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Periodic Table And Classification

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Periodic Table and Classification with Summary.

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    HE PERIODIC TABL ASIC CONCEPT A vertical column of elements in the Periodic Table is called a Group. A horizontal row of elements in the Periodic Table is called a Period. An element that has the properties of both a metal and a non-metal is called a Metalloid. The Group I elements in the Periodic Table is called Alkali Metals. The Group VII elements in the Periodic Table is called Halogens. The Group O or VIII elements in the Periodic Table is called Noble Gases. A block of metallic elements between Groups Il and Ill in the Periodic Table is called Transition Elements. A substance that increases (enhances) the speed of a chemical reaction is called a Catalyst.
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    ANALYSE: v/ Describe the Periodic Table Relate the position of an element in the Periodic Table to its portion number and electronic structure Recognize the similarities between the elements in the same group of the Periodic Table in terms of their electronic structure Periodic table is defined as The arrangement of known elements in certain groups and periods in such a way that elements with similar properties are grouped together." The early attempt to classify elements meaningfully was done by Dmitri Mendeleev, a He classified the elements based on the periodic law. Mendelee'ds Russian chemist in 1869. periodic law states that "the physical and chemical properties of the elements are a periodic function of their atomic masses". In Mendeleev's periodic table when elements were arranged in the increasing order of their atomic masses the physical and chemical properties of elements got repeated after definite intervals. The elements in vertical rows were called and the elements in horizontal columns were called periods. In Mendeleev's table groups there were nine groups and seven periods. About 60 elements known at that time were arranged in the periodic table. In 1913, Moseley by his studies on x-ray diffraction pointed that it is the atomic number not the atomic mass is responsible for the periodicity of chemical properties. He proposed the Modern Periodic Law. It states that: "The chemical properties of the elements are the periodic function of their atomic numbers". 1.1 EATURES OF THE PERIODIC TABLE 1. The list of elements arranged in order of their increasing Proton(Atomic) Numbers in a table is called Periodic Table . A Period is a horizontal row of elements and a group is a vertical column of elements.
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    2. The Periodic Table consists of Seven Periods and Eight Groups of elements. Ill IV V VI Vll 2 3 4 5 6 7 TRANSITION ELEMENTS oooooooooooooooooo oooooooooooooooooo • oooooooooooooooooo 3. Periods: a) The Periodic Table consists of seven periods of elements, numbered from 1 to 7. b) The Period run horizontally from left to right. 4. Groups: a) The Periodic Table has eight groups of elements, numbered from 1 to O. b) They run vertically from top to bottom. c) Group O is also called group VIII. TIPS FOR STUDENTS: With the exception of beryllium, elements in groups I and Il form ionic compounds. Elements of Groups V, VI and VII in Periods 2 and 3 can form both ionic and covalent compounds. For example, Na2 is an ionic compound but H2S is a covalent compound. 5. The block of element between Group Il and Group Ill is called the transition
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    elements or transition metals. 6. The position of an element and its electronic configuration a) Elements in the same group have the same number of valence electrons. This number is the same as the group number. For example: oxygen and sulfur from Group VI have six valence electrons. b) Elements in the same group have similar chemical properties. c) The table below shows the relationship between group number and electronic configuration. GROUP NUMBER OF VALENCE ELECTRONS EXAMPLE 1 1 11 2 111 3 B 4 c 5 6 Vll 7 Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine 8 (except valence electron Neon d) The electronic configuration of an element can be worked out if we know the position of the element in the Periodic Table. e) The group number of an element tells us the number of valence electrons. The period in which the element is placed tells us the number of shells of electrons in the electronic configuration.
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    ADDITIONAL PORTION: Periodicity is the repeatition of similar properties of elements placed in a group after of atomic numbers. The periodic table given by Moseley is called "long form of periodic table." At present 118 elements are known and they are arranged in 18 groups and 7 periods in the modern periodic table. In long form of periodic table the roman numerals for the groups are now replaced by groups 1 to 18 according to the recommendation made by International Union of Pure and Applied Chemistry ( IUPAC) XAMPLE An element in Group V and Period 3 has the following electronic configuration: Five valence electrons: Group V element Electronic configuration: 2, 8, 5 Four electron shells: Period 3 elements f) Conversely, from the electronic configuration of an element, we can work out the position of the Periodic Table . XAMPLE An element with the electronic configuration (2, 8, 8, 1) is a Group I and Period 4 element. One valence electrons: Group I element Electronic configuration: 2, 8, 8, 1 Four electron shells: Period 4 7. The position of an element and its proton If the proton number of an element is given, we can work out the period and group to which this element belongs .
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    An element has the proton number 16. Its electronic configuration is ( 2, 8,6 ). Six valence electrons: Group VI element Electronic configuration: 2, 8, 6 Three electron shells : Period 3 TEST YOURSELF: v/ Recognise the change from metallic to non-metallic character of the elements from left to right across a period of the Periodic Table v/ Relate the group number to the number of valency electrons and metallic or non- metallic character of the elements v/ Relate the group number to the ionic charge of an element 1.2 ERIODIC TRENDS 1. 1 2 3 4 5 The bold line in the Periodic Table divides the metals from the non-metals. KEY: Line dividing metals and Non - metals 1 Na METALS NON - METALS 11 Be Mg V Cr 111 Al In c v N AS Vll s Ar c Ni Cu
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    6 Fr Ra 7 2. Metallic and non-metallic character across a period. a) Elements on the left of the periodic Table are metals, while those on the right are non- metals. b) From left to right across a period, there is a decrease in the metallic properties and an increase in non-metallic properties . c) A metalloid is an element with properties that are intermediate between those of metals and non-metals. Examples are silicon, Si (Period 3), and germanium ,Ge (Period 4). d) If the proton number or electronic configuration of an element is given, we can predict the metallic or non-metallic character of the element of the element and hence, predict the properties of the elements. 3. Nature of oxides: a) In general, i) metals form basic oxides, ii)non-metals form acidic oxides and iii)metalloids form amphoteric oxides . b) However, there are exceptions. Some metals such as aluminium form an amphoteric oxide, A1203, and the metalloid, silicon, form an acidic oxide, Si02 . 4. Group number and the ionic charge of an element a) In The table below shows the relationship between group number and the ions formed by the element .
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    b) The elements in Group I, Il, and Ill are metals. They form ionic compounds. c) In the ionic compounds, the metals lose electrons to form positive ions. The charges on the ions of the elements in Group I, Il, and Ill correspond to the group numbers . d) The elements in Group V, VI, and VII are non-metals. They tend to form covalent compounds when they combine with non-metals . e) However, when they combine with Groups I and Il metals, they form ionic compounds. f) In the ionic compounds, the non-metals form negative ions. The charge on the ion corresponds to group number minus eight . XAMPLE Nitrogen is a Group V non-metal. The charge on its ion corresponds to: (5-8) Group number That is, N3- . g) The element in Groups IV, V, VI, and VII form covalent bonds by the sharing of electrons.
