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The Periodic Table Table of Contents 1.0 The Periodic Table... 3 1.1 History... 3 1.2 Description of the Modern Periodic Table... 3 1.3 Classification of Elements... 4 1.4 Atomic Number and Mass Number... 9 2.0 Some Periodic Properties of Elements... 10 2.1 The Sizes of Atoms... 10 2.2 Ionic Radius... 12 2.3 Ionization Energy... 12 2.4 Electron Affinity... 13 2.5 Metallic Character... 14 3.0 Quiz... 15 The Periodic Table 2
1.0 The Periodic Table 1.1 History Dmitri Mendeleev and Lothar Meyer independently proposed in 1869 the periodic law. This law states that when the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically. Many scientists during the time of Mendeleev and Meyer have attempted to come up with a systematic arrangement of the elements. Mendeleev succeeded where others failed for two reasons. First, Mendeleev left blank spaces in his table for undiscovered elements. Second, he corrected previously erroneous atomic mass values. Two of the atomic mass values he corrected were those of indium and uranium. The blanks in the periodic table that Mendeleev created came at atomic masses 44, 68, 72, and 100. He predicted that the blanks are those of the elements that were soon to be discovered. Shortly after the appearance of his 1871 periodic table, two of the elements Mendeleev predicted were discovered. In Mendeleev s table, similar elements fall in the same vertical groups. The properties of these elements changed gradually from top to bottom in a group. 1.2 Description of the Modern Periodic Table The modern periodic table arranges the elements in 18 groups, called the long form. The vertical arrangement of elements is called a group or family. The horizontal row of elements is called a period. A group number is located at the top of the periodic table while the period number is on the left. The first two groups (groups 1 and 2) in the periodic table are known as the s-block. The last six groups (groups 13 to 18) belong to the p-block. Together, the s-block and the p-block constitute the main group elements. Elements in the inner lower box (groups 3 to 12) of the periodic table are known as the transition elements or the d-block. The f-block elements are the elements extracted from the periodic table. They are the two rows of elements below the main periodic table. The upper row is known as the lanthanides. They belong to period 7. Elements in the lower row are called the actinides. They belong to period 8. Together the actinides and the lanthanides are called the inner transition metals. The group and period numbers aid in the location and classification of elements. For example, the element in group 2 and period 4 is calcium. The element in group 17 and period 2 is the highly reactive halogen fluorine. 3 The Periodic Table
The Modern Periodic Table Group 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Period 1 1 H 2 He 2 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 3 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 4 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 5 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 6 55 Cs 56 Ba * 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 7 87 Fr 88 Ra ** 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 Uub 113 Uut 114 Uuq 115 Uup 116 Uuh 117 Uus 118 Uuo 8 119 Uue 120 Ubn Lanthanides 57 La 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu Actinides 89 Ac 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr Atomic numbers of solids are in black. Those of liquids are in blue, and those of gases are in red. 1.3 Classification of Elements The elements are classified as metals, nonmetals, and metalloids. The periodic table is divided by a zigzag line on the right portion. On top of this zigzag line are the elements boron (B), silicon (Si), arsenic (As) and tellurium (Te). Under the zigzag line, you will find aluminum (Al), germanium (Ge), antimony (Sb), and polonium (Po). Elements in olive green colored blocks are the metalloids. To the left of the zigzag line are the metals. To the right are the nonmetals. Most metals are good conductors of heat and electricity. They are also malleable and ductile. They have moderate to high melting points. In general, nonmetals are non-conductors of heat and electricity. They are usually The Periodic Table 4
