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In a previous unit we discussed electrovalent and covalent bonding. While electrovalent bonding is between metals and non metals, covalent bonding is between non-metals. In the formation of electrovalent and covalent bonds valence electrons play very important roles and each valence shell of the bonded atoms attain inert gas stable configuration. For metal- metal bond, the valence electrons are so few that electron sharing to attain electron octet is not possible.

Electrovalent bonds cannot be formed as metals tend to lose electrons and not accept them. A negatively charged metal ion is not possible. How then do we explain bonding in metallic solids?

How can we explain the fact that ntendlic solids are good conductors of heat and electricity? What major differences are there in the structures-of ionic and metallic solids? The above questions will be answered in this unit. We shall also explain the origin of intermolecular forces that hold covalent molecules together in the bulk sample and account for their special properties. The above explanation of metallic bonding implies that the lattice forms a single large crystal.

This accounts for the high strength of metals. There is no direction to metallic bond and so the metallic lattice can be distorted easily by hammering and drawing. Metals are malleable and ductile. The free moving electrons conduct heat and electricity by their movement. The strength of the metallic bond depends on the attraction of the electron cloud to the positive cores in the metal lattice. The metallic bond strength increases with the number of valence electrons each metal contribute into the electron 'cloud'.

Take the example of Mg 2, 8, 2 2p6 Na 2, 8, 1 1s2 2s2 2p6 3s' Sodium is a softer metal than magnesium because for sodium only one valence electron per atom but for magnesium two electrons are donated per atom to the electron cloud.

Following the above argument compare the strength of the metallic bonding in magnesium with that in aluminium For metals in the same group of the periodic table, metallic strength decreases down the group. The increase in atomic size down the group is not accompanied by any increase in electron cloud strength.

This listed properties of metals are explained by the metallic bonding just explained. Table 5. In addition to these bonds there are other weaker attractive forces that exist between atoms and molecules. The existence of these weak attractive forces explains a number of physical properties of some compounds.

Because these forces are usually between molecules they are called intermolecular forces. For example Van der Waal's forces, dipole-dipole attractions and hydrogen bonding. A non polar molecule is one in which the electron pair for bonding is equally shared by the atoms involved in the bond formation.

Examples of non polar molecules are N2 , C12 , H2 , 02 etc i. Non polar bond may also exist between unlike atoms if they have the same electronegativity. For exam! The movement of electrons around an atom can lead to a momentary shift of more electrons to one side of the molecule than the other. During this shift an imbalance in charge exists with one side of the molecule slightly positive and the other slightly negative.

The positive end will attract the negative end of another molecule close to it. This attraction constitute a bond. This attractive force may be strong but because it is for a short time its effect is generally very small. The magnitude of this force increases with increasing number of electrons This force is present between all molecules atoms and ions. Its effect can be very large when there are many electrons in the molecules or atoms.

Take the case of the halogens Group VII elements fluorine, chlorine are gases, bromine is a liquid while iodine is a solid. Remember all of them exist as diatomic molecules and are only bonded together by van der waal forces, Van der Waal's forces are attractions between molecules which happen because of creation of temporary dipoles in all molecules.

The very large number of electrons in bromine and Iodine allows for substantial cohesive force between bromine and iodine molecules making bromine liquid and iodine solid at room temperature. Van der Waal's forces is sometimes called induced dipole- induced dipole attraction. The shared electron pair will be more under the control of the more electronegative atom. Take the example of HCI. Chlorine is more electronegative than hydrogen.

The shared pair of electron is controlled more by Chlorine. The chlorine end of the molecule will be slightly negative and the hydrogen end slightly positive e. This is dipole-dipole attraction. Though dipole-dipole interactions are not as substantial as full ion-ion interactions, they are stronger than Van der Waal's forces.

The table 5. Dipole interactions are only about one percent as strong as covalent and ionic bonds. In combination with these small electronegative elements, hydrogen carries a substantial positive charge. The attraction of this positive end with the negative end of another molecule will constitute a strong bond.

This bond is the hydrogen bond. Hydrogen bond is about 5 to 10 times stronger than ordinary dipole-dipole interaction. It is not as strong as ordinary covalent bonds between atoms in a compound. Hydrogen bonding is responsible for water being a liquid at room temperature rather than a gas. Hydrogen bonding explains the high boiling point of water compared to hydrogen sulphide see table 5. Hydrogen bonding explains why hydrofluoric acid is a weaker acid than hydrochloric acid.

No wonder the number of compounds are limitless. In this unit the types of bonding discussed are interatomic and intermolecular bonding. The metallic bonding is one of the major types of interatomic bonding and it explains very well the observed properties of metallic solids.

Weak bonding exists between molecules, atoms and ions as a result of instantaneous shift in electron distribution around atoms in compounds. This weak bonding can be substantial leading to solid structure of covalent compounds at room temperature.