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    COMMON ERROR X Group IV elements are all non- metals. ACTUAL FACTS v/ On going down a group, the metallic property increases. For example, in Group IV, carbon is a non-metal, silicon and germanium are metelloids,but tin and lead are metals. 5. The diagram below summaries the trends of the propreties of the elements in the Periodic Table oooooooooooooooooo oooooooooooooooooo ACROSS A PERIOD: Metallic properties decrease Non - Metallic properties increase ON GOING DOWN A GROUP:
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    Proton number increases Size of atom increases Metallic properties increase ERIODIC PROPERTIES Periodic properties of elements are those properties which are directly or indirectly related to the electronic configuration of the elements and show gradation in moving down the group or_along the period. Some of the important periodic properties and their trends are discussed below. TOMIC RADIUS AND IONIc RADIUS ATOMIC RADIUS: Atomic radius is defined as "the distance between the centre of its nucleus and electron in the last orbit." The exact value of atomic radius is difficult to determine because it is not possible to determine the exact position of an electron. In the atom, some times the electron may be very close to nucleus and sometimes it may be far away from the nucleus. There are 3 forms of atomic radius based on the nature of the bonding. They are COVALENT RADIUS VANDER WAAL'S RADIUS METALLIC RADIUS COVALENT RADIUS: In homoatomic molecules having same type of atoms the covalent radius may be defined as half the distance between the centres of the nucleus of two adjacent atoms which are similar joined by a singal convalent bond
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    Inter nuclear distance in single bond COVALENT RADIUS = 2 74 pm 37 pm FIG. COVALENT RADIUS IN 1-12 XAMPLE the inter nuclear distance between two hydrogen atoms is 74 pm. Therefore its covalent • +37 pm. radius IS VANDER WAAL'S RADIUS: It is defined as the " One half the inter nuclear distance between two similar atoms belonging to two neighbouring molecules of same substance in the solid state
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    The Vander Waal's forces of attraction is very weak in the gases and liquids hence it can only be determined in the solid state. The magnitude of force is more hence the Vander Waal's covalent radius is always more than the corresponding value of covalent radius. For example Vander Waal's radius of chlorine is while its covalent radius is 90 pm . 180 pm METALLIC RADIUS: It is half the distance between the centres of the nuclei of two adjacent metal atoms in a crystal. Since the metallic bond is weaker it must have more radius when compared to covalent radius. ARIATION OF ATOMIC RADIUS IN THE PERIODIC TABLE The atomic radii (covalent radii) of elements generally decrease from left to right in a period, while when we move down the the atomic radii increases. group 1. In the Second period elements the covalent radii decrease from lithium to Fluorine. It is given below. ELEMENT COVALENT RADIUS IN PM 2. In the group I ELEMENT COVALENT RADIUS IN PM 123 Be 90 80 c 77 N 75 o 74 72 alkali metals the variation of covalent radius is given below. 123 157 202 217 235
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    REASONS: On moving left to right across the period the effective nuclear charge increases and electrons are strongly attracted towards the nucleus from one element to another. This strong attraction is mainly because of the entry of new electron to the same shell. Thus the covalent radius decreases. As we move down a group there is increase in the number of shells and atomic size increases. The added electron enters the new shell. Even though a proton is also added, the effective nuclear charge is less, on the added electron in the new shell. Therefore the atomic radius increases down the group. ONIC RADIUS Ions are the charged particles which are formed by the loss or gain of one or more electrons from a neutral atom. A cation has positive charge while an anion has negative charge. Ionic radius is defined as follows. "It is the distance between the centre of the nucleus to the region up to which it can attract the electron cloud Since the electron cloud is extended to larger distance from the nucleus it is not possible to determine the ionic radius experimentally. Pauling by his x-ray diffraction ,suggested that if the radius of one of the ions is known, then the ionic radius of the other can be calculated by subtracting the value of known ionic radius, from the inter nuclear distance. The radius of Na ion is 95 pm, then the radius the inter- nuclear distance is In NaCl , 276 pm. of Cl-is 276-95 = 181 pm. The radius of a cation is always smaller than the neutral atom from which it is formed. This is because when an electron is removed from the outermost shell the neutral atom has become unipositive ion. It means the nucleus gains a proton. Now the nucleus can attract the remaining electrons more strongly, hence the electron cloud shrinks, resulting in decrease in ionic radius. It is shown below
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    XAM ELEMENT ATOMIC RADIUS (PM) IONIC RADIUS OF CATION ( pm) Li 123 60 Na 157 95 202 133 Mg 136 65 Ca 114 99 125 50 The radius of Anion is always larger than the neutral atom from which it is formed. This is because the added electron is less attracted by the nucleus and it also experiences shielding effect by the inner electrons. Thus it results in expanding of the electron cloud. This fact is given in the following table. ELEMENT ATOMIC NUMBER IONIc RADIUS OF CATION ( pm) SOELECTRONIC IONS o 74 142 75 171 99 136 114 181 133 196 125 219 They are the species of different ions which have same number of electrons but different magnitude of nuclear charges. If the ion has greater value of nuclear charge it has lesser ionic radius. The size of isoelectronic ions decrease with increase in atomic number. It is shown in the following table.
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    ELEMENT ATOMIC NUMBER NUMBER OF ELECTRONS IN THE ION IONIC RADII ( pm) 7 10 171 02 8 10 140 Na+ 11 10 95 Mg2+ 12 10 65 A 13+ 13 10 50 ONISATION ENERGY OR IONISATION POTENTIAL It is defined as " The minimum amount of energy needed to remove the most loosely bound electron from a neutral gaseous atom to form a positive ion Ionisation energy is denoted as Na(g) + IEI The energy needed to remove the first electron form a neutral gaseous atom called first ionisation potential IEI . The ionisation energies required to remove second and third electrons are IE2 and IE3 respectively. The ionization energy in kJ mol-I. The first ionization energy is always less than the second because, when one electron is removed from the neutral gaseous atom forms a unipositive ion The extra proton in the nucleus attracts the remaining electrons more strongly . Therefore it is difficult to remove second electron as more energy is required to overcome the nuclear which is given as follows. attraction. In summary IE3 > > IEI ELEMENT ATOMIC NUMBER IEI kJ mol 1 Li 3 520 Be 4 900 5 800 c 6 1086 N 7 1402 o 8 1314 9 1681
