rise in ionization energy values. Abnormally high ionization potential values in Beryllium, Magnesium, Nitrogen, and Phosphorus are explained on the basis of the extra stability associated with full S and half filled P sub shell respectively.
A break is observed between Be and B in the period 2 and between N and O in the same period. In the period 3 we observe break between Magnesium and Aluminum and age between Phosphorus and Sulphur
These breaks can be accounted as follow;
a) There is a large increase in the ionization energy as we pass from it to He. This is due to the stability of the double state of Helium
b) The break in the trend of increasing ionization energy between Be and B is due to extraordinary large ionization energy of Be
Be has two paired electrons in its 2S orbital. It needs energy first to split 2S electron pair and secondly to effect the electron removal. But for B there is only one electron in the 2P orbital which is easy to remove as it is more distant from the nucleus
A comparable explanation accounts for the break between N and O. Their electronic configurations are as follows;
For Nitrogen there is extra stability due to half-filled P-sub shell. In Oxygen the 2 sub level has a pair of electrons which shields the and electrons rendering them easier to remove, hence an abnormally lower ionization energy for Oxygen. We may take almost similar account for the elements in period 3 where there is break between Mg and Al and also between P and S.

1) Effective nuclear energy; The greater the effective nuclear charge the greater the ionization energy
2) Shielding Effect and Screening Effect; The greater the shielding and screening effect, the less the ionization energy
3) Radius ; The greater the distance between the nucleus and the outer electrons of an atom, the less the ionization energy
4) Sub level; An electron from a full or half –filled sub level requires additional energy to be removed

Ionic radii like atomic vary periodically with atomic number. The anions are much larger than the corresponding atoms while cations are usually much smaller. The explanation of the trends in atomic radii applies also to the very similar trends in ionic radii. Example
In period 4 the bivalent cation from Ti2+ to Z2+ show a d-blocks contraction similar to but larger than the corresponding contraction for atomic radius. The same is true of the Lanthanide contraction for the trivalent cations La3+ to Lu3+in period 6
Electron affinity along the period can be discussed in terms of non-metal and metals. It is generally observed that non-metal have higher values of electrons affinities than metals. Non-metals have higher values because they can readily gain an electron to give negatively charged ion. The process involving addition of electrons to neutral atom is exothermic because energy must be lost in order that the formed ion is stable.
Generally there is increase of electrons affinity across the group due to the increase of electronegativity which is a result of decrease in atomic radii of the respective elements. Also increase of effective nuclear charge increases electron affinity.

Electronegativity of the main group elements increase across each period. This is due to the increase in affective nuclear charge as well as decrease in atomic size. Therefore the halogens atom in every period has the highest electronegativity value than the rest members. Alkali are the least electronegative elements in each period.

The melting point of an element is the measure of the amount of energy (heat) which must be supplied to breakdown the regular arrangement of atoms or molecule in crystal.
It`s a temperature at which a substance changes from solid into liquid. There are different types of force which holds atoms and molecules together, example due to such variation in forces the melting point of the elements in a period do not change uniformity e. g Period 3 elements. Melting point increase sharply from Na+ to Mg. Sodium atom to contribute only one electron to the metallic crystals, but Mg contributes two electrons. This accounts for the lower Melting point of Sodium compared to Magnesium.
All has three electrons in the outer valence shell but contributes only two of them to the “electrons sea”. This is why it has melting point close to that of Magnesium. The third electron is held firmly to the extent that it is not contribute to the “sea of electrons”
Silicon is non-metal with some metallic properties like luster, electric conductivity and ability to form alloy with metals. Silicon has the highest Melting point in period 3 due to its “giant covalent structure”. The silicon giant structure is comparable to that of Diamond.
Phosphorus and Sulphur have relatively low Melting point because their molecules are held together by weak Van Der Waal’s forces. Sulphur has a higher Melting point than Phosphorus due to the differences in sizes of their molecules i.e
Chlorine is diatomic and is a gas at room temperature. Its melting point is very low due to a very low Van Der Waal forces holding the molecules.