Covalent bonding between unlike atoms will always lead to unequal share of bond electrons. Attraction between polar ends of molecules also account for the cohesive force between polar molecules, when the polar bond is between hydrogen and small electronegative elements. The cohesive energy of the dipole-dipole interaction can be very substantial. This may lead to abnormal behaviour of such compounds. It explains why water is a liquid instead of a gas at room temperature.

Senior Secondary Chemistry Textbook 2 Lagos. New School Chemistry Onitsha. Africana FEP Publishers. Recall that atoms are built of particles of three kinds: protons, neutrons and electrons. The nucleus of each atom is made of protons and neutrons. The number of protons the atomic number determines the electric charge of the nucleus, and the total number of protons and neutrons the mass number determines its mass.

In a neutral atom the number of electrons surrounding the nucleus is equal to the atomic number. The chemical and physical properties of an element are governed by the number and arrangement of the electrons Several attempts have been made since to group elements together based on recurring properties such as atomic weight. The most important step in the development of the periodic table was published in by Dmitri Mendelyeev, who made a thorough study of the relation between the atomic weights of the elements and their physical and chemical properties.

The word periodic means recur at regular interval. The initial arrangement has now been largely replaced following new knowledge about electronic structure of atoms.

The present periodic table is based on the recurrence of characteristic properties when elements are arranged in order of increasing atomic number.

In other words, the properties of the elements are the periodic function of their atomic number. When elements are systematically arranged in order of increasing atomic number, certain characteristics recur at regular intervals. The periodic table shows the arrangement of elements in seven horizontal rows and eight vertical columns as shown in table 6. The horizontal rows of the periodic table consist of a very short period containing hydrogen and helium, atomic number 1 and 2 , two short periods of 8 elements each, two long periods of 18 elements each, a very long period of 32 elements, and an incomplete period.

The elements in the period have the same number of shells and the number of valence electrons increases progressively by one across the period from left to right. For all members of the period the additional electron is added to the second shell hence the name period 2.

In general, every period starts with an element containing one electron in its outermost shell e. Li, Na, K and ends with an element whose outermost shell is completely filled e. He, Ne, Ar - the noble or inert elements.

The properties of elements change in a systematic way through a period. For example the first members of each period are all light metals that are reactive chemically, and this metallic character decrease across the periods which ends with unreactive inert gases.

The elements that appear in a vertical column belong to the same group or family. They have the same number of outer electrons or valence electrons and have closely related physical and chemical properties. The central elements of the long periods, called the representative elements have properties differing from those of the elements of the short periods. They are unstable and short-lived. Period 1 elements have one electron shell K ; period 2 elements have two electron shells K,L ; period 3 elements have three electron shells K,L,M ; etc.

The number of valence electrons in the atoms of the elements in the same period increase progressively by one from left to right. Across a given period, there is a progressive change in chemical properties. For example, metallic properties decrease across the period while non-metallic characteristics increases. The first three members of any period Groups 1 to 3 , except period 1 are metals while those of Group 4 to 7 and 0 are non-metallic in behaviour.

Using period 3 as an illustration, sodium, magnesium and aluminium are metallic and form mainly ionic compounds and basic oxides. To the right of the period, phosphorus, sulphur and chlorine are non-metallic and form mainly covalent compounds and acidic oxides. Hydrogen is placed in group IA for convenience only because of the single electron but does not have similar characteristic with other members of the group.

They react by losing this valence electron to form ionic or electrovalent bonds. The alkali metals are excellent conductors of electricity because the valence electrons are mobile. Because of their reactivity especially with water, the metals must be kept in an inert atmosphere or under oil. Sodium metal catches fire when in contact with water, so avoid dropping it in the sink in the laboratory The metals are useful chemical reagents in the laboratory, and they find industrial use in the manufacture of organic chemicals, dyestuffs and tetraethyl lead the anti-knock agent in gasoline.

Sodium is used in sodium - vapour lamps, and because of its high heat conductivity, in the stems of valves of airplane engines, to conduct heat away from the valve head. They have two electrons in their outermost shell and react essentially by forming divalent ionic bonds. Members of the group are trivalent since each of its atoms has three valence electrons and forms electrovalent compounds.

The oxide and hydroxide of aluminium are amphoteric - they have both acidic and basic properties. Their atoms each has four valence electrons and tend to form covalent compounds. Carbon is a non-metal, silicon and germanium are metalloids while tin and lead are metals showing a gradation from non-metallic to metallic character on going down the group. The compounds of carbon and hydrogen called hydrocarbons form a large class of organic compounds used as fuels e. Their atoms each has five valence electrons and show two common valence of 3 and 5.

Both of them are non-metals. They are electron acceptors in their reactions and form several oxides e. They are electron acceptors and are oxidising agents e. They are commonly called halogens. They are all non-metals and highly reactive.