brittle (solids). Many nonmetals are gases. The metalloids sometimes display the properties of metals, and sometimes, those of nonmetals. Through the color scheme of the periodic table above, we see that the majority of the elements are metals and that the nonmetals are confined to the right side of the periodic table. Most solids are metals. The only metal that is a liquid is mercury. The main group elements are also given special names. The table below summarizes the special names of the main group elements. The Main Group Elements and Their Special Names Group Special Names 1 Alkali Metals 2 Alkaline Earth Metals 13 Boron Family 14 Carbon Family 15 Nitrogen Family 16 Oxygen Family 17 The Halogens 18 The Noble Gases The group 1 elements or the alkali metals form cations with a charge of +1. They tend to lose 1 electron. These metals include lithium, sodium and potassium. The group 1 metals are silver colored. They are highly reactive, soft, and can be cut with a knife. Figure 1.1 The Alkali Metals. From left: lithium, sodium, potassium, rubidium, and cesium. The group 2 elements, also known as the alkaline earth metals, tend to lose 2 electrons to form cations with a charge of +2. They are also silver colored but are harder and denser than the group 1 metals. Their reactivity is also high but lower than those of the alkali metals. They have higher melting points. Some members of the family are calcium, barium, and strontium. Figure 1.2 The Alkaline Earth Metals. From left: beryllium, magnesium, calcium, strontium, and barium. 5 The Periodic Table
The group 13 elements are called the boron family. They tend to lose 3 electrons and form cations with a charge of +3. They are composed of the metalloid boron and the metals like aluminum, gallium and indium. Boron is a shiny, black substance in its pure form. It is very hard but much too brittle for practical use. Aluminum ranks third on the list of the ten most abundant elements in the earth s crust. It is ductile and a good conductor of heat and electricity. Gallium was one of the elements predicted by Mendeleev. Figure 1.3 The Boron Family. From left: boron, aluminum, gallium, indium, and thallium. The group 14 elements of the periodic table are composed of carbon, silicon, germanium, tin, lead and ununquadium. This group is also known as the carbon family. Among the members of the group, only silicon is the abundant substance on earth s crust. The carbon family has 4 valence electrons. They tend to have oxidation states of +4. Figure 1.4 The Carbon Family. From left: carbon, lead, silicon. C Pb Si Some members of the group 15 elements are nitrogen, phosphorus, arsenic, antimony and bismuth. They are also known as the nitrogen family. These elements have five valence electrons, 3 electrons short for a complete octet. Consequently, they tend to gain 3 electrons from other atoms and form anions with oxidation states of -3. These elements are capable of forming double and triple bonds with other elements. This is the reason for the notable stability of their compounds. All the elements in the group are solids at room temperature except for nitrogen, The Periodic Table 6
which is a gas. The most important member of the group 15 family is nitrogen, being the most abundant gas in the earth s atmosphere. Figure 1.5 The Nitrogen Family. From left: nitrogen, phosphorus, arsenic, antimony, and bismuth. The group 16 elements, also known as the oxygen family, are composed of oxygen, sulphur, selenium, tellurium, and polonium. Oxygen is a gas at room temperature. It is necessary for combustion and is abundant on earth. Sulfur is a brittle yellow solid. Large amounts of sulfur is used in the manufacture of the important acid, sulphuric acid. Selenium conducts electricity when exposed to light so it is used in the manufacture of photocopiers and solar cells. Polonium is very volatile and is radioactive. Tellurium is primarily used as a semiconductor. The group 16 elements have 6 valence electrons and are 2 electrons short for a complete octet. They tend to gain 2 electrons and form anions with an oxidation state of -2. They are capable of forming double bonds with other elements. Figure 1.6 The Oxygen Family. From left: oxygen, sulfur, selenium, and tellurium. The group 17 elements have 7 valence electrons. They are very willing to gain 1 more electron for a complete set of eight electrons in their valence shells. They are also known as the halogens. They are composed of the elements fluorine, chlorine, bromine, iodine, and astatine. At room temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. Group 16 is the only group in the periodic table that contains all the three states of matter. Figure 1.7 The Halogens. From left: fluorine, chlorine, bromine, and iodine. 7 The Periodic Table