Chemical properties along the period depend on the change of the elements from strongly metallic to non-metallic. Sodium and Magnesium are strongly metallic while Aluminum is weakly metallic element. Silicon, Phosphorus, Sulphur and Chlorine are non–metallic elements. Thus metallic properties decrease across the period from left side of the periodic table to the right


The hydrides of period 3 elements include NaH, MgH2, AlH3, SiH4, PH3, H2S and HCl. Sodium hydride is strongly ionic while Magnesium hydride is largely ionic but the bonds are covalent. Aluminium hydride is covalent.


Sodium .Magnesium, and Aluminum yields hydrogen gas and metal hydrides, they are basic in nature because they react with water to form base
Silicon hydride (silane) evolve hydrogen with water in alkaline medium (catalyses the reaction)
Phosphine (PH3) is a non–polar covalent compound and hence does not react with water. Phosphine is a non–polar covalent compound due to small difference in electronegativity values between Hydrogen and Phosphorus
The hydrides of Sulphur (H2S) and Chlorine (HCl) are polar covalent compounds. They hydrolyze in water to form acid.
Therefore the hydrides of the strongly metallic elements tend to form alkaline solution with water while those of the non-metals form acidic solutions.


The Chlorides of period 3 elements include NaCl, ,MgCl2, ,AlCl3, SiCl4, PCl3, PCl5, S2Cl2. The Sodium and Magnesium chlorides are ionic salts .The rest of the chlorides are covalent in nature. As metallic character decreases along the period ionic character of the chlorides decreases.


When the ionic chloride is added to water, there is an immediate attraction of polar water molecules for ions in the chlorides. The solid chloride either dissolves to form free ions example or react to form new substances.
The hydrolysis of chlorides varies across the period Sodium chloride is not hydrolyzed in water probability due to a large size of Na+ ion leading to low polarizing power for water molecules. The chlorides of the rest elements have covalent characters. As metallic character decrease along the period, ionic character of the chlorides decreases as well.
The extent of hydrolysis of chlorides varies across the period. Magnesium chloride is not hydrolyzed but its hydrated crystals undergo hydrolysis when heated to give hydrogen chloride and a basic Magnesium chloride salt.
Aluminum chloride is easily hydrolyzed by water to produce an acidic solution. Al3+ ion is very small in size and it is highly charged. Thus its polarizing power is high, due to its high polarizing power it forms hydrated ions ([Al(H20)6]3+). The Al3+ ion strongly polarizes the O-H bond of the water molecules in the [Al(H2O)6]3+ ion to the extent that the bonds break to release the hydrogen protons. The solvent water molecules abstract protons from polarized water molecules to form the acidic hydroxonium ion H3O+, The H3O+ I is formed as follows
The chloride of Silicon Phosphorus and Sulphur hydrolyses completely in water to form acidic solution or solid
Note: As metallic character decreases along the period, ionic nature of the chloride decrease while the extent of hydrolysis increases


The hydroxides in period 3 include NaOH, Mg(OH)2, Al( SiO(OH)2, P(OHPO(OH )3,SO2(OH)2 and ClO3(OH),Cl(OH). The hydroxides from silicon to chlorine are not true hydroxide but they are oxy–acids.
Sodium and Magnesium hydroxides are the true hydroxide. They have strong tendency of releasing the OH group. Aluminum hydroxide is Amphoteric.
The acidic character of non-metal hydroxides is due to the tendency of releasing or donating H+ ion when dissolved in water. Silicon, Phosphorus, Sulphur and Chlorine are electronegative enough to withdraw electron by inductive effect from the O-H bond thus, facilitating the release of hydrogen as a proton H+.
The acidity of the hydroxide of Si, P, S and Cl increases with the increase in the electro negativity of respective elements. Solubility decrease from NaOH to due to decrease in metallic character or ion character of the hydroxides across the period


Sodium and Magnesium are strongly metallic and their oxides are ionic. Aluminum oxide is ionic but not basic as these sodium and magnesium.
The oxides of sodium and magnesium are strong bases while aluminum oxides are amphoteric. The oxides of the remaining elements are all acidic.
Solubility of the oxides of period 3 elements decreases along the period as metallic properties decreases. The oxides of sodium and magnesium form hydroxide with water or steam because of the protective film of oxide. The oxides of phosphorus sulphur and chlorine reacts with water to form acidic solution. Silica (SiO2) does not react with water but it’s acidic