The halogens show great similarity in their properties e. Group 0: Helium He , Neon Ne , Argon Ar , are the familiar members of this group which are commonly referred to as rare gases or noble gases. They have no bonding electrons because the outermost shell is completely filled hence the group name zero.

Members of the group exhibit similar properties which are different from those of the halogens that come before them and alkali metals that come after them. This is a confirmation that the end of a period has been reached. All the transition elements have the following characteristics. You should have learned that when elements are arranged in order of increasing atomic number, certain properties recur at regular intervals. Furthermore, you should have learned that the periodic table of elements serve to justify the trend of behaviour exhibited by elements.

It has served to introduce you to the periodic Table. The units that follow shall use the atomic orbital model to further justify the classification and explain the gradation of properties of elements based on the periodic table.

You have learned in unit 2 about the contributions of Rutherford and Bohr to atomic structure in order to obtain a model of the atom. Their contributions went a long way to explain some of the observation about the atom. The Rutherford's model of an atom as consisting of a central positively charged nucleus and the negatively charged electrons some distance away from the nucleus, is still acceptable.

However, classical electromagnetic theory denies the possibility of any stable electron orbits around the nucleus. In Bohr's model of the atom, the electron was restricted to being found in a definite regions i. In the Wave Mechanics Model, however, there is a slight chance that the electron may be located at distances other than in the restricted orbits. Despite this, we still accept Bohr's scheme for quantisation of energy in the atom and that the lowest energy level of the atom is the most stable state.

Although Bohr's contribution was remarkable, particularly his quantisation of energy, theory to explain the spectral lines for hydrogen atom; it has the following limitations: a The Bohr model failed to account for the frequencies of the spectral lines for complex atoms other than hydrogen.

The present day picture of the atom is based on wave mechanical or quantum mechanical treatment. The treatment reflects on the wave-nature of the electron and the quantisation of energy in the atom.

Although these treatments are fundamentally mathematical in nature, it describes the electron as point charge and that the density of the cloud at a specified point gives only the probability of finding electrons at that point.

We shall look at how this new thinking will help our understanding of the atom and the observed relation between electronic arrangement in atoms and the chemical behaviour of elements. The quantum theory attempts to understand how electrons are arranged in the atom based on wave and quantum mechanics treatment. The electron is visualised as a point charge. The density of this point charge varies in different locations around the nucleus and gives a measure of the probability of finding the electron at a specified point.

The region or space, around the nucleus, in which an electron in a given energy level is most likely or probable to be found is defined as an orbital. So rather than describing a fixed Bohr orbit in which electrons are located, the modem theory gives a probability description of atomic orbitals. The results of the quantum mechanical treatment of the atom is summarised below.

This designation is retained in the quantum model but to represent distinct energy levels and not shells or orbits. In otherwords, the quantum model recognises different quantised energy levels around the nucleus.

Each principal quantum number n corresponds to a particular energy level and has integral values of 1, 2, 3, 4, etc. Electron with the largest 'n' value has the most energy and occupies the highest energy level; and therefore the most easily removable or ionisable electron.

The maximum possible number of electrons in an energy level is given by 2n2. The subsidiary quantum number, 1, has integral values ranging from 0, 1, 2, Table 7. Number of sub-levels Names of the sub-levels ,-. Rather, the location of electron is defined in terms of probabilities which is described by the orbital. A region in space where there is a high probability of finding an electron in an atom is called an orbital.

The density cloud of the electrons defines the shape of the orbital. The electrons that move about to produce a spherical symmetrical cloud around the nucleus is an s- electron residing in an s-orbital. The p-electrons move about three axes, x, y and z that are at right angles to one another, producing a dumb-bell cloud around the nucleus along each axes. They are called the p-orbitals and are distinguished from each other by N, Py and Pz in line with the direction of the electron cloud. The geometrical representation of the d and f orbitals are more complex and beyond the scope of this programme.

However, before we can apply the quantum numbers to express the electronic configuration of atoms, there are two important rules that you should be familiar with. The principle simply means two electrons in an atom cannot behave in an identical manner. The way in which electrons are arranged in an atom is determined by the order of the sub-levels on a scale of increasing energy level. This is so because electrons are found in the lowest possible energy level, the ground state which is the most stable state of an atom.

A simple representation of the orbitals on an energy scale is given in Fig. To keep a check on the spin of the electron, arrows of opposite spins are used to distinguish two electrons in an orbital. One of the advantages of the electronic configuration of elements using quantum numbers is that it showed the basis for the periodic classification of element. In other words, the key to the periodicity of elements lies in the electronic configurations of their atoms.

The orbital arrangement of electrons clearly showed the great usefulness of the Period Table as it explains the groups and characteristic properties of elements. The correlations between electronic configuration and the physical and chemical behaviour of elements will be discussed in details in subsequent units.

