Chemistry pdf download

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Introduction 33 6. Electron affinity 51 8. The kinetic theory and change of state 72 Graham's, Avogadro's and Gay Lussac's Laws Monitoring the rates of Chemical reactions ' 3.

Le chatelier's principle and pressure change 5. Le chatelier's principle and concentration change 5. Effect of catalyst on reversible reactions 5. Assignment of oxidation number 9. Objectives Definition and scope of organic chemistry 1. Uniqueness of carbon atom Representation of organic molecules 1. Classification of organic compounds 1. Open-chain or aliphatic compounds 1. Many materials that we use everyday, directly or indirectly are products of chemical research and examples of useful products of chemical reactions are limitless.

What then is matter? Matter is anything that has mass and occupies space. Matter is classified into solid, liquid and gas. Glass, sand and most metals are examples of solids. Water and kerosene are examples of liquids.

Air and cooking gas are examples of gases. The above classification is commonly referred to as physical classification. Matter can also be classified into elements, compounds and mixtures. This latter classification is referred to as chemical classification. It is remarkable that all these substances, solids, liquids, gases, elements, compounds and mixtures are built up from simple basic units.

What is the basic unit of matter? What are the building blocks of matter and what laws govern the interaction of matter? The above are some of the questions that will be answered in this unit. A pure substance that can be broken down into elements is called a compound. There are over chemical elements. Some occur naturally as free elements, mixed with other elements or compounds.

Some are very rare, while most occur in combined state in compounds. The table 1. Table 1. Examples of metals are copper and iron. The general characteristics of metals are lustre, good conductor of heat and electricity.

Metals can be rolled and hammered into sheets and drawn into wires. They are used for roofing and electrical cables respectively. All metals are solids at room temperature except mercury which is a liquid at room temperature.

About 75 percent of the elements are metals. Unlike metals, non metals do not have characteristic lustre. Many are gases at room temperature and others are solids except bromine which is a red brown liquid at room temperature. Non metals are non-conductors of heat and electricity. They cannot be rolled into sheets or drawn into wires like the metals. Oxygen, nitrogen, carbon and iodine are examples of non metals.

The properties of compounds are different from those of the elements from which they are formed. A lot of energy is often required to split compounds into the constituent elements. There are limitless number of compounds. Sodium chloride, common salt , water and calcium trioxocarbonate iv , marble are examples of compounds.

The composition of a mixture varies and the components are separated by physical methods. Such physical methods include, heating, cooling, dissolution, filtration and distillation. Air and petroleum are examples of mixtures. For example both will rust when exposed to air and moisture and both will conduct heat and electricity.

The atom is the smallest unit of an element that can take part in a chemical reaction. It is the smallest unit of the compound that can take part in a chemical reaction. The atom is to the element as molecule is to the compound. The basic unit of matter in chemical reactions is the atom. Atoms and molecules are the building blocks of matter. A theory serves as a guide to new experiments.

When proved incorrect or inadequate by experiment, a theory is discarded or modified so that new experimental facts can be accounted for. This latter statement is true of Dalton's atomic theory. Dalton's atomic theory form the basis of theory of the atom.

It has been modified in the light of new experimental facts about the atom. This implies that the atom is divisible. The existence of atoms of the same element having different masses have been proved with the use of an instrument called mass spectrometer. Atoms of the same element having different masses are called isotopes. Irrespective of masses, atoms of the same element have same chemical properties.

Evidences for this assertion came from results of experiments of early scientists like Faraday, Thompson and Millikan. The negatively charged particle in matter is the electron. It has negligible mass. The proton is the positively charged particle. It carries the same magnitude of charge as the electron and is very much heavier than the electron.

The third particle is the neutron, a neutral particle with a mass approximately equal to that of the proton. These three particles are constituents of the atom except the hydrogen atoms that do not contain neutrons. The number of each particle present in the atom varies from one element to another. For the atom matter to be electrically neutral, the number of protons must equal the number of electrons.

Atoms of the same element will have the same number of protons and electrons but may have different numbers of neutrons. Such atoms will have different masses and are called isotopes. The measurement of masses of atoms is not possible because they are very small. Their masses can however be compared to give relative atomic and molecular masses of elements and compounds.

On this scale 1 atom of carbon isotope is given a mass of 12 atomic mass units. With the use of the mass spectrometer it has been possible to determine fairly accurately the relative atomic masses of elements. The relative molecular mass of a compound is the sum of the relative atomic masses of the elements present in the chemical formula of the compound. A chemical equation uses these symbols and formulae to summarise a chemical reaction.