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    IE2 kJ mol 1 IE3 kJ mol 1 7297 11810 1757 14850 2427 3638 2352 4619 2858 4576 3388 5296 3375 6045 ariation of IE in the eriodic table Ionisation energy is expected to increase on moving across the period and the value decrease down the group. am le 1. In the second period the El values increases as we move from left to right of the periodic table. ELEMENT IEI( kJ mol-I ) 2. In the I group ELEMENT IEI ( kJ mol-I ) 520 Be 899 B 800 c 1086 N 1402 o 1314 1681 Alkali metals, the first ionisation energy values decreases. 1312 520 Na 496 419 403 2080 475 In a period we notice that noble gases have highest value while the alkali metals have lowest values. REASONS: On moving along a period, the electron enters the same shell. It leads to increase in effective nuclear charge and decrease in atomic size. This results in increase in the forces binding the
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    electron with the nucleus. Hence IE increases, when we move down a group the IE decreases because the new shell is added and electron is present far away from the nucleus. It is less attracted by the nucleus. Moreover the shielding effects (Repulsion by inner electrons) causes the electron less tightly bound to nucleus, it requires very less energy to pull the electron. Therefore the ionisation energy decreases. The variations of IEI for second period and first group are represented in the following graphical representation. 2000 1500 IEI( KJ mol-I) 1000 899.5 Li(5 0.3) 500 2 4 Na ( 2080.4 ) F ( 1680.8 ) N ( 1402.1 ) o ( 1313.7) c ( 1086.2 ) B ( 800.6 ) 6 8 10 ATOMIC NUMBER ( Z )-....................+ FIG. VARIATION OF IE ACROSS SECOND PERIOD 550 500 450 IEI( KJ mol-I) 400 Li ( 520.3 ) Na ( 495.8 ) K ( 418.9 ) Rb ( 403 ) Cs ( 374.3)
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    350 5 15 25 30 35 40 45 50 55 ATOMIC NUMBER ( Z ) FIG. VARIATION OF IE DOWN THE FIRST GROUP LECTRON AFFINITY It is defined as " The energy released when a neutral gaseous atom accepts an electron to form an anion Electron affinity is denoted as EA Similar to ionisation energies, when a uninegative anion is formed it is called first electron affinity EAI The EA1energy released while EA2 energy is absorbed. This is because a uninegative anion has to take up electron cloud against the repulsion by the electron could. Therefore energy must be supplied. VARIATION OF ELECTRON AFFINITY IN THE PERIODIC TABLE Electron affinities normally increase along a period and decrease down the group with few exceptions. XAMPLE 1. In the elements of second period, it increases from lithium to Neon. ELEMENT EAI kJ mol 1 60 Be o 83 c 122 N o o 141 333
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    2. In the group I Alkali metals, the values decrease which is given in the following table ELEMENT EAI ICJ mol 1 73 60 Na 53 48 46.9 45.5 Electron affinity value is the measure of tightness of extra electron to the atom. REASONS: Along a period, there is increase in the nuclear charge and decrease in atomic size. Therefore the electron is more tightly bound and greater is the energy released EAI values of Be N and Ne are very low because of high symmetric electron configurations and high ionisation energies. Hence EAI value is zero. In a group EA decrease with atomic number. This is because the added electron is far away from nucleus hence loosely bound. Fluorine has lower electron affinity than chlorine because of smaller size. (Figure No. 1.5) 350 325 IEI( Kj moh) 300 275 10 20 30 40 50 60 FIG. ELECTRON AFFINITY IN HALOGENS
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    Higher values of EA indicate greater tendency to gain electrons and hence good oxidizing capacity. LECTRONEGATIVITY It is defined as " The relative electron attracting tendency of an atom for the shared pair of electron in a covalent bond It has no specific units. Greater the value of electronegativity , higher is its tendency to attract the shared pair of electrons towards itself. It is a relative value. Fluorine ( 4 0 ) is highest electronegative element. Where as Cesium ( 0.7 ) is least electronegative element. ARIATION OF ELECTRONEGATIVITY IN PERIODIC TABLE Electronegativity value increases across the period and it decreases down the group. XAMPLE 1. In the second period, the value of electronegativity increases from Lithium to Fluorine. ELEMENT ELECTRONEGATIVITY ANALYSE: Li 1.0 Be 1.5 2.0 c 2.5 N 3.0 o 3.5 4.0 Describe the Group I elements (the alkali metals) and predict their properties using the Periodic Table.
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    1.3 ROUP 1 ELEMENTS - ALKALI METALS 1. Group I elements are called alkali metals because they react with water to form alkalis. 2. The diagram below shows the properties of Group I elements in the Periodic Table. Na 000000 ALKALI METALS IN THE PERIODIC TABLE 3. Physical properties of alkali metals: a) The physical properties of Group I is metal charge gradually down the group. b) The density of the alkali metal generally increases down the group. Lithium, sodium and potassium are less dense than water. Rubidium and caseium are denser than water. c) They are relatively soft and can be cut easily with a knife. d) They have low melting and boiling points compared with other metals. The melting boiling points decreased down the group. e) They are good conductors of heat and electricity.
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    f) The following table shows the melting and boiling points, density and appearance of same Group I elements. ALKALIS METAL LITHIUM, Ll SODIUM, NA POTASSIUM, K RUBIDIUM, RB MELTING POINT (oc) 180 98 64 38 BOILING POINT (oc) 1330 890 688 688 DENSITY (gcm3) 0.53 0.97 0.86 1.53 APPREANCE Shiny, silver solid Shiny, silver solid Shiny, silver solid Shiny, silver solid 4. hemical ro erties of alkali metal a) They have similar chemical properties because all the elements have similar electronic configuration , that is, one electron in the outermost shell. By losing their valence electron, an alkali metal attains a noble gas configuration. b) The Group I metals are very reactive. They react rapidly with air and vigorously with water.
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    ALKALI METÄi) LITHIUM SODIUM POTASSIUM They are kept under oil to prevent coming into contact with air and water. eaction of alkali metals with water 5. a) All the alkali metals reacts with cold water to form alkali (soluble metal hydroxides) and hydrogen gas. + 2H20(l) NaOH (aq) OBSERVATIO • Reacts quickly. • Lithium floats on the water • No flame is seen. • Reacts very quickly. • Sodium melts and burns with a yellow flame. The molten sodium darts around the water surface. • Reacts violently. • Potassium melts and burns a lilac flame. 2Li(s) + 2H20(l) water QUATION N2Li (aq) + H2(g) lithium hydrogen Lithium 2Na (s) Sodium 2K (s) + Hydroxide sodium Hydroxide water 2H20 (l) hydrogen Potassium water —+2KOH(aq) + H2 (g) potassium hydrogen Hydroxide c) The reaction between alkali metals and water is a redox reaction. In the reaction between alkali metals and water is a redox reaction. Na (s) -.Na+(aq) + e- .. oxidation Sodium acts as the reducing agent when it reacts with water .