This is a follow up to what you learned about Rutherford and Bohr models of the atom. You should have also learned that the position of electrons can be defined only in terms of the probability of finding it in a region in space referred to as orbitals. Furthermore, you learned about the four quantum numbers used for characterising an electron.

You need to be aware of how to write the orbital electronic configurations of elements based on these four quantum numbers. It has served to introduce you to orbital electronic configuration. The unit on Period Table II shall build upon this treatment of the electrons in the atoms of elements.

Senior Secondary Chemistry Textbook 3 Lagos. New School Chemsiny. Introduction The periodic table consists of elements arranged in groups and periods based on their atomic number. The number of electrons in a neutral atom of an element gives the atomic number. The electrons occupy electronic shells. The elements in any period have the same number of electronic shells and the number of valence electrons increases progressively by one across the period from left to right. When elements are arranged in this way, it was observed that the properties of elements recur at regular intervals.

That is, the properties of the elements are a periodic function of their atomic number. The periodic properties give rise to vertical columns of groups or families of elements with the same number of outer or valence elements.

The elements in any group have closely related physical and chemical properties. One great advantage of this, is that it is only necessary to learn the properties of each group rather than the properties of each individual element. Recall that the Bohr theory of shells was the basis for the arrangement of electrons in atoms.

Recent discover; r ::bout the atoms have necessitated a revision of this idea. The electrons in atoms are now believed to occupy regions in space around the nucleus called orbitals rather than fixed shells. Orbitals are simply regions in space around the nucleus where the probability of finding an electron is high, they are usually denoted by s,p;d, and f- orbitals. Recall the Pauli's exclusion principle and Hund's rule in the arrangement of electrons in the energy levels.

The summary below will refresh your memory. The modern periodic classification The Periodic Table groups atoms of the element according to their electronic configurations. Elements with one s electron in their outer shell are called Group I the alkali metals and elements with two s electrons in their outer shell are called Group H the alkaline earth metals.

These two groups are known as the s block elements, because their properties result from the presence of s electrons. Elements with three electrons in their outer shell two s electrons and onep electron are called Group III, and similarly Group IV elements have four outer electrons, Group V elements have five outer electrons, Group VI elements have six outer electrons and Group VII elements have seven outer electrons.

Group 0 elementS have a full outer shell of electrons so that the next shell is empty; hence the group name. Groups III, IV, V, VI, VII and 0 all have p orbitals filled and because their properties are dependent on the presence ofp electrons, they are called jointly the p block elements, In a similar way, elements where d orbitals are being filled are called the 61 block, or transition elements.

In these, d electrons are being added to the penultimate shell one shell before the outer shell. Finally, elements where f orbitals are filling are called the f block, and here the f electrons are entering the antepenultimate ot second shell from the outer shell shell. A summary of the block arrangement of elements based on the outermost energy levels for s- and p- block elements; and the orbitals being filled for d- and f- block elements. Instead of listing the elements, the periodic table arranges them into several rows or periods, in such a way that each row begins with an alkali metal and ends with an inert gas.

The sequence in which the various energy levels are filled determines the number of elements in each period, and the periodic table can be divided into four main regions according to whether the s,p,d or f levels are being filled. The similarity of properties within a group and the relation between the group and the electron structure is emphasized.

The d block elements are referred to as the transition elements since they are situated between the s and p blocks. Hydrogen and helium differ from the rest of the elements because there are no p orbitals in the first shell. Helium obviously belongs to Group 0, the inert gases, which are chemically inactive because their outer shell of electrons is full.

Hydrogen is more difficult to place in a group. It could be included in Group I because it has one s electron in its outer shell, or in Group VII because it is one electron short of a complete shell. Hydrogen is included in both these groups in the periodic table, although it resembles neither the alkali metals nor the halogens very closely.

The unique properties of hydrogen are largely due to the extremely small size of hydrogen atoms. Thus there is a case for placing hydrogen in a group on its own, or omitting it from the periodic table altogether.

The periodic table therefore provides an organized structure to the knowledge and understanding the chemistry of the elements. Apart from this, there is also a variation of atomic properties of elements in the periodic table.

Some of these properties are atomic and ionic sizes, ionization energy, electron affinity and electronegativity. However with the aid of modem techniques such as X-ray and electron diffraction, it is possible to determine the distance between covalently bonded atoms. For example, the distance between the nuclei of oxygen atoms in an oxygen molecule is 0.

The atomic radius or sizes of any atom is taken to be one-half the distance of closest approach between the nuclei of atoms in the elemental substance. In other words, as the atomic number increases across any period, the size of the atom decreases. Recall that as we move across a period one electron is added increasingly from one element to the next and the electrons are being added to the same shell at about the same distance.

