GENERAL CHEMISTRY

To view the Notes, click the following links below:

1. THE ATOM

2. CHEMICAL BONDING

3. ENVIRONMENTAL CHEMISTRY

4. SOIL CHEMISTRY

5. CHEMICAL ANALYSIS

ANSWERS FOR SOME QUESTIONS

GLOSSARY

PERIODIC TABLE

TABLE OF ELEMENTS SHOWING ATOMIC NUMBERS, ATOMIC MASS, AND ATOMIC RADII

TABLE OF SOME FUNDAMENTAL PHYSICAL CONSTANTS

INORGANIC CHEMISTRY

To view the Notes, click the following links below:

1. PERIODIC CLASSIFICATION

2. SELECTED COMPOUNDS OF METALS

3. EXTRACTION OF METALS

4.TRANSITION ELEMENTS

GLOSSARY

ORGANIC CHEMISTRY

To view the Notes, click the following links below:

1. INTRODUCTION TO ORGANIC CHEMISTRY

2. ALIPHATIC HYDROCARBONS

3. AROMATIC HYDROCARBONS

4. HALOGEN DERIVATIVES OF HYDROCARBONS

5. ALCOHOLS AND PHENOLS

6. CARBONYL COMPOUNDS

7. CARBOXYLIC ACIDS AND THEIR DERIVATIVES

8. AMINES

9. POLYMERS

GLOSSARY 

HOW TO BE GOOD AT CHEMISTRY – PART 1 

1.
Dig Deep on practice problems.

Just
doing lots of practice problems will not necessarily make you a better problem
solver. You will never see an exam problem that looks exactly like a practice
problem, so doing every problem possible is not a good strategy. Instead, when
you work out a practice problem, we have given you, make sure that you can
explain why and when you would make each step in your solution.

2.
Conceptual understanding is the goal

The
point of working on problems is NOT to get the “right answer.” By working on
problems, you will become more familiar with central concepts and ideas in
chemistry. Having a conceptual understanding means that you can: Take a core
chemistry idea and apply it to a new situation or problem. Think about core
ideas and their implications. Predict and explain chemical systems. Use
critical thinking to solve problems

3.
Self-study

‘Self-help
is the best help.’ You might have heard of this proverb. The same goes for studies.
Before you attend any class at tuition or school, remember to study the lesson
beforehand by yourself. The benefits of this habit are many. You will get the
opportunity to understand the lesson at your own pace. Also, while studying by
yourself, you will come across unknown terms and concepts that you could have
ignored in class. Your prior reading of the chapter can also help you connect
to the topics from previous grades, thus expanding your knowledge base.

……………

NORGANIC CHEMISTRY

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

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 PERIODIC TABLE CLASSIFICATION OF THE ELEMENTS (MENDELEEV’S AND LOTHAR MEYER 1869)

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

STRUCTURE OF PERIODIC TABLE

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.

1. S-BLOCK ELEMENTS

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

 2. P-BLOCK ELEMENTS

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

3. d-BLOCK ELEMENTS
   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
   
   4. f-BLOCK ELEMENT

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

NOTE;

1. The atoms of all elements of the same period have the same number of shells which are partially or fully occupied by electrons.
2. 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.

PERIODIC TRENDS IN PHYSICAL PROPERTIES

1. DOWN THE GROUP

1.   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

The2^nd 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

1. 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

2. ACROSS THE PERIOD
         I.    ATOMIC RADII(ATOMIC SIZE)

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.                                       

II.      IONIZATION 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;

1. 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
2. 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.

FACTORS AFFECTING IONIZATION ENERGY

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

 
III.  IONIC RADII

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

vvvvvvv          
IV.     ELECTRONIC AFFINITY

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.

1.                        
   ELECTRONEGATIVITY

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.

1.                     
          MELTING POINT

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.
_.

PERIODIC TRENDS IN CHEMICAL PROPERTIES ACROSS PERIOD THREE (Na to Ar)

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

a)     HYDRIDES

The hydrides of period 3 elements include NaH, MgH_2, AlH_3, SiH₄, PH_3, H₂S and HCl. Sodium hydride is strongly ionic while Magnesium hydride is largely ionic but the bonds are covalent. Aluminium hydride is covalent_.

REACTION OF HYDRIDES WITH WATER

Sodium .Magnesium, and Aluminum yields hydrogen gas and metal hydrides, they are basic in nature because they react with water to form base

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.

b)    CHLORIDES

The Chlorides of period 3 elements include NaCl, ,MgCl_{2, ,}AlCl_3, SiCl_4, PCl_3, PCl₅,  S₂Cl_2. 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.

REACTION OF CHLORIDES WITH WATER

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 extent of hydrolysis increases

c)     THE HYDROXIDES

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

d)    OXIDES

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.   

REACTION OF OXIDES WITH WATER

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

DIAGONAL RELATIONSHIP BETWEEN THE ELEMENTS

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;

1. The size of the ion
2. 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 AND MAGNESIUM.

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

QUESTION

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.