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    21--120 (1) + H2 (g) + 20H-(aq) . Overall reaction: .. reduction 2Na (s) + 2Na+ (aq) + 201-1- (aq) + H2(g) 6. The alkali metals are powerful reducing agent because they give up their valence electrons easily. 7. The reactivity of Group I elements increases down the group because the atomic size increases. The larger the atomic size, the further the valence electron is from the nucleus, and thus, the easier for the atom to lose the valence electron. Reactivity increases Density increase Melting and Boiling points decrrase Reducing power increases Most Reactive
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    XAMPLEI How and under what conditions do you expect rubidium to react with water? OLUTIO Rubidium is below potassium in Group I of the Periodic Table. It reacts violently with water at room temperature to give rubidium hydroxide and hydrogen. 2Rb(s) + 2RbOH(aq) + H2(g) TEST YOURSELF: Describe the group VII elements ( the halogens ) and predict their properties using the Period Table GROUP Vll ELEMENTS - HALOGEN 1. 2. 3. The Group VII elements are also called halogen in the Periodic Table. The diagram on the right shows the position of halogens in the Periodic Table. The halogens are non-metals and exit as diatomic covalent molecules, F2, C12, Br2 and 12. Their ions, such as fluoride ion bromide ion (Br-) and iodide ion
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    4. OOOOOOOOOOOOOOOOOO OOOOOOOOOOOOOOOOOO oooooooooooooooooo HALOGENS IN THE PERIODIC TABLÉ h sical ro erties of halo ens a) All halogens have low melting and boiling points. b) All the halogens are coloured and are non-conductors of electricity. c) On going down the group, i ) The colour intensity increases , ii) The melting and boiling points increase. d) The following table shows the physical properties of some halogens. HALOGEN FLUORINE, F2 COLOURPHYSICAL STATE Pale yellow gas MELTING POINT oc -219 BOILING POINT(OC) -188
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    CHLORINE, C12 BROMINE, Br2 Iodine, 12 Greenish-yellow gas Reddish-brown liquid Purplish-black solid -101 114 -35 59 184 5. hemical ro erties of Halo en . a) Halogens have similar chemical properties because they similar electronic structure, that is seven electrons in the outermost shell . b) Halogens react vigorously with metals to form salts called halides. The metal halides are ionic compounds. Iron Chlorine Iron (Ill) Chloride (ionic compound ) c) They also react with non-metals to form covalent compounds. XAMPLE Hydrogen chlorine hydrogen chloride (covalent compound )
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    d) Halogens are powerful oxidising agents because they gain electrons readily to form halide ions. XAMPLE C12(aq) + Chlorine 2C1-(aq) Chloride 6. Unlike Group I elements, the reactivity of halogens decreases on going down the group. 7. Displacement reactions of halogens: a) A displacement reaction is a reaction in which one element takes the place of another element. b) A more reaction halogen will displace a less reactive halogen from its halide solution as shown in the following table. C12 Br 12 As we go down the halogens group , Reactivity Decreases Melting and Boiling points increase Colour intensity increases HALOGEN CHLORINE KC1(aq) Least Reactive KBr(aq) Reddish-brown solution formed. KI (aq) Brown solution formed.
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    C12(aq) + 2Br-(aq) C12(aq) + 21- (aq) 2C1- (aq) + 12 (aq) BROMINE IODINE NO REACTION NO REACTION Tips for Students: Brown solution formed. Br2(aq) + 2 1-(aq) 2Br-(aq) + 12(aq) No Reaction 1. When writing the chemical equations for the displacement reactions of Halogens, the state symbols for the reactants are all (aq).this is because the Halogens and the halides are dissolved in water. 2. halogens with halides (ions of halogens) are two different identities. For example: Chlorine is C12 and chloride is CI-. Similarly,bromine is Br2 and bromide is Br- . 8. A displacement reaction is a Redox Reaction. In the displacement reaction between chlorine and potassium bromide, chlorine oxidises bromide ions to bromine and is itself reduced to chloride ions . Reduced (aq) + 2Br- (aq) 2C1- (aq) + Br2 (aq) Oxidised
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    1.5 COMMON ERRORS Potassium fluoride ( x potassium bromide ( KBr ) and potassium iodide ( KI ) are coloured substances. ACTUAL FACTS The halogens, F2, C12, Br and 12, are coloured substances, but KF, KBr and KI are white solids and their aqueous solutions are colourless. ROUP ELEMENTS -NOBLE GASES: 1. The Group O or Group VIII elements are called Noble Gases. 2. The diagram below shows the positions of Group O elements in the Periodic Table. oooooooooooooooooo NOBLE GASES IN THE PERIODIC TABLE 3. PHYSICAL PROPERTIES OF NOBLE GASES: They exist as single atom and are known as Monatomic Elements. a) They are Colourless gases at room temperature and pressure. b) He Ar
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    c) They have Low Melting and Boling Points that increase on going down the group. d) They are insoluble in Water. 4. CHEMICAL PROPERTIES OF NOBLE GASES: a) The noble gases are Inert (Chemically Unreactive) because they have stable electronic configurations . b) Their outermost shells are fully occupied, with two valence electrons for helium atom and eight valence electrons for the other noble gases. 5. USES OF NOBLE GASES: NOBLE GAS USES HELIUM NEON ARGON For filling weather balloons, advertisement balloons and airships In advertising signs ( neon lights ) In light bulbs, in welding and making steel PROPERTY Low density, non - flammable Glows brightly when electricity is passed through Provides and inert atmosphere to prevent the filament in light in light bulb or the steel in the furnace from reaction with oxygen
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    1. The transition elements (or transition metals) are located between Groups Il and Ill of the Periodic Table . Iron, Copper, Nickel, Silver and Gold are some examples of Transition Elements . 1 11 111 IV V VI Vll TRANSITION ELEMENTS oooooooooooooa ooooooooooooooon Y Zr Nb Mo Tc Ru Rh Pd Ag Cd TRANSITION METALS IN THE PERIOD TABLE 2. Melting Point and Density: All transition elements are metals. Unlike the alkali metals, transition elements have high melting points and high density . The table shows the melting points and density of three transition elements, iron, nickel and copper, as compared to that sodium.
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    ELEMENT IRON (Fe ) NICKEL ( Ni ) COPPER ( cu ) SODIUM ( ) MELTING POINT 1535 1453 1083 98 DENSITY ( g/cm3 ) 7.86 8.90 8.96 0.97 3. Oxidation States: a) Group I metals have a fixed oxidation state of + 1 and group Il elements have a fixed oxidation state of + 2, but a transition element has variable oxidation states ( i.e., more than one oxidation state ) in its compounds . The oxidation states of the transition elements, chromium and manganese are shown below . COMPOUNDS OF TRANSITION ELEMENTS Chromium ( Ill ) Oxide Potassium Dichromate ( VI ) Manganese ( Il ) Chloride Manganese (IV ) Oxide CHEMICAL FORMULAE Cr203 K2Cr207 MnC12 Mn02 OXIDATION STATE OF THE TRANSITION ELEMENT
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    Potassium Manganate (VII ) KMn04 4. b)Because of their variable oxidation states, transition elements can form more than one compound with another element . For EXAMPLE: Iron ( Il ) Chloride, FeC12 and Iron ( Ill ) Chloride,FeC13. Copper ( I ) Oxide, Cum and Copper ( Il ) Oxide, Cuo. Coloured Compounds: Transition elements form coloured compounds . For example, Iron(ll) sulfate is green, Copper ( Il ) sulphate is blue and potassium manganate (VII) is purple. The colour of the anhydrous compound of a transition element may be different from its hydrated ( Crystalline ) compound. For example, anhydrous copper ( Il ) sulfate, CuS04, is white, but the hydrated (Crystalline) copper (Il) sulfate, CuS04.5H20, is blue. Similarly, anhydrous cobalt (Il ) chloride is blue, but hydrated cobalt (Il) chloride is pink . Because of this property, both anhydrous copper ( Il ) sulfate and anhydrous cobalt ( Il ) chloride are used as a test for the presence of water .