At the same time, protons are also being added to the nucleus. Increase in the number of proton, increases the nuclear charge which progressively exert a stronger attraction upon the electrons around it and would pull them towards the nuclei. As the nuclear charge increases with atomic number across a period, the attractive force exerted by the nucleus on the outermost electrons of the atom increases hence the atomic radius or size decreases across a period.

For example, on moving from Lithium to beryllium, the number of charges on the nucleus is increased by one, so that all the orbital electrons are pulled in closer to the nucleus. In a given period, the alkali, metal is the largest atom and the halogen the smallest.

Table 8. Recall the lithium in period 2 has two shells; sodium in period 3 has three shells while potassium in period 4 has four shells. In general, as we go down the group, atomic size increases with atomic number see Table 8. The size of an ion called ionic radii is different from atomic sizes. Ionic sizes are measured in electrovalent compounds. The ionic radius of a given compound is the distance between the centre of one ion and the centre of its nearest neighbour of opposite charge.

A positive ion is formed by removing one or more electrons from an atom. When this happens, the number of positive nuclear charge is more than the number of negative electronic charge, hence the electrons are pulled in. A positive ion is therefore smaller than the corresponding atom and the more electrons removed that is, the greater the charge on the ion , the smaller it becomes e.

The number of positive nuclear charge is now less than the number of negative electronic charge hence the pull on the electrons is reduced. In general, ionic radii of negative ions are greater than the corresponding atomic radii i. Ionization energy If energy is supplied to an atom, electrons may be promoted to a higher energy level.

If sufficient energy is supplied, the electron may be completely removed, giving a positive ion. Since it is possible to remove one, two or three E 2nd I. E 3rd I. As the distance decreases, the attraction of the positive nucleus for the electron will increase, hence more energy is required to remove the outermost electron hence the ionization energy will increase.

Note that the screening effect remain almost the same across a period since electrons are added to the same shell. Table 3 shows the first ionization energies of the first twenty elements. These show a general upward trend from Li to Ne and from Na to Ar. The values for Ne and Ar are the highest in their periods because it requires a great deal of energy to break a stable filled shell of electrons.

There are several irregularities. The high values for Be and Mg are attributed to the stability of a filled s for N and P indicate that a half-filled p level is also particularly stable. The values forlevel. E 2ad I. Fig 8. First ionization energies of the elements in the first two short periods. This trend is shown with the alkali, metals from Li to Na to K The energy released when an extra electron is added to a neutral gasesous atom to form a univalent negative ion is termed the electron affinity.

Since energy is given off in the process, electron affinity has a negative value. Electron affinities depend on the size and effective nuclear charge of the atom. Moving from left to right across a period, electron affinities decreases i. Down a group of the periodic table, electron affinities increase i.

The reason for the observed trend is that atoms with smaller atomic radii tend to have a stronger attraction for electrons and thus form negative ions more readily. The tendency of an atom in a molecule to attract bonded electrons to itself is termed the electronegativity of the atom.

Generally, small atoms attract electrons due to closeness of the nucleus more than large ones and are therefore more electronegative. Atoms with nearly filled shells of electrons will tend to have higher electrcnegativity because of the desire to have a stable filled shell than those with sparsely occupied shells. The electronegativities of elements decrease down a group and increase across a period.

The reason for the trend is that down the group, atomic size increases and effective nuclear charge decreases hence electron attracting power electronegativity of the atom decreases.

From left to right of a period, the opposite effect is observed, atomic size decreases and effective nuclear charge increases, these combine to increase electronegativity. The most electronegative elements are the reactive non-metals e. Fluorine at the top right-hand corner of the periodic table while the least electronegative elements are the reactive metals e.

See Table 8. You should have also observed how this periodic properties vary down a group and across the period of a Periodic Table. You need to be aware of the reasoning behind the observed trend.

It has served to introduce you to the variation of atomic properties - atomic size and radius, ionization energy, electron affinity.

African-Fep Publishers. A chemical reaction takes place always between large number of reactant particles. The products that are formed also contain a large number of product particles.

Chemists therefore use a large number of particles as a base unit when comparing amounts of different substances reacting or are formed in chemical reactions. This basic unit is the mole and the mole concept is one of the most important concepts in Chemistry. The mole concept is applicable to all chemical processes.

In this unit the mole will be defined and the concept applied to chemical calculations involving masses and volumes of chemical substances. Now 0. The elementary particles may be molecules, atoms, ions, electron etc and must be specified.

Avogadro number has been determined experimentally and is 6. A very large number indeed. A mole of a substance therefore provides a quantity of material that can be measured for use in the laboratory.

The molar mass of a compound is the number of grams of the compound needed to make up one mole of the compound i. With this new definition of the mole you can calculate the number of i moles ii particles iii atoms etc in a given mass of a substance of known formula. Table 9. The concept of a mole is central in this type of calculations. The mass of the magnesium oxide is found to be 0. Example 2 Zinc oxide is found by chemical analysis to contain Determine the formula of zinc oxide.