The group 18 elements are also known as the noble gases. It is made up of the elements helium, neon, argon, krypton, xenon, and radon. These gases are highly unreactive. This is because they have a complete set of eight valence electrons and are stable. They do not easily react with other atoms because of their existing stability. Hence, they are also called inert gases due to their unreactivity. These gases are colorless and odorless. They have important applications in lighting, welding, and space exploration. Figure 1.8 The Noble Gases. From left: helium, neon, argon, krypton, and xenon. The transition metals are also known, in general, as the d-block elements. They are the elements in the lower inner box of the periodic table. Examples of transition metals are cobalt, chromium, nickel, manganese, and copper. The transition metals are known to form colored compounds. Examples of these colored compounds are shown in the picture below. Figure 1.9 Some Transition Metals. From left: aqueous solutions of: cobalt(ii) nitrate, Co(NO 3 ) 2 (red); potassium dichromate, K 2 Cr 2 O 7 (orange); potassium chromate, K 2 CrO 4 (yellow); nickel(ii) chloride, (purple). The inner transition metals are the elements in the two rows below the main periodic table. The first row elements are called the lanthanide series. They are the elements with the atomic numbers 58 to 71. The lanthanides form +3 ions as their principal species. Below are some examples of the lanthanide metals. The Periodic Table 8
Figure 1.10 The Lanthanide metals. From left: lanthanum, cerium, praseodymium, neodymium, samarium, europium, gadolinium, terbium, dysprosium, holmium, erbium, thulium, ytterbium, and lutetium. The lower row in the inner transition metals is also known as the actinide series. This row is composed of 14 elements that have atomic numbers 90 to 103. The atomic numbers of the actinides are even more uncertain than the lanthanides. Only minute amounts of some of the elements in this series are obtained because of their instability. This series is composed of actinium, thorium, protactinium, uranium, and others. 1.4 Atomic Number and Mass Number The elements are represented as where X is the symbol for the element, A is the mass number and Z is the atomic number. The atomic number is the number of protons in an atom. Since a neutral atom has equal number of protons and electrons, then the atomic number is also the number of electrons. It would only differ if the atom becomes an ion. The mass number is the number of protons and the number of neutron combined together. This is so because the mass of the electrons are negligible. The mass of the whole atom is only contributed by the protons and the neutrons. In many books, the positively charged proton is represented as p +. The negatively charged electron is e - and the neutron is simply n or n 0 because it has no charge. A = p + + n 0 Z = p + = e - From the above formula, the number of protons, neutrons, and electrons can simply be obtained. As an example, let us take a look at sodium. It is represented as. For sodium, A=23 and Z=11. It simply follows that p + =11, e - =11, and n 0 =12. Sodium has a mass number of 23 and an atomic number of 11. It contains 11 protons, 11 electrons, and 12 neutrons. How many protons, 9 The Periodic Table
electrons, and neutrons does a magnesium atom have? The symbol for magnesium is. Magnesium has an atomic number of 12 and a mass number of 24. A = 24 and Z =12 Since the number of protons is equal to the number of electrons, Z = p + = e - p + = 12 and e - = 12 The number of neutrons can be obtained from the given mass number. A = p + + n 0 n 0 = A p + n 0 = 24 12 = 12 So a magnesium atom has 12 protons, 12 electrons, and 12 neutrons. 2.0 Some Periodic Properties of Elements Elements in the same period have the same outermost energy level, also called the valence shell. Their valence energy levels correspond to their period number. So the valence shell of sodium in period 3 is the 3 rd energy level or the 3 rd shell. The elements in the s-block have the s valence subshells. It means that the valence electron of an element in the s-block is found in the s subshell of the atom. As an example, let us have sodium again. Sodium has 11 electrons. Its electron configuration is 1s 2, 2s 2, 2p 6, 3s 1. The outermost or the valence electron is located on the s subshell of the 3 rd energy level. The valence electrons of the p-block elements are located in their p subshells. The same is also true with the d-block and the f-block elements. The valence electrons are so important because it is they that give rise to the different properties of the atoms. The valence electrons are also the ones involved when an atom reacts with another atom or when they combine together to form a compound or molecule. 2.1 The Sizes of Atoms The size of an atom is taken to be the distance between the nuclei of two bonded atoms. As the number of electrons increases in an atom, more energy levels are needed to accommodate the electrons. From left to right of a period in the periodic table, the number of electrons increases. However, these elements have the same valence shells. This means that the added electrons are occupying the same energy level. As we go from left to right in a period, more electrons are occupying the same outermost energy level. This also means that more electrons are attracted to the nucleus. This causes the atom to shrink as the electrostatic force of attraction between the nucleus and the electrons get more intense. As we go from left to right in a period, therefore, the size of the atom decreases. The Periodic Table 10