Chemical symbols consist of the first one or two letters of the name of the element. Some symbols do not correspond with the elements names; these symbols are derived from the Latin names of the elements.

It is important that you know the symbols for as many of the common elements as possible. The Latin names for copper, lead and gold are in bracket. The formula of a compound gives the proportion of the different elements present in the compound by mass. By the law of constant composition the proportion by mass of the different elements present is fixed for a pure sample of the compound irrespective of the method of preparations.

Example 1. Having determined the relative atomic mass of the elements that make up the compound, you can proceed to determine the empirical formula of the compound. Since the relative masses of the elements in a compound depend partly on the masses of the atoms and also on the relative number of each atom of each element involved in the combination to form the compound.

Determine the formula of the compound. Now determine a formula of a compound of calcium, carbon and oxygen with 40 percent Ca, and 48 percent oxygen. Two more laws will be discussed in this section.

A consequence of this law is that a chemical equation must always be balanced to account for all atoms present on the reactant side, and on the product side of the reaction. CO and CO 2 or SO2 and S03 the masses of one element which react with a fixed mass of the other are in a ratio of small whole numbers.

Show that this is in agreement with the law of multiple proportion. For the 1st compound 1. Ratio of 0 mass combining with 1. In the reshuffling, one compound is converted to another.

The smallest units that can take part in chemical reactions are the atoms and molecules. A chemical equation is often used to summarise the reaction that has taken place. A chemical equation gives the reactants and products of the reaction and the quantities of the reactants and products in correct ratio in accordance with the law of conservation of matter. Sometimes chemical equations give the physical states of the reactants and products. The elementary particles that form the basic units of elements and compounds in chemical reactions are atoms and molecules.

Though atoms are composed of more fundamental particles, they are not split in chemical reactions. Dalton's atomic theory is the basis of modem atomic theory and explains satisfactorily the laws that govern chemical reactions of elements and compounds.

The atom has a substructure of its own. This is the subject of the next unit. They are neither created nor destroyed. A chemical equation must be balanced for the law of conservation energy to be satisfied. Chemistry is a quantitative science and as such, amounts of matter used in reactions must be known. Calculate the relative atomic mass of chlorine element. Senior Chemistry Textbooks I and 2. Lagos Longman Publishers. Osei Yaw Ababio New School Chemistry. Onitsha Africana-Fep Publishers.

Dalton's theory proposed that atoms are indivisible units of matter. The atom is the smallest unit of matter that can take part in a chemical change. Dalton's atomic theory satisfactorily explained the laws of chemical combination but could not explain why substances react the way they do. Why is oxygen able to react with maximum of two atoms of hydrogen as in water? Why do some elements exist only as diatomic molecules? Why are some elements very reactive and some inert?

Dalton's law could not explain electrolysis neither could it explain the different masses of atoms of the same element. Today we believe that the atom has a substructure of its own. The atom consists of much smaller particles that we call protons, neutrons and electrons. The relative masses and charges of these particles are given in unit 1. What are the evidences in support of this new picture of the atom? How many of these particles are present in the atom of elements?

How are these particles arranged in the atom? These are some of the questions that will be answered in this unit. This evidence was closely followed by the discharge tube experiment.

A heated metal cathode emitted negatively charged particles. This beam of particles is called cathode rays and the particles are the electrons. Thompson worked on cathode rays and confirmed that they are negatively charged. Their charge mass ratio was determined and found to be —1. Millikan determined the electronic charge in his famous oil drop experiment in The charge of the electron is —1. The mass of the electron was calculated. It is 9. The existence of positively charged particles was confirmed in a modified version of the set up used by J.

Thompson by another scientist by name Goldstein. The proton is about times as heavy as the electron and carry a charge equal but of opposite sign to that of the electron. Protons and electrons are present in all atoms. The evidence of radioactivity by H. Becquerel further demonstrated the existence of subatomic particles.

Becquerel observed that certain substances spontaneously emit radiations. The most important of these radiations are the alpha, beta and gamma radiations. Chadwick later confirmed the existence of a neutral particle in the atom and called this a neutron.

This neutron has a mass approximately equal to that of the proton. Table 2. The atomic number is the number of protons in an atom of the element and for a neutral atom, the atomic number is also the number of electrons. The sum of the number of protons and neutrons is the mass number. Isotopes are atoms of the same element with different mass numbers.