Solution Assuming we analyse g sample. Example 3 2. What is the simplest formula of mercury oxide. Given that the atomic mass of oxygen is 16, determine the atomic mass of X. A balanced chemical equation of the reaction is all that is required.

The percentage yield gives the ratio of the experimental yield to a theoretical yield assuming complete reaction.

Take the Iasi example Suppose 3. This is called molar volume. Now consider the reaction in the last section. In the next unit the use of mole concept in volumetric analysis and solution preparations will be discussed.

Learn to use the mole concept and you will be in a position to solve a mole of problems. The mole concept is applicable to gas reactions as well as reactions with solid and liquid substances. The mole is a measurable quantity of substance and is more relevant to experiments in quantitative analysis. Give the number of moles of atoms in the following: 60g Carbon, 6. The percentage composition by mass of a compound of sodium, sulphur and oxygen are Water of crystallisation is 50 percent by mass.

Senior Secondary Chemistry Textbook 1. Africana-Fep Publishers. The mole concept is particularly useful in predicting yield and calculating yield from experiments. This is very important for industrial or la' oratory processes that involve reversible reactions.

Yield and yield percent calculations allow chemists to assess the efficiency of chemical processes and look for ways of improvement where applicable and possible. In this unit we continue our discussion of the mole concept. Applications in solution preparation, volumetric analysis and electrolysis will be the focus.

More calculation using the mole concept to assess the efficiency of some processes will also be discussed. A molar solution is prepared by dissolving one mole of the solute in small amount of the solvent and then make the solution up to 1. A molar solution can also be prepared from amount less than or equal to one mole of the solute.

Example It is required to prepare The solute and solvent are weighed and mixed in the stated proportions. The final volume of solution is immaterial. Note From the definitions and methods of preparation, both solutions are not the same. The preparation of solutions in volumetric titrimetric analysis is done in standard volumetric flasks and solution concentrations are expressed in mol dm-3 Not all solutions are molar solutions.

The concentration of a solution is calculated from the amount of solute and the volume of solution. Example Calculate the concentration of a solution containing 8. You must remember that though a solution has one concentration, the amount of solute will be different for different volumes of the solution. A volume of sea water will taste the same whether you test a cup of it or a bucket or a drop.

The amount of salt you recover from sea water however depends on the volume of sea water evaporated. Example Calculate the amount of sodium chloride recoverable from i 1.

Concentration of sodium chloride in the sea water is 0. Solution a i 1. From the average titre the calculation of the concentration is done using mole concept. Example Calculate the percentage purity of the sodium hydroxide sample. Mole concept allows for calculation to know how much solvent must be added to get the required concentration.

Dilution becomes the only way of making dilute solution of common acids that are available commercially as concentrated acids e. Substituting Example What volume of water must be added to cm 3 of a 0. The Faraday is quantity of charge carried by 1 mole 6. In these calculations the Faraday is used as base unit of electricity. Example 4, cov'ombs of electricity passed in an electrolysis process. The volume of hydrogen gas liberated was 0.

Calculate the efficiency of the electrolysis process. Solution 1 Faraday E 6. More applications of the concept will still be in subsequent units. It is hoped that the examples in these two units will assist you in subsequent units where you will be required to use the concept in calculations. The mole concept is used to asses the efficiency of an electrolytic process. A solution contains 2. Calculate the atomic mass of and identify M.

New School Chemistry New Edition. No one has ever seen an atom or a molecule. None of the most powerful microscopes known to us can help us see such particles. Matter exists in three states.

These states are solid, liquid and gas. Ice, water and steam are good examples of these states. Let us take a state like the solid state. The particles in the solid are tightly connected together by forces of cohesion.

The forces holding the particles of a solid restrict their movement, so that they are held in fixed positions. Solids have defmite shapes and volumes and are very difficult to compress. Liquids are hard to compress, have no definite shape but posses definite volumes. A gas occupies the whole volume of the container, has no definite shape and is very compressible.

Can you explain why if a bottle of perfume is opened at one end of a room, the smell is perceived all over the room? The kinetic theory explains the differences in the behaviour of matter in different states.

The changes that occur when matter is heated are also explained by the theory. The kinetic energy of a body is the energy it possesses as a result of its motion. The higher the velocity of a body, the higher its kinetic energy. Can you now explain why accidents with very fast moving bodies cars, stones etc are very fatal? In any given sample of matter, some molecules have very high energies while some have very low kinetic energies.

The average kinetic energy of the particles increases with increasing temperature of the matter. A suspension of sulphur powder in water when viewed under a microscope will demonstrate Brownian motion.

Brown was the first scientist to observe this behaviour, hence the name Brownian motion. Diffusion occurs in solids, liquids and gases. A drop of liquid bromine in a closed jar of air vaporises and spreads evenly throughout the jar. A crystal of a soluble coloured solid when dropped in water will after sometime colour the entire volume of water. CuSO 4. Diffusion is fastest in gases and slowest with solids. The swelling of bean seed in water is an example of osmosis.