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    COMMON ERRORS Transition elements are x coloured. Only transition metals form x coloured compounds. 5. As Catalysts: ACTUAL FACTS Transition elements are not coloured . It is their compounds that are coloured. Few non - transition metals also form coloured compounds besides transition metals. For example, led (Il) oxide( Pb0) is a yellow solid. a) A Catalyst is a substance that increases the speed of a reaction. b) Transition elements are often used as catalysts. For Example, platinum is used as a catalyst to speed up the reaction between nitrogen and hydrogen to form ammonia . (g) + 3H2(g) 2NH3 4500C, 250 atm c) Compounds of transition elements are also used as catalysts to speed up chemical reactions. For Example, copper ( Il ) sulphate is used to speed up the reaction between zinc and dilute sulfuric acid to form zinc sulfate and hydrogen gas.
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    CuS04 (Catalyst) Zn (s) + H2S04 (aq) ZnS04 (aq) + H2 (g) d) Transition elements and their compounds are often used as catalysts in industrial processes. TIP FOR STUDENTS: Compounds of some non-transition metals can also act as catalysts. STRUCTURAL QUESTIONS AND ANSWERS 1.) In A modern Periodic Table: a) What is the link between an element's electronic configuration and the Group it is in? b) Describe how the elements are arranged in the modern Periodic Table in terms of atomic number. Answer: a.) The group number of an element (for groups I to V Il) corresponds to the number of electrons in its outermost shell. b.) The elements are arranged in the order of increasing atomic number. 2) (You may refer to the Periodic Table to answer this question). Give the chemical symbol of: a) b) c) d) e) f) A metal used as the catalyst in the Haber process: A metal which is a liquid at room temperature and pressure; The most reactive non-metal; An element which consists of diatomic molecules and constitutes 78% of air An element in Period 2 which forms an ion of the type X + An element in Period 2 which has four electrons available for bonding.
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    Answer: 3) The properties of some elements are described below. For each of the elements, state which position of the Periodic Table, A, B, C, D, E or F, it is found in a) This element exists as a colourless gas at room temperature. It is a poor heat conductor. It is used as a gas to fill up glass light bulbs in incandescent lighting because it does not react with their filaments. Position : b) This element has a melting point of 2334 0 C. it is a versatile catalyst. It has oxidation states of +2, +3 and +4 in its compounds. Position : c) This element exists as a silvery-white solid at room temperature. It is strong and light and is used in the manufacture of electronic components. It is highly flammable when shaved into thin strips and produces a brilliant white flame when burnt in air. It has oxidation state of +2 in its compounds. Position : d) This element is used in the galvanization of steel to prevent corrosion. It burns in air to give an oxide which reacts with both acids and alkalis. It has an oxidation state of +2 in its compounds. Position : e) This element is used in the galvanization of steel to prevent corrosion. It burns in air to give an oxide which reacts with both acids and alkalis. It has an oxidation state of +2 in its compounds. Position : f) This element has a melting point of 2210 C. it is used widely in the manufacture of glass and ceramics. It reacts with oxygen to give a colourless, flammable gas. Position :
  • 38
    g) This element exists as a very soft, silvery-white solid at room temperature. It is a component of many low-melting alloys. It has an oxidation state of +3 in many of its compounds. Position : h) This element exists as a very soft, silvery-white solid at room temperature. It is a component of many low-melting alloys. It has an oxidation state of +3 in many of its compounds. Position : Answer: i) Which element in Period 3 has the highest melting point? Explain in terms of structure and bonding, why this element has a high melting point. ii) Which element in Period 3 has the lowest melting point? Explain, in terms of structure and bonding. Why this element has a low melting point. b) Name the Period 3 elements which are good conductors of electricity. c) An element X, in Period 3, reacts with oxygen to give an ionic compound with the formula XO. Identify the element X. a) An element Y, in Period 3, burns in air on gentle heating to produce a colourless gas with the formula Y02. Identify the element Y. Answer: i.)Silicon has the highest melting point. Silicon has a giant molecular structure. A large amount of energy is needed to overcome the strong covalent bonds holding the Si atoms together. Hence, silicon has a high melting point. ii. Argon has the lowest melting point. Argon has a simple atomic structure. A small amount of energy is needed to overcome the weak forces discrete argon atoms being together. Hence, argon has a low melting point.
  • 39
    b) Sodium, Magnesium and aluminium. c) Magnesium d) Sulfur. 5. The table below shows the formulae and melting point of oxides formed from six elements in Period 3. Formula of element Si s s Group 1 11 111 Formula of oxide Melting point PC Na2 0 A1203 Si02 P406 P4010 1280 2900 2040 1610 24 Sublimes at 300 oc -75 17 a) Explain, in terms of structure and bonding, the variation in melting points of the oxides listed in the table. b) Explain why when some magnesium oxide powder is mixed with water at room temperature, a solution of around pH 9 is obtained. c) Write balanced chemical equations for the reactions of water with the oxides of sulphur.Suggest the pH of the solutions formed. d) One of the oxides listed in the table does not dissolve in water but reacts with both acids and bases. Identify this oxide. e) Write the formulae of the two oxides of selenium (Group VI). Answer: 5.
  • 40
    a) Na20, MgO and A1203 have giant ionic structures (lattices) in which the positively charged ions and the negatively charged oxide ions (02-) are held together by strong electrostatic attraction (ionic bonds). A large amount of energy is needed to overcome the strong ionic bonds holding the positive ions and 02- ions together. Hence, these ionic oxides have high melting points. Si02 has a giant molecular structure. A large amount of energy is needed to overcome the strong covalent bonds holding Si and O atoms together. Hence, Si02 has a high melting point. P406, P401(), S02 and S03 have simple molecular structure. A small amount of energy is needed to overcome the weak intermolecular forces holding the discrete molecules together. Hence, these covalent oxides with simple molecular structures have low melting points. b) Magnesium oxide reacts with water to form magnesiu hydroxide. Mgo + Magnesium hydroxide is only sparingly soluble in water. Only a small amount of hydroxide ions gets into the solution. Hence, the resulting solution is a weak alkaline (pH 9) c) S02+ H20 S03 + H20 H2S03 H2S04 In both cases, acidic solutions are produced. The pH of the solutions is around 2. d) Aluminium oxide. e) Se02, Se03 6.) Phosphorous (group V) reacts with oxygen to give two different oxides. (a) Fill the relevant formulae and oxidation states in the table below. (b) Explain why the phosphorous oxides do not conduct electricity when molten. (c) Arsenic (As) is in the same Group of the Periodic Table as phosphorous. Write the empirical formula of the two oxides of arsenic.
  • 41
    Answer: Molecular formula of oxide Em irical formula of oxide Oxidation state of phosphorous in oxide P406 P203 +3 P4010 P205 +5 b) The phosphorous oxides have simple molecular structures in which the molecules in which the molecules are held together by weak intermolecular forces. They do not conduct electricity because of absence of mobile charged particles (such as ions or electrons). c) AS203 and AS205 7.) Some physical and chemical properties of metallic element M are listed below. Colour Silvery white Density/ gcm- Melting 3 0.89 oint/0C 63 Boiling oint/0C 759 Reaction with water Bubbles vigorously and catches fire to form aqueous MOH (a) (i) Deduce the Group of the Periodic Table to which element M belongs. (ii) Write a balanced chemical equation,with state symbols , for the reaction between element M and water. (b)M reacts vigorously with fluorine to form a salt MF.When 0.78 g of M is reacted with fluorine, 1.16g of MF is obtained . (i) Write a balanced chemical equation , with state symbols , for the t reaction of element M and fluorine. (ii) Calculate the number of moles of fluoride ions in the salf formed , and hence determine the relative atomic mass of M.