We shall see in the subsequent units that the chemical property of an element depends on the number as well as the arrangement of the atomic electrons. This explains why atoms of the same element with different masses have the same chemical properties. I2 C and 14c These are isotopes of carbon, mass number 12 and 14 and neutron numbers 6, 6 6 and8respctivly.

It was necessary to propose a model for the atom. Thompson proposed that the atom could be viewed as positive matter in which electrons are uniformly distributed to make it neutral at every point. This view was dropped because of the findings of two other scientists, Rutherford and Marsden. They bombarded a thin gold foil with fast moving alpha particles. They found that most of the alpha particles pass through the foil undeflected. Some were deflected as large angles while very few were sent back on their paths.

Alpha particles are positively charged and many of them passing through the foil undeflected suggested that most of the gold foil was empty space. He proposed that the atom consisted of a tiny positively charged nucleus. The nucleus is centrally placed in the atom and the electrons surround it. The very small number of deflections of alpha particles suggest that the nucleus occupies a very small portion of the atom.

For heavy particles such as the alpha particle to be so deflected suggests that the nucleus is a centre of heavy mass and positive charge. The protons and neutrons occupy the nucleus while the electrons are arranged around the nucleus and move in orbits around it, as planets around the sun.

This is Rutherfords nuclear model of the atom. When a voltage is applied across electrodes that are sealed in a partially evacuated glass tube, the space between the electrodes glows. This is not observed in practice. Intact energy absorption and emission by elements is discontinuous. Have you ever seen the rainbow in the sky? The colours you see range from violet to red with no sharp line separating one colour from the other.

This is a continuous spread of colours and is called a continuous spectrum. The different colours are component colours of light. In the laboratory the separation of light into its component colours also happens when light passes through a glass prism. Light from the vapour of an element does not give a continuous spectrum. Each element has its own characteristic bright lines in particular positions.

This is a line spectrum suggesting that light energy absorption or emissons by elements is only at particular energies characteristic of the element. On the basis of the above observations Niel's Bohr proposed a model for the atom in which electrons move round the nucleus only in allowed orbits numbered serially.

The orbit closest to the nucleus is assigned number 1 and is the orbit of lowest energy. Allowed transitions are transitions from one orbit to another and will lead to emission or absorption of the energy difference between the orbits. In the light of Bohr's model, there are electronic energy levels in atoms corresponding to different orbits of electron motion that are allowed.

These energy levels are sometimes referred to as electron shells and designated K, L, M, N etc. Bohr's model gave satisfactory explanation of the hydrogen spectrum. The theory is limited however in its explanation for multi electron atoms.

The wave mechanical treatment of the atom overcomes this limitation of Bohr's theory and is the subject of a subsequent unit. N etc. When 8 electrons are accommodated into the M shell however, there is extra stability and the next 2 electrons go into the N shell.

Subsequently any electron goes into the M shell until it contains the maximum of 18 electrons. Study the above table and note the following a He and Ne; each has full shell of electrons and are both inert gases. These pair of elements with similar electron arrangements also have similar chemical properties. Now give the electron arrangement for each of the following: Carbon 6 Phosphorus 15 and potassium 19 Which of the elements in the table will have similar chemical properties with phosphorus.

Note Isotopes have the same number of electrons and therefore have the same electron arrangement. Their chemical properties are the same.

The proton number is not affected by ion formation but the electron number increases or decreases. This observation and the full shell arrangements for the inert gases, suggest that a full shell electron arrangement is a stable electron arrangement. This is supported by results of experimental investigations by early scientists. Three fundamental particles are present in the atom.

The positive charge and mass are in a small centre of the atom. This centre is the nucleus and is surrounded by electrons moving round in allowed orbits of fixed energy. Electron assignment to these energy levels show that elements with similar properties have similar electron arrangement.

The noble inert elements have closed shell electron arrangement suggesting that a closed shell arrangement of electrons is a stable configuration. This observation is used to explain chemical bonding. Some of these evidences are revealed and discussed in this unit. This model was modified by Bohr who proposed electronic energy levels, sometimes referred to as electronic shells around the nucleus. Stable ions also have closed shell electron arrangement. The atomic numbers are in brackets Be 4 Mg 12 K 19 Si 14 and Cl 17 b Which of the listed elements in a above are likely to: i be metals.

Senior Secondary Chemistry Textbook 2. Longman Publishers. Osei Yaw Ababio. Africana-FEP Publishers. The number of positive charges in the nucleus of a neutral atom determines the number of electrons that surround the nucleus.