ANOMALOUS BEHAVIOUR OF THE FIRST ELEMENT IN A GROUP OF THE PERIODIC TABLE

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,

1. The first  element in a group has the smallest  atomic  and  ionic size  when compared with the rest group  members
2. The first  element in a group has the highest ionization energy
3. The first element in a group has the highest electronegativity
4. The first element in a group has the higher electronic  affinity

ANOMALOUS BEHAVIOUR OF LITHIUM 

1. Lithium  forms covalent  compounds  while  other alkali  metals  form ionic  compound  example  LiCl is  a covalence  compound while NaCl  is ionic
2. 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

1. 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)
2. When burnt in air ,lithium gives the monoxide while other alkali  metals

1. 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

ANOMALOUS BEHAVIOUR OF BERYLIUM

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

(From air)

The fumes   are due   to hydrolysis of  in water where    is given out.

ANOMALOUS BEHAVIOUR OF FLUORINE

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 ;

1. i) Its  smaller  atomic and ionic size
2. ii) Absence of d-orbital

iii)  Its highest  electronegativity  value

1. 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 combines directly with carbon while other halogens have no effect on it.

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

METAL OXIDES

Definition; An oxide is binary compound made up of oxygen and other elements, example  MgO, PbO, Al₂O_3, H₂O_2, NO_2, 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 OF_2- oxygen difluorine)

GENERAL METHODS IN PREPARATION OF METAL OXIDES.
There are two methods of preparing metal oxides

1. a)  DIRECT METHOD
2. b)  INDIRECT METHOD
3. DIRECT METHOD OF PREPARATION  OF METAL OXIDES

In this method a metal reacts with a reagent or oxygen or air to give metallic oxide

1. I)  A metal  oxide  may be prepared  by burning  a metal in air or oxygen
2. INDIRECT  METHOD OF PREPARATION  OF METAL OXIDES

In this method, the metal oxide is obtained by heating carbonates, hydroxides and nitrates etc

TYPES OF METALLIC OXIDE

The metal oxides may be classified as follows:

1)  BASIC OXIDE

2)  ACIDIC OXIDES

3)  AMPHOTERIC OXIDES

4)  PEROXIDES

5)  SUPEROXIDE

6)  MIXED OXIDES

1. BASIC OXIDES

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

Sodium ferrite

4. PEROXIDES

Peroxides are compounds containing the peroxide ion  Na₂O₂  Peroxides of alkali metals and alkaline earth metals can be prepared.

1. a)  By heating the metal in the presence of excess oxygen or air

1. b)  By heating the monoxides of the metal alone or in the presence of oxygen/air.

1. c)  By the action of oxygen or air on the metal dissolved in liquid ammonia. This method is for the preparation of K₂O_2, Rb₂O₂ and  Cs₂O₂
2. d)  By the action of H₂O₂ on metallic  salt solution in the presence  of an alkali.

         PROPERTIES OF PEROXIDES

• 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.

5.  SUPEROXIDES

These are oxides of metals containing the ion  the known superoxides are these of potassium (KO₂), Rubidium (RbO₂) and Caesium (Cs₂O₂)

Preparation
Superoxides of K, Rb, and Cs can be prepared by burning the metal in excess of oxygen or air

          PROPERTIES OF SUPEROXIDES.

• 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.

Ionically;

•  Superoxides are paramagnetic in nature due to the presence of one unpaired electron in  ion.

6.  MIXED OXIDES

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)

PROPERTIES

• 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

FERROUS –FERRIO OXIDE (Fe₃O₄) (FeO.Fe₂O₃)

Fe₃O₄ occurs naturally as Magnetite. It may be prepared by heating iron with oxygen or steam

 PROPERTIES

– 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

  HYDROXIDES OF THE METALS

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

PREPARATION OF METAL HYDROXIDES

There are two methods of preparation of metal hydroxides.

1. A)  DIRECT METHOD 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

Example

1. B) INDIRECT METHOD OF PREPARATION OF METAL HYDROXIDES

In this method the metal hydroxide is prepared by

1. a) The action  of water on the metal oxide;

   example

NB: the metal hydroxides which are prepared by action of water on metal oxides are soluble in water.

1. 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

1. c)  Precipitation of a metal hydroxide  by adding ammonia solution or sodium  hydroxide solution  to a solution of salt of the metal

1. 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

PROPERTIES OF METAL HYDROXIDES OF THE SELECTED METALS

1. ALKALI  METAL HYDROXIDES (MOH)

PHYSICAL PROPERTIES

-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)

CHEMICALS PROPERTIES

-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

USES OF HYDROXIDES OF Na AND K

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.