The first element in every group is the smallest and has the highest electronegativity compared to the rest group member as a result the first elements have properties which differ from the rest group member but similar to those at the next lower elements diagonally. The kind of relationship in which the elements which are diagonally located in the period’s table have similar properties is known as diagonal relationship.
Diagonal relationship may also be explained in terms of polarizing power of the diagonal elements .The polarizing power of an element is the ability of positive ion to polarize the negative ion .When positive and negative ions approach each other their shape are distorted. The extent by which the ion is able to undergo distortion is called polarizability. The effect is polarization is as follows;
NB; Polarization is the distortion or deformation of an electron cloud of an anion by a cation.
If polarization is quite small the ionic bond results and it is high the electrons in the anion are drawn towards the cation to the extent that a covalent bond is formed .The polarizability and polarizing power of an ion is affected by;
i. The size of the ion
ii. The charge on the ion
Thus polarizing power is high for a small ion which has high effective nuclear charge. On moving across a period from left to right, ionic radii decreases and effective nuclear charge increases. On descending a group ionic radii increases while the effective nuclear charge decreases. Hence the diagonal elements have similar polarizing power. Due to these elements have similar properties. The diagonal relationship can be represented as follows
Element Li Be B C N O F
Electronegativity 1.O 1.5 2.0 2.5 3.0 3.5 4.0
Element Na Mg Al Si P S Cl
Electronegativity 0.9 1.2 1.5 1.8 2.1 2.5 3.0
The relationship is most significant in the following pairs Li and Mg, Be and Al, Be and Si etc


Lithium resembles magnesium and differs from the other alkali metals as follows;



outline factors that enable elements to have diagonal similarities

A. similar electronegativity
B. similar atomic and ionic sizes.
C. have ions with similar polarization power.


The first element in every group of the period show some properties which are not shown to other elements in the respective group .The element is said to show anomalous behavior when its properties differ with those of the rest group member. The anomalous behavior of the first element in a group is due to the following factors,
A. The first element in a group has the smallest atomic and ionic size when compared with the rest group members
B. The first element in a group has the highest ionization energy
C. The first element in a group has the highest electronegativity
D. The first element in a group has the higher electronic affinity


I. Lithium forms covalent compounds while other alkali metals form ionic compound example LiCl is a covalence compound while NaCl is ionic
II. Lithium reacts with nitrogen gas on heating to form ionic nitride while other alkali metals do not react
III. Lithium reacts slowly with cold water while other alkali metals react vigorously
IV. Lithium forms hydrated chloride while the rest group member form anhydrous chloride example LiCl.2H2O and NaCl (Li can polarize water due to high polarizing power)
V. When burnt in air ,lithium gives the monoxide while other alkali metals form peroxide and super oxide
Lithium monoxide
Sodium peroxide
Potassium peroxide
VI. Lithium does not form acetylide with ethyne (acetylene ) while other alkali metals form acetylide with ethyne
VII. Lithium is only alkali metal whose salts may undergo hydrolysis
VIII. The hydroxides of lithium decomposes on heating to monoxide and water while hydroxide of other alkali metals sublime undecomposed
IX. Lithium nitrate decomposes on heating into Lithium monoxide, nitrogen dioxide and oxygen while the nitrates of other alkali metals decompose into nitrites and oxygen.
X. Lithium carbonate decomposes gently heating into monoxide and carbon dioxide while the carbonates of other alkali metals are stable .They decompose at higher temperature.
XI. Lithium hydroxide is less soluble in water and hence a much weak base than sodium hydroxide a and potassium hydroxide
XII. Lithium chloride is deliquescent while chloride at the rest group members aren’t deliquescent
XIII. Lithium sulphate do not form alums while the sulphate of the rest group form alums


1. Beryllium react with concentrated solution of alkali to form hydroxo-complexes and hydrogen gas while other alkaline earth metal do not react
2. The oxides and hydroxides of beryllium are amphoteric while those of other alkaline earth metals are basic
3. Beryllium chlorides hydrolyze in water while the chlorides of the rest group members do not.
BeCl2 + H2O → Be(OH)2 + 2HCl (Hydrolysis)
4. Beryllium chloride dimerize in vapour state while the chloride of the rest group member do not dimerize.