  • 42
    Answer: 7.) a) i) M belongs to Group I ii) 2M (s) + 2H20(1) b) i) 2M (s) + F2 (g) 2MOH (aq) + f-12 (g) 2MF (s) ii) Mass iof fluoruide ions = 1.16 — 0.78 = 0.38 No. of moles of fluoride ions = 0.38 / 19 = 0.02 No. of moles of M + ions = 0.02 Relative atomic mass of M = 0.78 / 0.02 = 39 8.)Sodium is an element in Group I of the Periodic Table.the atomic number of sodium is 11. (a) Describe what will be observed when a small piece of sodium is put into a bowl of cold water containing a small amount of Universal Indicator. (b) Liquid sodium reacts with hydrogen gas to give an ionic solid , sodium hydride(NaH) (i) Explain why this is a redox reaction. (ii) Draw a ' dot and cross' diagram to show the electronic structure of sodium hydride . You Only need to show the valence electrons. (iii) Explain in the terms of structures and bonding, why sodium hydride has a high melting point. (c) Rubidium is another element in Group l. the atomic number of rubidium is 37. (i) Use your knowledge of Group I elements to complete the table below. S mbol of rubidium Formula of rubidium chloride Formula of rubidium h dride Formula of rubidium sulfide Formula of rubidium sulfate Formula of product of the reaction between rubidium and water
  • 43
    (ii) Answer: Explain why the reaction between rubidium and water is more vigorous than the reaction between sodium and water. a) The piece of sodium rapidly melts into a silvery-white ball, which runs on the surface of the water making a hissing sound, and soon catches fire. The Universal Indicator turns from green to blue. b) i) In this reaction, Na (oxidation state 0) is oxidized to Na+ (oxidation state +1) while H2 (oxidation state o) is reduced to H- (oxidation state -1). Hence, this is a redox reaction. iii) Sodium hydride has a giant ionic structure (lattice) in which the positively- charged sodium ions (Na+) and the negatively-charged hydride ions (H-) are held together by strong electrostatic attraction (ionic bonds). A large amount of energy is needed to overcome the strong ionic bonds holding the positive ions and negative ions together. Hence, sodium hydride has a high melting point. Formula of rubidium chloride Formula of rubidium h dride Formula of rubidium sulfide Formula of rubidium sulfate Formula of products of the reaction between rubidium and water RbC1 RbH Rb2S Rb2S04 RbOH and H2 ii)Rubidium is in period 5 while sodium is in period. Rubidium has a larger atomic radius than sodium. The outermiost electron of rubidium is further away from the nucleus as compared to eth outermost electron of sodium. The attraction between the nucleus and the outermost electron in rubidium is not as strong as that in sodium. Hence, this electron can be removed more easily. Thus, rubidium is more reactive than sodium and the reaction between rubidium and water is more vigorous. 8.) Cesium(Cs) is an element in Group I of the periodic Table. The atomic numbers of caesium is 55. (a) Explain why cesium must be kept under dry argon.
  • 44
    (b) Explain why cesium reacts with ice at - 1 IOOC but lithium does not. Answer: a) Cesium is a very reactive metal which reacts vigorously with oxygen and water in air. Thus, it should be stored under dry argon (an inert gas). b) Cesium is in period 6 while lithium is in period 2.Cesium has a larger atomic radium than lithium. The outermost electron of cesium is further away from the nucleus as compared to the outermost electron in cesium is not as strong as that in lithium. Hence, this electron can be removed more easily. Thus,cesium is more reactive than lithium. 9.) The elements in group V Il of the periodic Table are also known as halogens. (a) Com lete the table below. Atomic numbers 9 17 35 53 Elements Molecular formula State at room temperature and ressure gas Li uid Colour Pale Yellow Black (b) The halogen react with hydrogen to form hydrogen halides. (i) (ii) (iii) Explain why all the hydrogen halides have low melting and boiling points. Pure hydrogen halides are poor conductors of electricity but hydrogen halides diluted with water are good electrical conductors. Explain this phenomenon. The table below shows the conditions required for the hydrogen and some halo ens to take lace. Halo en Fluorine Conditions and rate of reaction Explosive under all conditions; ra id reaction
  • 45
    Chlorine Bromine Explodes in the presence of li ht, ra id reaction Reacts when heated in the presence of appropriate catalyst; slow reactoion The reactivity of the halogen show a trend down the group. Describe the trend. (c) Astatine is another group V Il elements, found below iodine in the periodic Table. (i) (ii) (iii) How many electrons are there in the outermost shell of astatine? Suggest the appearance of astatine at room condition. Write the formula of hydrogen astatide. Atomic number 9 17 35 53 b) Answer: Element 1 Molecular formula C12 Br2 12 State at room temperature and ressure Gas Gas Li uid Solid Colour Pale ellow pale yellow- reen Red Black i)The hydrogen halides have simple molecular structures. A small amount of energy is needed to overcome the weak intermolecular forces holding the discrete molecules together. Hence, the hydrogen halides have low melting points. ii)The hydrogen halides have simple molecular structures in which the molecules are held together by weak intermolecular forces. They do not conduct electricity because of the absence of mobile charged particles (such as ions or electrons).