The arrangement of the electrons around the nucleus determines the chemical properties of the element. Presently there is no theory that predicts the stability of the nucleus but a number of empirical observations suggest that the presence of neutrons account partly for nuclear stability. Except for the hydrogen atom, all nuclei contain neutrons. As the proton number increases, the neutron number also increases. For heavy elements the neutron number far exceeds the proton number in the nucleus.

Can nuclei react? What happens when a nucleus is unstable? Can a stable nucleus be made unstable? These are some of the questions to be answered in this unit.

According to Rutherford the nucleus is about 11I00, times the size of the atom. Except for hydrogen all elements have more than one proton in their nuclei. From your knowledge of elementary magnesium in Physics, you learnt that like charges repel each other whereas unlike charges attract each other. As mentioned in the introduction, no theory predicts nuclear stability. Empirical observations suggest that neutrons are partly responsible for nuclear stability.

Table 3. Contrary to the postulate of Dalton, an element is destroyed and a new one is created, when a radioactive element like uranium disintegrates. Like in simple chemical reactions symbols are used to summarise the nuclear reaction in an equaiton. A nuclear reaction equation must be balanced to account for all particles and charge. It breaks down emitting particles of Thorium and Helium. This is called radioactivity and is an example of a nuclear reaciton.

Very heavy metals with atomic number greater than 83 are radioactive. Some isotopes of light elements are also radioactive. A radioactive isotope is called a radioisotope. Examples of radioisotopes are 14 92U and 6 Radioisotopes decay at different rates. The half-life is a measure of the stability of a radioisotope. They are very stable compared to 21 3i Na and 84 p0 half lives of 60 seconds and microseconds respectively. Some stable isotopes of elements can be made radioactive.

It helps to keep 14r, activity constant in the atmosphere. These radiations are mainly of three types. Decay of light radio isotopes is usually accompanied by one or at most two of the three radiations. Nuclear fusion reactions require very high temperatures. The energy from a fusion reaction initiates more reactions.

If not controlled, can lead to explosion. The reaction releases three neutrons which If the chain reaction is not can initiate more fission reactions. This is an example of a chain reaction. The energy from fission reactions are not as high as in fusion and the reactions in fission do not require very high temperatures to initiate.

There are fission reactors used in generating electricity. Nuclear transformation implies changing one element to another by the reactions of atomic nuclei. These reactions are many. Many isotopes of elements have been prepared by this method and used in chemical, biological and medical researches. Some heavy elements have been produced by the radioactive action of alpha particles, i.

Helium atom. This is one use gamma radiaiton is put into. F-radiation also destroys healthy cells as well and too much exposure to it can do more harm than good. The extent of damage depends on the energy and type of radiation. The effect of radiation is also cumulative and small doses over a long period of time will also cause serious damage to biological systems.

Radioactive waste is very dangerous and must be disposed properly to avoid necessary exposure to its hazards. Nuclear reactions unlike chemical reactions which involve valence electrons, nuclear reactions involve protons and neutrons. Nuclear reactions are much more exothermic than chemical reactions.

Many atomic nuclei are unstable. Some occur naturally and some are man-made. Unstable nuclei emit radiations with characteristic properties. The emitted radiations find application in various fields of human endeavour but also pose danger to users and non users alike. Radioactive waste must be properly disposed to avoid unwanted effects. Radioactive materials must always be handled with care.

The radiations have properties that make them detectable. Osei Yaw Ababio, The forces that hold atoms together in compounds are called chemical bonds. The combination of chemical elements to give a compound is a chemical reaction. Some elements are very reactive and exist in nature only in combined states, e.

Few elements are relatively imreactive and exit rarely as free elements. They are called noble or rare elements, e. Most elements have intermediate reactivity and exist as free elements as well as in chemical compounds e.

There are some non metallic elements that exist only as diatomic molecules in the free state. These elements also occur in combined states. In the previous unit, you were shown that arrangement of electrons in atoms showed some correlation between electron arrangement and properties. Li 2, 1 , Na 2,8,1 , K 2,8,8,1 all have similar configuration with one electron each in their outermost shell.

They are metals. F 2,7 and C1 2,8,7 all need one electron to complete their outermost electron shell. They are non-metals. The inert or noble elements He 2 , Ne 2,8 , Ar 2,8,8 all have complete shell arrangement of electrons. The electron arrangement in stable ions of metals and non metals also show that complete shell of electrons is a stable configuration e.

In chemical bonding therefore elements tend to attain the noble or inert gas configuration. The outermost shell electron arrangement is therefore very important in determining the type of bond.