Example

1. Alkali metal hydroxides are used in precipitation reaction

Example

1. Caustic soda (NaOH) is used in the manufacture of silk, paper and soap
2. Caustic potash (KOH) is used to manufacture soft soaps

2. ALKALINE EARTH METAL HYDROXIDES (M(OH)₂₎

           PHYSICAL PROPERTIES

– 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

REACTION WITH AMPHOTERIC METALS

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

REACTION  WITH CARBONDIOXIDE

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

CHEMICAL PROPERTIES

1.  ACTION  WITH ACIDS  AND ALKALIS

•   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

2. ACTIONS WITH CARBON DIOXIDE

•  Moist hydroxides absorb CO₂ from air forming carbonates

–  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

3. ACTION OF HEAT

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.

Example

4.  ACTION WITH AMMONIUM SALTS

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.

Example

5. ACTION WITH SULPHUDIOXIDE  

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.

USES

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.

………….

Exercise (H/W)

1. Preparation of alkanes from petroleum, coal, natural gas.
2. Read on method of preparation known as cracking. Preparation of alkanes from petroleum and natural gas.

–       Petroleum is formed from the remains of tiny marine organisms that died and sank to the bottom of the sea millions of years ago.

–      Petroleum is a mixture of many organic compounds, since the organic compounds are lighter than both the rock and the water they move upwards through the adjacent rock. Sometimes the organic Compound are trapped in porous rocks that are called reservoirs beneath impermeable rocks.

–     Example of reservoir is limestone.

–     Reservoirs from which petroleum can be extracted by drilling are referred to as oil fields.

–    The petroleum obtained is referred to as a crude oil.

Fractional distillation of crude oil:

–    Major components of crude oil are:-

(i)      Residuals (coke, asphalt, tar).

(ii)     Lubricating oils.

(iii)  Fuel oils.

(iv)  Diesels.

(v)    Kerosene.

(vi)  Naptha.

(vii) Petrol.

(viii) Petroleum gas.

The components of the crude oils are known as fractions and different fractions are separated by heating them in a process known as Fractional distillation and it is done in a distillation tower called a still.

The oil is first evaporated by heating . The vapour rises up and the tower acts as a giant heat exchanger removing heat from the gases as they rise up. Temperature falls to 20 by the time vapour reaches the top. The vapour condenses as they rise up.

The heavier ones i.e. those with higher boiling points condense first. Gaseous fractions pass out at the top.

Cracking

Some of the fractions obtained from the fractional distillation of the crude oil are converted into new products.

Cracking is the conversion of large molecules of organic compounds into compounds with smaller molecules.

There are two methods;-

(i)   Thermal cracking.

(ii)  Catalytic cracking.

In thermal cracking the large molecule organic compound is heated to a high temperature until its molecule break apart.

In catalytic cracking, a catalyst speeds up the cracking process.

(e)     Preparation of Alkanes from alcohols:

By reduction of alcohols:

When alcohols are hated with concentrated hydroiodic acid and red phosphorus at 423k under high pressure, alcohols can be reduced to alkanes.

Physical properties of alkanes:

(a)     Boiling point and melting point:

Alkanes have low melting point and boiling point.

Reason: Since alkanes are non-polar molecules with weak Van – der – Waal forces between them then low temperature is required to break the bond hence low melting and boiling point.

Increase in molecular mass leading to the increase in melting & Boiling points

NB:

If you compare straight chain and branched chain of same molecular mass; straight chain has higher M.P &B.P. Branched chained isomers have lower boiling points and melting point than straight chain isomer.

Reason: Branched chains are more compact hence have less surface area. This is why they have low M.P and B.P. Straight chains have higher surface area.

(b)   Solubility:

They are soluble in non polar organic solvents but insoluble in polar compounds e.g. Water.

Chemical properties of alkanes:

In general, alkanes are non reactive (inert) compared to other classes of organic compounds.

Reason:

(i)    They don’t have a functional group.

(ii)   There bonds are quite strong i.e. c – c (strong bond). Large energy is needed to break the bond hence less reactive.

(iii)      These two bounds are almost non polar and therefore neither electrophilic nor nucleophilic substitution reaction can take place. Can’t react with electron – loving  species or a proton loving species (Nucleophilic).
–          Electrophilic reacts with a negatively charged species.

–          Nucleophilic reacts with a positively charged species.

Alkanes can undergo the following reaction;-

(i)     Substitution reaction.

(ii)    Oxidation reaction.

(iii)   Thermal decomposition (cracking).

All these reactions take place at high temperature or under the presence of light energy.

1. SUBSTITUTION REACTIONS
2. a) Halogenations

–     This is addition of halogens.

–     The reaction between alkane and halogen is known as free radical substitution reaction.

–     Free radical substitution reaction is the reaction in which a free radical substitutes atom/atoms in a molecule.

1. What is a free radical?

A free radical is an atom or group of atoms which consist of unpaired electrons.

Example:

v  Function of UV – light is to give out a free radical.

Mechanism of reaction

Homolytic sharing of electron i.e. equal sharing of electrons go to each chlorine atom

NOTE:

Free radicals are very reactive. It wants to become stable.