  • 46
    However, when diluted with water, the hydrogen halides react with water are able to conduct electricity because of the presence of mobile ions. iii)The reactivity of halogens decreases down the group. ii) Black solid iii) HAt 10.) Chlorine, bromine and iodine are halogens. Given the following chemicals, describe how will you carry out a series of experiments to show the relative chemical reactivity of the three halogens. Chemicals provided; aqueous chlorine aqueous bromine aqueous iodine aqueous potassium chloride aqueous potassium bromide aqueous potassium iodide your answer should include; A table with the reagents which you will use and the observation you will see in the each experiment ionic equation for any reaction which takes place Answer: 10.)Aqueous chlorine is mixed sequentially with aqueous potassium chloride, aqueous potassium bromide and aqueous potassium iodide. This experiment is repeated with aqueous bromine and aqueous iodine. The observations are shown below. Aqueous chlorine Aqueous bromine Aqueous iodine KCI (a ) No visible change; solution remains ale reen No visible change; solution remains oran e No visible change; solution remains dark brown KbR (a ) Solution turns orange due to the formation of Br2 No visible change; solution remains oran e No visible change; solution remains dark brown Solution turns dars brown; 12 formed. Solution turns dark brown; 12 formed No visible change; solution remains dark brown
  • 47
    The experiments above show that: Chlorine displaces bromine from potassium bromide. C12 + 2 Br- Br2 + 2 Chlorine displaces iodine from potassium iodide. C12 + 21- 12+ 2 Cl- Bromine displaces iodine from potassium iodide. Br2 + 21- 12 + 2Br- Hence, chlorine is the most reactive while iodine is the least reactive among the three halogens. 11.) For each of the followings reactions involving Group V Il elements, predict whether or not a reaction will take place ,write a balanced chemicals equation for a the reaction. (a) Lithium is heated with chlorine gas. (b) Sodium is heated with astatine. (c) Aqueous bromine is added to aqueous potassium bromide. (d) Aqueous iodine is added to aqueous potassium bromide. (e) Chlorine is bubbled through aqueous potassium iodide. Answer: a) Yes, a reaction takes place. b) Yes, a reaction takes place. c) Yes, a reaction takes place. d) No reaction takes place. e) Yes, a reaction takes place. 2Li + C12 2Na + At2 Br2 + 2KAt C12 + 2K1 2 LiC1 2NaAt At2 + 2KBr 12 + 2KC1
  • 48
    12.) Argon is the most abundant noble gas on Earth. (a) element chlorine exist as diatomic molecules while argon is a monoatomic gas. give reason (b) Argon is often used in light bulbs. Give one property of argon that makes it suitable for this application. (c) Describe how pure argon is obtained commercially. Answer: a) Argon has eight electrons in its outermost shell. The electronic configuration of argon (2, 8, 8) is very stable. Thus, argon is unreactive and exists as single atoms. On the other hand, the electron configuration of chlorine is 2, 8, 7. To achieve the stable electronic configuration of argon, each chlorine atom bonds covalently with another chlorine atom, forming diatomic molecules. b) Argon is inert and will not react with the hot tungsten filament. c) Argon is obtained from the fractional distillation liquified air. 13.) The table below contains about the first three Grou 0 elements. Name Helium Neon Ar on Atomic number Relative atomic Density at r.t,p,/g 2 10 18 mass 4 20 40 -3 dm 0.167 0.833 1.67 a) Suggest why group elements are also known as noble gases. b) Describe how the boiling points of group 0 elements change down the group. c) Helium is used to fill small balloons. Explain what properties of helium make it suitable for this application. d) Krypton is another group 0 element, found below argon in the periodic table.
  • 49
    i. State the relative atomic mass of krypton ii.Calculate the density of krypton at r.t.p. Answer: a) Helium has two electrons in its outermost shell while the other Group 0 elements have eight electrons in their outermost shells. All the Group 0 elements have stable electronic configurations. Thus these elements are uncreative and seldom react with other substances. In addition, these elements have vey low melting and boiling point and exist as gases under room conditions. b) The boiling points of the elements increase down the group. c) helium has a very low density and is unreactive (not flammable) d) i) relative mass of crypton is 84 ii) Density of krypton at r.t.p = 84 / 24 = 3.5 g/dm3. 14.)The table below shows the physical properties of some transition metals. Metal Titanium Vanadium Chromium Man anese Iron Cobalt Nickel Co er Density gm- 4.50 5.96 7.20 7.20 7.86 8.90 8.90 8.92 Melting point/ Boiling point/ Relative 1660 1890 1244 1244 1535 1495 1455 1083 3287 3380 1962 1962 2750 2870 2730 2567 electrical conductivit 1 1.5 5 1 8.5 12 12 49 a) List two differences between the physical properties of Group I metals and the transition metals.
  • 50
    b) Identify the metal with the highest electrical conductivity. Explain why it is often used to make electrical wiring but is NOT usually used to make overhead power lines. c) Titanium is used to make many aircraft parts. Relate this application of titanium to its physical properties. d) Calculate the number of moles of nickel atoms that are found in a 10 cm3 block of nickel. Answer: a) The densities of Group I metals are low while the densities of transition metals are high. The melting points of Group I metals are also lower than the melting points of the transition metals. b) Copper. Density of Copper is too high so it is usually not used to make overhead power lines. c) Titanium has a low density (compared to the other transition metals) and is strong. d) Mass of a 10 cm3 block of nickel = 10 x 8.90 = 89 g No. of mol of nickel = 89 / 59 = 1.51 (3 s.f.) 15.) Rubidium (Rb) and ruthenium (Ru) are metals. Some properties of rubidium and ruthenium and their compounds are given in the table below. Metal -3 Density /kg m 1532 12370 Melting point / oc Compounds 39 2334 formed (colour of com ound) RdF (white) RbCl (white) RuF3 (brown, RuF4 (yellow), RuF6 (brown), RuC12 (brown), RuC13 (brown)
  • 51
    Based on the data above, suggest where the metals could be found in the periodic table. Support your answer by making reference to the physical and chemical properties of the metals. Answer: Rb can be found in Group I. it has low melting point and low density. The oxidation state of Rb is its compounds is +1 ) (each atom of Rb gives away one electron). Ru is a transition metal. It has high melting point and high density.lt forms compounds of variable oxidation states and these compounds are coloured. 16.) The Contact process is an individual process for producing sulfuric acid. During this process, sulfur dioxide is converted to sulfur trioxide. This is done by reacting sulfur dioxide gas and oxygen gas in the presence of vanadium (V) oxide. In the absence of vanadium (V) oxide, the conversion of sulfur dioxide and sulfur trioxide is very slow. The chemical equation for this reaction is 2S02+ 02 2 S03 In the presence o vanadium (V) oxide, this conversion takes place faster and in two steps. Step 1: Step 2: S02 + V205 4 V 02 + 02 S03 + 2C02 2V205 a) What is the role of vanadium (v) oxide in the contact process? b) Based on the information given above, explain why only a small amount of vanadium (V) oxide is needed in the process. c) Identify the oxidizing and reducing agents in steps 1 and 2. Answer: a) It acts as a catalyst. b) V205 used up in step 1 is regenerated in step 2. c) Step 1: Oxidising agent is V205 and reducing agent is S02 Step 2: Oxidising agent is 02 and reducing agent is V02
  • 52
    17.) Antonie Lavoisier, a French scientist is often cited as the father of modern chemistry. In 1789, he published a book Elementary Treatise of Chemistry — and in this text, he explained that an element is a substance that could not be broken down by any known method of chemical analysis. The table of elements proposed by Lavoisier is show below. Based on the experiments he had conducted and the above observations made, he classified the substances into four element groups. Simple gas-like substances Light Heat Oxygen Nitrogen Hydrogen Non-metallic elements Sulfur Phosphorous Carbon Metallic elements Simple earthy Antimony Silver Arsnic Bismuth Cobalt Copper Tin Iron Manganese Mercurys Molybdenum Nickel Gold Platinum Lead Tungsten Zinc elements Lime (calcium oxide) Magnesia (magnesium oxide) Barite (barium sulfate) Alumine (Aluminium oxide) Silice (silicon dioxide) a) Your knowledge of the modern periodic table is that all the metallic elements listed by Lavoiusier metals? b) i) From your knowledge of the modern periodic table, which one of the four element groups proposed by Lavousier are definitely not elements but compounds? ii) Suggest why the scientists of the 18th century thought that these substances were elements.