Beryllium chloride Dimer
5. Beryllium form fluoro-complexes while other alkali earth metals do not.
6. The chloride of beryllium readily dissolved in organic solvents while the chlorides of other alkali earth metals do not readily dissolved in organic solvents.
7. Beryllium do not react with water or steam while other group members can react with either cold or boiling water or steam.
8. Beryllium oxide do not react with water while the oxide of the rest group members react with water to form hydroxide.
9. Beryllium do not react with either dilute or concentrated nitric acid while the rest group member react with both dilute and concentrated
10. Beryllium carbide hydrolyses in water to form methane while the carbides of other group members give ethyne

CaC2(s) + 2H2O(cold) → Ca(OH)2 + H – C ≡ C – H ((C = C)-2 Oxidation state)
11. Beryllium chloride fumes in moist air while chloride of the rest group members do not.
(From air)
The fumes are due to hydrolysis of in water where is given out.


Like other first elements in every group fluorine exhibit some properties which differ with the rest group members.
The anomalous behavior of fluorine is due to ;
i) Its smaller atomic and ionic size
ii) Absence of d-orbital
iii) Its highest electronegativity value
iv) Its higher electron affinity
The Anomalous behaviors of fluorine are as follows
1. Fluorine is a monovalent while other halogens show covalencies of 3 and 5. Chlorine and Iodine also show a valence of 7. Fluorine is a monovalent because it has no d-orbital while other members have d-orbital.
2. The elements in the periodic table show their highest oxidation states when combined with fluorine, example and SF8. This is due to highest electronegativity value of fluorine.
3. Fluorine forms hydrogen fluoride molecules which are strongly hydrogen –bonded. The hydrides of the rest halogens do not form hydrogen bonds between their molecules
4. The solubility of fluorides often differ markedly from the solubilities of other halides of the same metal, example the alkaline earth metal halides are very soluble in water except the fluorides which are insoluble .On the other hand silver fluoride is the only soluble silver halide. The fluorides of alkaline earth metal have higher lattice energies than hydration energies.
5. Metals show their highest degree of ionic character when combined with fluorine example and are ionic but Al and are covalent compound.
6. Hydrogen fluoride form acidic salts containing the biflouride ion (HF2). The halide of rest halogens form normal salts only, example etc.
7. Fluorine is the only halogen more electronegativity than oxygen and it often behaves differently to other halogens in reaction with oxygen containing compounds example water (F displaces O from water unlike others). Fluorine liberates oxygen from water while other halogens do not
8. Fluorine evolve oxygen from hot concentrated alkalis while other halogens form chloride and chlorate (v)
(Chlorate ion)
With cold dilute alkalis fluorine given a fluoride and difluorine monoxide while chlorine form chloride and hypochlorite.
(Difluorine monoxide)
(Hypochlorite ion)
9. Fluorine combines directly with carbon while other halogens have no effect on it.

(i.e COMPOUNDS OF Na, Mg, Ca, Al, Fe, Zn, Cu AND Pb)


Definition; An oxide is binary compound made up of oxygen and other elements, example MgO, PbO, Al2O3, H2O2, NO2, SO2 etc
Therefore metal oxides are binary compounds made up of oxygen and metal, example PbO, FeO, , etc
NOTE: The binary oxygen – fluorine compounds are not called oxides of fluorine, but are called fluorides of oxygen since fluorine is more electronegative than oxygen (example OF2- oxygen difluorine)

There are two methods of preparing metal oxides


In this method a metal reacts with a reagent or oxygen or air to give metallic oxide
I) A metal oxide may be prepared by burning a metal in air or oxygen

II) A metal oxide may be prepared by passing steam on red heat metal
III) A metal oxide may be prepared by reacting a metal with an oxidizing agent like HNO3

In this method, the metal oxide is obtained by heating carbonates, hydroxides and nitrates etc
The metal oxides may be classified as follows:


These are oxides which react with acids to form salt and water only. They also combine with acidic oxides to form salts. Basic oxides may be ionic or covalent
Acidic oxide are formed by metals in their higher oxidation states,
These oxides are generally covalent in nature. They dissolve in water to form oxy-acids and hence are called acid anhydrides.