Figure 2.1 Sodium has 1 valence electron while sulfur has 6. They both belong to period 3 of the periodic table. The electrostatic force of attraction is greater in sulfur. Sulfur is smaller that sodium. The above pictures are sodium and sulphur atoms. They both belong to period 3 of the periodic table. Their outermost energy level is the 3 rd energy level. Sodium has only one valence electron, while sulphur has six. There will be a greater electrostatic force of attraction between the valence electrons of sulphur and its nucleus, causing the atom to become smaller. As we go down a group in the periodic table, the number of electrons also increases. However, the added electrons occupy a new energy level. This causes the atom to become bigger as the size of an atom is determined by its energy levels. There are 11 electrons in a sodium atom while there are 19 in a potassium atom. The valence shell of sodium is the 3 rd energy level while it is the 4 th energy level in potassium. Clearly, although they both have one valence electron, potassium is bigger than sodium because it has four energy levels. Figure 2.2 Sodium has three energy levels. It belongs to period 3. Potassium has four. It is in period 4. 11 The Periodic Table
In general, atomic size increases from top to bottom in a group and decreases from left to right in a period. 2.2 Ionic Radius Ionic radius is the size of an ion. An ion is a charged atom. To form an ion, an atom loses or gains electrons. When electrons are gained, the resulting ion has a negative charge. The number of electrons gained is the number following its charge in its symbol. If the ion is S -2, then it is a sulphur atom that has gained two electrons. A negatively charged ion is called an anion. When an atom loses electrons, it gains a positive charge. The number of electrons lost is the number following its charge in its symbol. If the ion is Ca +2, then it is a calcium atom that has lost two electrons. A positively charged ion is called a cation. When an atom loses valence electrons to form a cation, there is an excess of nuclear charge over the number of electrons in the resulting cation. The nucleus draws the electrons in closer. This causes an overall contraction of the ion. Thus, cations are smaller than the atoms from which they are formed. For isoelectronic cations (cations with the same number of electrons and identical electron configuration), the more positive the charge, the smaller the ionic radius. When an atom gains electrons, the atom has to accommodate them by welcoming them in its valence shell or by expanding the size of the atom by the addition of a new energy level. Moreover, the added electrons experience a repulsive force from the electrons already in the atom. This causes the electrons to spread out more and thus cause the ionic size to expand. In general, anions are larger than the atoms from which they are formed. For isoelectronic anions (anions with the same number of electrons and identical electron configuration), the more negative the charge, the larger the ionic radius. 2.3 Ionization Energy Ionization energy, I, is defined as the energy needed to remove an electron from a gaseous atom or ion. It is an endothermic process as energy must be absorbed. Energy is needed to break the attractive force occurring between the valence electron and the nucleus. The energy needed to remove the first valence electron is called the first ionization energy, I 1. To remove the second valence electron, second ionization energy is needed. It would always be easier to remove an electron that is not tightly bound to the nucleus. This would require less amount of energy. This energy is usually expressed in kj (kilojoule) or in ev (electron volts). Ionization of an atom can be represented by the following. X (g) X - (g) I 1 = + kj/mol X - (g) X -2 (g) I 2 = + kj/mol The Periodic Table 12