All the above evidences confirm that particles of matter are in motion as postulated by the kinetic theory. Because of the strong cohesive force, the particles are held in fixed positions and can only rotate or vibrate about a mean position. This explains why solids have definite shapes and volumes and are very difficult to compress e.

The particles in the liquid state are further apart than in the solid. The kinetic energy of the liquid particles is higher and particles are not fixed in positions. There is some motion that allows the liquid to maintain a fixed volume but no fixed shape e g ethanol, water and kerosene.

In the gas state the particles are in constant random motion in all directions at very high velocities. There is virtually no force of attraction between the particles explaining why a gas diffuses freely filling all available space.

A gas has no definite shape or volume. The large empty spaces between the gas particles explain why gases are very compress Fig I. On the other hand gas molecules are in perpetual and random motion at very high velocities because of lack of cohesive forces between the particles.

When solid matter is heated, the average kinetic energy of the particles increases and changes in the nature of matter occurs from solid-liquid-gas At a temperature characteristic of a particular solid, the particles that are fixed in position acquire sufficient energy kinetic energy to overcome the cohesive force keeping them in fixed positions.

So the solid gradually change to the liquid form. The temperature at which this happens is called the melting point of the solid and the phenomenon is called melting. The melting point is characteristic of the solid and is often used as a criterion of purity for the solid substance.

A pure solid will have a sharp melting point i. As the temperature increases the particles acquire sufficient energy to overcome the cohesive energy of the liquid state.

The particles become free, move more randomly independent of each other. The liquid has gradually been turned to gas vapour. This is vaporisation. The temperature at which there is massive vaporisation from within the bulk of the liquid is the boiling point. At the boiling point, vapour molecules escape from the inside of the containing vessel to the surrounding space.

The boiling point is also a criterion of purity for liquid substances. Again pure liquids have a sharp boiling point. For instance the boiling point of water is C but vaporisation can take place below that temperature. This is evaporation. Evaporation is most rapid at the boiling because the liquid particles have maximum kinetic energies.

Evaporation also occurs at temperatures below the boiling point. This is most likely when a liquid sample is placed in an open container. The high energy particles on the liquid surface can vaporise into the surrounding space. The loss of high energy particles from the liquid surface will result in a decrease in the liquid volume as well as a decrease in the average kinetic energy of the liquid sample.

What is the effect of evaporation below the boiling point on liquid temperature? When the gas cools it returns directly to the solid state.

This process is called sublimation and is a useful method for separating a mixture of substances when only one of the substances sublimes, e. Fig This is a typical heating graph. This energy which is not used to raise the temperature is called the latent heat of vaporisation and the latent heat of fusion at the boiling and melting points respectively. The latent heat is used to supply the particles energy to overcome the cohesive forces in the liquid or solid state. When the substance cools the reverse changes occur.

As the vapour condenses and the liquid freezes the lalent heats are evolved. With virtually no attractive force between the gas particles you should expect the physical behaviour of gases to be much different from those of the solids and liquids.

This special behaviour of gases will be the subject of the next two units. When matter is heated change of state occurs. The kinetic theory explains this as the result of the higher average kinetic energy of the particles at the higher temperature. Melting and boiling occur at temperature characteristics of the matter.

These temperatures are called melting and boiling point respectively. Arrange A,B,C,D in the order of increasing melting point. Give a reason for your order of arrangement.

Can you recall the following as explained in that unit9 Melting point, boiling point, vaporisation and condensation. You will also recall that particles in the gas state are in random motion in all directions and at very high speed with virtually no force of attraction between the particles. The physical behaviour of a gas is very much different from those of the solid and liquid. This physical behaviour of gases was investigated by early scientists and that led to the establishment of gas laws named after them.

There is a need therefore to increase the postulates of the kinetic theory to account for the special behaviour of gases. This is the first of two units on gases and you will learn about Boyle's law and Charles' law and the general gas equation.

Statement of the gas laws will be examined and the gas behaviour as established by each law explained by the kinetic molecular theory. These assumptions are only true for an ideal gas. They constitute what is called the kinetic molecular theory. They specifically deal with the gas molecules. These following six statements describe the behaviour of an ideal gas , 1. A gas consists of small identical particles called molecules moving randomly in all directions colliding with each other and also with the walls of the containing vessel.

There is no force of attraction between the gas molecules. Molecular collisions are perfectly elastic i e no energy is lost when molecules collide with each other or with the container wall.

The vol nue of gas molecules is negligible compared to the container volume. The co. The temperature of the gas is directly proportional to the average kinetic energy of the molecules.