These are oxides with both basic and acidic properties. They react with both acids and bases. Amphoteric oxides include ZnO, Al2O3, BeO, PbO, SnO2, Cr2O3,
i) AS BASES: They react with acids to form salt and water only

ii) AS ACIDIC OXIDES: These react with bases to form salt and water

Sodium metastannate

Sodium ferrite


Peroxides are compounds containing the peroxide ion Na2O2 Peroxides of alkali metals and alkaline earth metals can be prepared.
a) By heating the metal in the presence of excess oxygen or air

b) By heating the monoxides of the metal alone or in the presence of oxygen/air.
c) By the action of oxygen or air on the metal dissolved in liquid ammonia. This method is for the preparation of K2O2, Rb2O2 and Cs2O2
d) By the action of H2O2 on metallic salt solution in the presence of an alkali.
  • Stability of peroxides increases with increasing of the electropositive character of the metal
  • Peroxides are more stable in dry state than when they are in solution form
  • Many peroxides are highly hydrated due to hydrogen bonding,
example Na2O2.8H2O, CaO2.8H2O, BaO2.2H2O etc
  • They dissolved in water to form alkaline solution and hydrogen peroxide
(Ice-cold water)
  • When treated with dilute mineral acids peroxides give H2O2

  • Peroxide give O2 on heating and hence act as oxidizing agents
NB: PbO2 is not peroxide since this oxides does not give water when treated with mineral acids.


These are oxides of metals containing the ion the known superoxides are these of potassium (KO2), Rubidium (RbO2) and Caesium (Cs2O2)
Superoxides of K, Rb, and Cs can be prepared by burning the metal in excess of oxygen or air
• They are yellow solids
  • The stability of these superoxides is in the order KO2 Rb CsO2
  • They are strong oxidizing agents
  • They are hydrolyzed by water forming H2O2 and oxygen.



These are oxides composed of two simple oxides. The two simple oxides may be of the same metal or different metal in different oxidation states, example Red–Lead (Pb3O4) is combination of 2PbO and. Due to this can be written ionically
Also magnetite (Fe3O4) is a combination of FeO and Fe2O3
The mixed oxides with different metals are as follows;
Magnesium ferrite (MgFe2O4) MgOFe3+2O3
ZnFe2O4 nOFe23+O3 (Zinc ferrite)

  • It is a brilliant scarlet (bright red) solid insoluble in water
  • Red lead behaves chemically as if it were a loose compound of lead monoxide (PbO) and lead dioxide ( PbO2). For example it reacts with dilute HNO3 on warming to give and water whereby PbO2 is left with no reaction
  • The reaction above is a redox reaction indicating that is a reducing agent
USES; Red-lead is used as pigment in oil plants


Fe3O4 occurs naturally as Magnetite. It may be prepared by heating iron with oxygen or steam
– The compound is black in color
– Very strongly ferromagnetic
– The compound is inactive chemically
– React with acids as a double oxide giving a mixture of ferrous and ferric salts in solution


These are compounds of metals which contains hydroxide ions (OH) as the only negatively charged ion, example NaOH, Mg(OH)2, Zn(OH, Fe(OH)3
There are two methods of preparation of metal hydroxides.
The hydroxides which can be prepared by this method are those composed of strongly electropositive metals. Example LiOH, NaOH and Ca(OH)2
In this method the metal hydroxide is prepared by
a) The action of water on the metal oxide;


NB: the metal hydroxides which are prepared by action of water on metal oxides are soluble in water.
b) Action of calcium hydroxides (milk of lime) on a solution of carbonate, example preparation of NaOH and KOH. These metal hydroxides are prepared by precipitation the unwanted ions and layering a solution of the alkali. For instance when potassium carbonates and (Ca(OH) solution are mixed and then allowed to settle, a solution of potassium hydroxides may be decanted
c) Precipitation of a metal hydroxide by adding ammonia solution or sodium hydroxide solution to a solution of salt of the metal

d) Electrolysis of a solution of the metal chloride, example preparation of NaOH. Alkali metal chlorides form conducting solutions and since these metals are highly in the electrochemical series, their ions remain in solutions during electrolysis and hydrogen evolved at the cathode. Preferential discharge of chlorides ions enable hydroxyl ions formed by ionization of the water to accumulates in the solution. As a result dilute solution of the metal hydroxide is produced