As we go from left to right of a period in the periodic table, the atoms get smaller. This also means that the valence electrons become closer the nucleus and are tightly bound. To remove an electron that is tightly bound to the nucleus, it would require greater amount of energy. Therefore, ionization energy increases as we go from left to right of a period in the periodic table. As we go down from top to bottom of a group in the periodic table, atomic size increases. This also means that the valence electrons become farther away from the nucleus. These valence electrons are loosely held. It would then be less difficult to remove these electrons. Therefore, ionization energy decreases as we go down from top to bottom of a group in the periodic table. Ionization energy is inversely proportional to the atomic size. Figure 2.3 The first ionization energies of the first 88 elements.. 2.4 Electron Affinity Ionization energy concerns the loss of electrons. Electron affinity, EA, is a measure of the energy change that occurs when a gaseous atom gains an electron. When an atom gains an electron, energy is given off. The process is exothermic. Following the thermochemical convention, electron affinity is a negative quantity because it is energy released. The magnitude of the electron affinity speaks about how tightly an added electron is bound to the nucleus. If it is more negative, the added electron is tightly held by the nucleus. Electron affinity can be represented by the following. 13 The Periodic Table
X (g) + e - X - (g) AE = - kj/mol It is easier for an approaching electron to be attracted by the nucleus if the nucleus is nearer to it. That is, if the atom is smaller. This would give off greater amount of energy due to the successful binding between the approaching electron and the nucleus of an atom. The electron affinity would be more negative. On the other hand, an approaching electron would find it difficult to bind to a nucleus of a bigger atom because the nucleus is quite far from it. This would release less amount of energy. The electron affinity would be less negative. From this, it can be deduced that electron affinity increases (more negative) as we go from left to right of a period in the periodic table. It decreases (less negative) as the atoms get bigger or if we go down from top to bottom of a group in the periodic table. 2.5 Metallic Character Metallic property is a property of a metal that arises from its ability to lose electrons. This also means that the more easily an atom loses electrons, the more metallic it is. Loosely held electrons are the ones most easily lost by an atom. These are the electrons that are far from the nucleus. These electrons can be found in bigger sized atoms. Thus, bigger atoms that tend to lose electrons are more metallic as they tend to lose electrons more easily and form cations. As we go from left to right of a period in the periodic table, the atomic size decreases. The valence electrons are nearer the nucleus. The atoms do not tend to lose their electrons as these valence electrons are more tightly bound to the nucleus. Thus, their metallic character decreases. As we go down from top to bottom of a group in the periodic table, the atomic size increases. The valence electrons are farther from the nucleus. The atoms tend to lose their valence electrons more easily as these valence electrons are less tightly bound to the nucleus. Therefore, the metallic character of the atoms increases. The most metallic element in the periodic table is francium. The closer the element is to francium, the more metallic it is. In general, the metallic character of atoms increases down a group, and decreases across a period. The Periodic Table 14
3.0 Quiz I. Identify the Element Given the Period and Group Numbers Group (G)/Period (P) Element G2, P4 G16, P2 G18, P5 G11, P4 G12, P6 II. Fill in the Boxes Atom/Ion Z A p + e - n 0 Symbol U +2 12 24 12 10 12 Mg +2 V 24 52 X -1 35 45 Y +3 27 14 Z 50 69 Q -1 10 10 III. Identification 1. is the only liquid metal. 2. The horizontal row of elements is called a. 3. The vertical arrangement of elements is called a or. 4. The modern periodic table is greatly similar to the periodic table done by. 5. The element calcium is found under group and period. 15 The Periodic Table
6. The s-block and the p-block elements are collectively called the. 7. The elements on this block have valence electrons on their d subshell.. 8. The first row of the inner transition metals is also known as the. 9. is the most important element under group 15 of the periodic table. 10. is the most metallic element. It is also one of the most reactive elements in the periodic table. 11. is the smallest atom in group 17. 12. is the only group in the periodic table that contains all the three states of matter. 13. is the energy needed to remove an electron from a gaseous atom or ion. 14. is a measure of how easily an approaching electron can bind to the attractive force of the nucleus. 15. The metals usually form colored compounds. They are the elements found on the inner lower box of the periodic table. IV. Arrange the following atoms in the order of increasing Atomic size 1.,,, 2. Na, Mg, Be, Ca Ionic Size 3. Mg +2, Ca, Ca +2 4. S -2, S, O -2 Ionization Energy 5. Na, Ne, P, Ar, K 6. B, C, Al, Si Electron Affinity 7. Ca, Cl, Na The Periodic Table 16
8. K, F, Ba Metallic Character 9. Na, Ba, Mg 10. F, Cl, I 17 The Periodic Table