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    iii) Identify two terms in the list which do not refer to an element or a compound. Answer: Yes b) i) The simple earthy elements. ii)The scientists did not have the instruments to separate the elements in these compounds. For instance, electrolysis could not be done to extract the reactive meals from the salts. iii) Light and heat. 18.) 0.78 g of a Group I element M reacts with oxygen to give 0.94 g of an oxide M20 a) Calculate the relative atomic mass of M and identify M. b) 0.94 g of M20 reacts further with more oxygen gas on heating to give a peroxide with the formula M202. Write a balanced chemical equation with state symbols, for the reaction between M20 and oxygen. c) i) Write the formula of the anion present in M202 and state the oxidation state of oxygen in this anion. ii) Explain why this anion is expected to behave as an oxidizing agent. Answer: 18.) a) 4M + 02 Mass of 02 reacted = 0.94 - No. of moles of 02 reacted — No. of moles of M reacted Ar of M = 0.78 / 0.02 = 39 2 M20 0.78 = 0.16g - 0.16 / 32 = 0.005 = 4 x 0.005 = 0.02 Thus, M is Potassium. b) 2M20 (s) + 02 2M202 (s)
  • 54
    c) i) 022-, the oxidation state of O is -1 ii) 022- is expected to behave as an oxidizing agent because it can be easily reduced to the stable oxide ion Of. From 022- to 02- the oxidation state of O is changed from -1 to -2. 19.) When chlorine is bubbled into a solution of iodine in hot aqueous sodium hydroxide, the two halogens react, forming a white precipitate W, sodium chloride and water. W has the following composition by mass: 1 H O 16.9 % 46.7 % 1.1% 35.3% a) Calculate the empirical formula of W. b) What s the oxidation state of iodine in W? c) i) Given that 7 moles of chlorine reacts with 1 mole of iodine in the given reaction, construct a balanced chemical equation for this reaction. ii) Explain why this is a redox reaction. Answer: 19.)( a) % b mass No. of moles Simplest mole ratio 16.9 23 0.735 2 1 46.7 127 0.368 1 1.1 1 1.1 3 o 35.3 16 2.206 6 Hence, the empirical formula is Na21H306. b) Let the oxidation state of I in Na21H306 be x. 2(+1) + X + 3(+1) + 6(-2) = O
  • 55
    Hence, the oxidation state of I = x = +7 c) i) 7C12 +12 + 18NaOH Na21H306 + 14 NaC1 + 6H20 ii) In this reaction, iodine is oxidized. The oxidation state of iodine changes from 0 in 12 to +7 in Na21H306. On the other hand, chlorine is reduced. The oxidation state of chlorine changes from 0 in C12 to -1 in NaCl. Hence,this is a redox reaction. SUMMARY AND KEY POINTS 1.) A vertical column of elements in the Periodic Table is called a Group. 2.) A horizontal row of elements in the Periodic Table is called a Period. 3.) An element that has the properties of both a metal and a non-metal is called a Metalloid. 4.) The Group I elements in the Periodic Table is called Alkali Metals. 5.) The Group VII elements in the Periodic Table is called Halogens. 6.) The Group O or VIII elements in the Periodic Table is called Noble Gases. 7.) A substance that increases (enhances) the speed of a chemical reaction is called a Catalyst. 8.) Periodic table is defined as " The arrangement of elements in certain groups and periods in such a way that elements with similar properties are grouped together."
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    9.) The chemical properties of the elements are the periodic function of their atomic numbers. 10.)The Periodic Table consists of Seven Periods and Eight Groups of elements. 11.) The block of element between Group Il and Group Ill is called the transition elements or transition metals. 13.) Elements in the same group have the same number of valence electrons. This number is the same as the group number. Elements in the same group have similar chemical properties. 14.) The electronic configuration of an element can be worked out if we know the position of the element in the Periodic Table. 15.) The group number of an element tells us the number of valence electrons. The period in which the element is placed tells us the number of shells of electrons in the electronic configuration. 16.) Periodicity of elements is the repeatition of similar properties of elements placed in a group according to atomic numbers. The periodic table given by Moseley is called "long form of periodic table." 17.) Elements on the left of the periodic Table are metals, while those on the right side are non- metals . 18.) From left to right across a period, there is a decrease in the metallic properties and an increase in non-metallic properties . 19.) A metalloid is an element with properties that are intermediate between those of metals and non-metals. Examples are silicon, Si(Period 3), and germanium,Ge (Period 4).
  • 57
    20.) If the proton number or electronic configuration of an element is given, we can predict the metallic or non-metallic character of the element of the element and hence, predict the properties of the elements. 21.) In general, i) metals form basic oxides, ii) non-metals form acidic oxides and iii) metalloids form amphoteric oxides . 22.) In the ionic compounds, the metals lose electrons to form positive ions. The charges on the ions of the elements in Group I, Il, and Ill correspond to the group numbers . 23.) The elements in Group V, VI, and VII are non-metals. They tend to form covalent compounds when they combine with non-metals . 24.) Periodic properties of elements are those properties which are directly or indirectly related to the electronic configuration of the elements and show gradation in moving down the group or_along the period. 25.) Atomic radius is defined as "the distance between the centre of its nucleus and electron in the last orbit." There are 3 forms of atomic radius based on the nature of the bonding. They are i) COVALENT RADIUS ii) VANDER WAAL'S RADIUS iii) METALLIC RADIUS 26.) The atomic radii (covalent radii) of elements generally decrease from left to right in a period, while when we move down the group the atomic radii increases. 27.) Ionic radius is defined as " the distance between the centre of the nucleus to
  • 58
    the region up to which it can attract the electron cloud " 28.) The radius of a cation is always smaller than the neutral atom from which it is formed. This is because when an electron is removed from the outermost shell, the neutral atom has become unipositive ion. It means the nucleus gains a proton. Now the nucleus can attract the remaining electrons more strongly, hence the electron cloud shrinks, resulting in decrease in ionic radius. 29.) Isoelectronic Ions are the species of different ions which have same number of electrons but different magnitude of nuclear charges. The size of isoelectronic ions decrease with increase in atomic number. 30.) Inert gases are elements present in the zero group. They are called noble gases because they hardly react with other elements. 31.) The elements with atomic numbers 58 (Ce) to 71(Lu) are called lanthanides. they are also called Rare earth elements. 32.) The elements with atomic numbers 90(Th) to 103(Lr) are called actinides. they are radioactive elements. 33.) Lanthanides and Actinides series of elements have similar properties and have been placed in two rows at the bottom of the periodic table. They are collectively known as Inner- transition elements. 34.) Metals have a tendency to form a cation, thus reacting easily with an anion. 35.) Non-metals have a tendency to form an anion, thus reacting easily with a cation. 36.) Shortest period consisting of 2 elements H and He is period 1. 37.) Elements of the same period have the same number of electron shells.
  • 59
    38.) Ionisation Potential is the amount of energy required to remove an electron from the outermost shell of an atom and to make it a positive charge ion. Ionisation potential increases from left to right across a period. It decreases from top to bottom going down a group. 39.) Electron affinity is the amount of energy released when an electron is gained by a neutral atom. Electron affinity increases from left to right across a period. It decreases from top to bottom going down a group. 40.) Electronegativity is the tendency of an atom of an element in a molecule to attract the shared pairs of electrons towards itself. Electronegativity increases from left to right across a period. It decreases from top to bottom going down a group.

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