-The hydroxides of group IA metals are white crystalline solids
-They melt at moderate temperature without decomposition except (LiOH)
-They are deliquescent solids
-They are very soluble in water (form alkali solutions)
-The basic strength of the alkali increases down the group, example calcium is the strongest base
-When cold and dilute alkali’s reacts with chlorine to form metal chloride and hydrochlorite
-When hot and concentrated alkali’s react with chlorine to form metal chloride and chlorate (v)

In the two reactions chlorine undergoes disproportion
– Most NaOH and KOH absorbs CO2 from the air whereby a metal carbonates is formed
– The hydroxides of Group IA metal reacts with acids to form salts and water only, example undergo neutralization reaction
  1. Owning to their highly basic character alkali metal hydroxides are used to absorb acidic gases, example CO2
  2. Alkali metal hydroxides are used in neutralization reaction.
  1. Alkali metal hydroxides are used in precipitation reaction
  1. Caustic soda (NaOH) is used in the manufacture of silk, paper and soap
  2. Caustic potash (KOH) is used to manufacture soft soaps
– They are white crystalline solids
– Solubilities increases considerably down the group from beryllium hydroxides (Be) to barium hydroxides (Ba). Beryllium hydroxide is insoluble in water.
Solubility of calcium hydroxide decreases with rise in temperature, the others increase, magnesium slightly but strontium and barium hydroxide greatly. Increase in solubility down the group is due to the fact that lattice energy decreases faster than hydration energy (Be() is essentially covalent because of the high polarizing effect of the small
– Group IIA hydroxides are much less soluble
-The hydroxides of Na and K precipitates some metals from their soluble salts (example, Aqueous solutions of their salts) as hydroxides

– Both NaOH and KOH liberates ammonia gas when added to ammonium salts
Zinc, Aluminium, Lead and Tin react with hydroxides of sodium and potassium to form complexes, example aluminate, plumbate, zincate and stumnate.

Aluminate ion
Zincate ion
Plumbate ion
When CO2 is bubbled through aqueous solutions of the NaOH and KOH the carbonates are formed,With excess of the CO2 the hydrogen carbonates are formed.

Of group IA elements due to the decrease in metallic character of the elements (example Group IA elements are more electropositive than their corresponding Group IIA elements). Also the decrease in solubility may be due to decrease ionic character of the hydroxides from Group IA to Group IIA
NB; A Suspension of slaked lime(calcium hydroxide) in water is called Milk of lime

  • Beryllium hydroxide is amphoteric. It reacts with excess sodium hydroxide forming a solution of sodium beryllate
Sodium beryllate
The other hydroxide of group IIA metals do not react with alkalis but react with acids to form salt and water only
– When is bubbled through lime water (Ca(OH)2) white precipitate of CaCO3 are formed. This causes the lime water to turn milky. The milky colour disappears when excess CO2 in bubbled through it. The milky colour disappears because calcium carbonates is converted into calcium hydrogen carbonate which is soluble in water
Clean solution White precipitate (milky)
(Milky) (Excess) Clear solution
The temperature at which the hydroxides begin to decompose increases down the group from about 3000C for beryllium hydroxide and magnesium hydroxide to about 7000C for barium hydroxide.

All the hydroxides except Be(OH) 2 react with aqueous ammonium salts to give ammonia gas. The ammonia gas is easily identified because it turn alkaline to litmus paper.
Sulphur dioxide turn lime water milky due to calcium sulphite formed. When excess SO2 is added the milky colour disappears (example, a clear solution is formed). The milky colour disappears due to the formation of calcium bisulphite which is soluble in water.


1. Lime water is used to test for carbon dioxide.

2. A suspension of Magnesium hydroxide in water (milk of magnesium)is used as an ant-acid.

3. Ca(OH)2 is used in making builders mortar (mixture of slaked lime, sand and water).

4. A mixture of Ca(OH)2 is used in making bleaching powder.

5. Ca(OH)2 is used for neutralizing acids in the soil.

6. A mixture Ca(OH)2 and water (white wash) is used for coating walls and ceiling.

7. Ca(OH)2 is used in water softening.

8. Ca(OH2) is used in sugar refining filtered.

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