Note: The words in bold typeface is the statement of the law. Boyle's law describes the relationship between the pressure and volume of a gas at constant temperature. The above implies that the product of the pressure and volume is always a constant.

What this means is that as the gas pressure increases, the volume decreases and vice versa as long as the temperature is constant. Boyle's law can also be presented graphically V PV Fig. Era 'nples 1. Calculate the gas volume when the pressure is increased to mm Hg pressure at constant temperature. A gas occupies a volume of 1. What will be gas pressure when the gas expands into 2. Answers 1. You will also recall that for a fixed mass of gas at constant temperature, the number and the average kinetic energy of the molecules remain constant for the number of molecular collisions per unit area of the containing vessel to increase under a condition of fixed mass and average kinetic energy of the gas molecules, the area and therefore the volume of the containing vessel must decrease.

This implies that an increase in gas pressure is accompanied by a corresponding decrease in the gas volume. The reverse of the above is true when the gas pressure is reduced at constant temperature.

The volume of matter generally increase with increasing temperature but the increase is most pronounced for gases. This temperature at which the gas volume is theoretically zero is the lowest temperature that can be reached.

It is called the absolute zero temperature. The Kelvin temperature scale represented with a capital K has this temperature as its starting point and measures temperatures upwards from it. The Celsius and the Kelvin absolute scales are related by the equation. Note: Temperatures on the Kelvin scale are in K units with no degree sign. OR V1 - V2 where V1 is the volume at T 1 and V2 is the volume at T 2 Charles' law can also be stated in an alternative form when the gas volume is constant and pressure changes P with temperature.

Charles' law can also be presented graphically as shown below. At what temperature will the volume be 2. This will result in more random motion of the molecules. For a fixed mass of gas, the number of gas molecules is constant and at constant pressure of gas the number of collisions per unit area of the containing vessel is also constant. This implies that an increase in the temperature of the gas at a constant pressure leads to a corresponding increase in the gas volume.

In the alternative, consider the case when the volume of a fixed mass of gas is kept constant and the temperature is increased. At the higher temperature, the more random motion of the gas molecules will lead to more molecular collisions per unit area on the wall of the container and thus increase in the gas pressure.

The value and units of the gas constant R depends on the units of P, V, and T. P in atm. But when pressure is in Nm-2 , volume in in' and temperature in Kelvin, the value of R is 8. A mass of oxygen gas occupies a volume of 2 x 10 4 cm' at a pressure of 1.

Calculate the volume when the pressure is increased to 1. Solution 1. Some of the assumptions of the theory may not always be true for real gases, explaining why real gases sometimes deviate from ideal behaviour.

The quantities, pressure, temperature, volume and amount mole are important in measurements and calculations involving gases. These parameters describe the behaviour of the gas and a change in one of them will result in modification of their behaviour. The theory is used to explain Boyle's and Charles' laws.

Apart from giving the statements of the two laws they are graphically illustrated and employed in simple calculations The general gas equation was derived from Boyles and Charle's laws.

Note that in calculations involving the gas laws, the Kelvin temperature is used. Identify two assumptions in the Kinetic theory that can account for observed deviations of real gases from ideal behaviour 2. Calculate the pressure of 0. State the effect if any on i the pressure ii the volume iii the number of moles, when the gas is heated. Onitsha African-Fep Publishers. The ideal or perfect gas does not exist.

It is only hypothetical. Real gases deviate from ideal behaviour and do not obey the gas laws perfectly. Imagine a real gas subjected to low temperatures and very high pressures.

The gas molecules are compressed into a very small volume and are not very energetic. Intermolecular forces of attraction e. Vander waals and dipole-dipole attractions come into play between the gas molecules.

The volume of gas molecules will also become significant when compared with the space occupied by the gas. In this unit you will learn more about gas behavior in Dalton's, Graham's, Avogadro's and Gay Lussac's laws. These laws are also adequately explained by the kinetic theory and like the previous ones, are obeyed perfectly by ideal gases only.

Chemical Thermodynamics. Essential Electromagnetism: Solutions. Ultraviolet light and its uses. An introduction to polymer-matrix composites. Introduction to Inorganic Chemistry. An Introduction to the Quantum Theory for Chemists. Unconventional Thermocouples. Chemical Engineering Vocabulary. Glossary of Combustion. Fundamentals of Green Chemistry. Membrane filtration processes. Introductory Maths for Chemists. Learn Calculus 2 on Your Mobile Device. Fundamentals of Reaction Engineering - Examples.

Intermediate Maths for Chemists. Engineering Mathematics: YouTube Workbook. Concise College Chemistry - Part 1. Fundamentals of Reaction Engineering. Transport Phenomena in a Physical World. Introduction to Polymer Science and Technology. Chemistry: Quantum Mechanics and Spectroscopy I. Partial Differential Equations. Introduction to Chemistry. Concepts In Scientific Writing. Introduction to Vectors.

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