Electrovalent bonding involves electron transfer from the valence shell of one atom to the valence shell of the other. One atom loses electrons to become positively charged and the other gains electrons to become negatively charged. The positively and negatively charged ions are called cations and anions respectively.

The ionic bond results from the attraction between these oppositely charged ions This type of bonding is usually between metals and non-metals. F: The formula of LiF written as above is the electron dot formula Lewis structure. The brackets around the fluorine are intended to show that all eight electrons are the exclusive property of the fluoride ion F. Mother example is the bond between sodium and chlorine.

Except for helium, He 2 the inert gas configuration corresponds to eight electrons in the outershell. The electronic theory of valency as postulated by Kossel and Lewis was prompted by the remarkable stability of the rare gas elements.

This stability is associated with the presence in the atoms of a group of eight electrons in the outer shell. This completeness appears to be the source of stability in rare gases. The tendency for atoms to have eight electrons in their outermost shell is explained by the octet rule. Note that the rule does not always hold In cases like these, other stable configurations explain ion stability.

The number of bonds to a particular atom depends on the number of electrons gained or lost to attain stable configurations for example.

Write the formula of the compound. Ionic compounds are usually solid a its: sting of regular arrangement of equal number of positive and negative charges. For Lif and NaCI there will oe equal numbers of cations and anions.

This regular arrangement of cations and anions in the solid crystal is called the lattice. The structure of sodium chloride is illustrated in the Fig 4. It is easier to explain the binding forces in the union between sodium ion and chloride ion, in the formation of sodium chloride since their opposite ionic charges attract each other.

Its however difficult to comprehend the manner of bondage between non ionic or non polar atoms. Its however difficult to comprehend the manner of bonding between non ionic or non polar atoms. Lewis, , came up with a tenable explanation, suggesting that non ionic molecular compounds arise from the sharing of electrons among atoms, resulting in a form of bonding which was called the covalent bond.

This type of bonding involves sharing of electron pairs rather than complete transfer. The binding force results from the attraction of the shared electron pairs by the nuclei of the atoms involved in the bonding. Look at the following examples. The number of electron pairs shared depend on the number of electrons each atom must share to attain an inert gas configuration.

Covalent compounds form molecules and depending on the intermolecular forces between the molecules they may be gas 02, H2 HCI or liquids Br2 H2 0 or low boiling solids candle wax. HCI 4. This is not the case with co-ordinate covalent bond. The electron Pair is attracted by both nuclei of the bonded atoms. Lone pairs are electron pairs that are not used in bonding to other atoms The other atom must have a vacancy in its valence shell to accept the lone pair.

The bond formation also results in inert gas configuration for both atoms. Co-ordinate covalent bonding is common with metal complexes. The molecules donating the electron pairs are called ligands and the metal ion the central atom. The formation of the ammonia complex with copper ions in solution. It is a complex ion and has charges located on a group of atom. Electrovalent compounds consist of cations and anions in their solid structure When an ionic solid dissolves in water or is melted, these ions become free.

This explains why ionic compounds are good electrolytes when molten or in solution. Most covalent compounds are gases at room temperature because they consist of molecules held together by weak intermolecular forces.

Dative bonding is an important type of bonding that helps to explain the structure and properties of additive compounds and complex ions. Water and ammonia have lone electron pairs and take part in dative bonding with the hydrogen ion. Dative bonding is a special type of covalent bonding. Senioir Secondary Chemistry Textbook 1, Lagos.

Each bond type gives charabteristic properties to the compounds that are formed. 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.

Essential Electrodynamics: Solutions. Basic Physical Chemistry. Statistical Thermodynamics and Spectroscopy. Fundamentals of Chemistry Part II. Electrically Driven Membrane Processes. Essential Descriptive Inorganic Chemistry. Micro- and Nano-Transport of Biomolecules. Atomistic Models. Essential Electromagnetism: Solutions.

Ultraviolet light and its uses. Chemical Engineering Vocabulary: Bilingual. Inorganic and Applied Chemistry. Chemical Thermodynamics. Unconventional Thermocouples. Chemical Engineering Vocabulary. An introduction to polymer-matrix composites.

Introduction to Inorganic Chemistry. An Introduction to the Quantum Theory for Chemists. Glossary of Combustion. Fundamentals of Reaction Engineering - Examples. Learn Calculus 2 on Your Mobile Device. Introductory Maths for Chemists. Intermediate Maths for Chemists.

Engineering Mathematics: YouTube Workbook. Concise College Chemistry - Part 1.

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