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Orbital filling table (1869) Dmitri Mendeleev(Russian chemist) shows a first version of the periodic table. He noticed that classifying the elements.

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Orbital filling table (1869) Dmitri Mendeleev(Russian chemist) shows a first version of the periodic table. He noticed that classifying the elements.

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Presentation on theme: "Orbital filling table (1869) Dmitri Mendeleev(Russian chemist) shows a first version of the periodic table. He noticed that classifying the elements."— Presentation transcript:

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2 Orbital filling table

3 (1869) Dmitri Mendeleev(Russian chemist) shows a first version of the periodic table. He noticed that classifying the elements by their atomic mass a periodicity in certain properties could be seen. The first table consisted of 63 elements. Periodicity: the regular repeating of properties according to the arrangement of elements in the PT.

4 *Henry Mosely(British) discovered nuclear charges of all known elements and that chemical properties of elements are related to their atomic numbers but not atomic weights.

5 He stated that elements should be arranged in order of increasing atomic numbers. So, todays periodic table was formed.

6 In modern periodic table, elements are listed in order of increasing atomic numbers. Elements with similar chemical properties are placed in the same vertical columns.

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10 1A:Alkali metals 2A:Alkaline earth metals 3A:Earth metals 4A:Carbon Family 5A:Nitrogen Family 6A:Oxygen Family 7A:Halogens 8A:Noble(Inert)gases

11 GROUPS/FAMILIES: The vertical columns The vertical columns Elements in the same group have Elements in the same group have -similar chemical properties(exception: H in 1A group) -Same number of valence electrons and orbitals.(exception:He in 8A group) - The effective nuclear charge (the charge acting on valence electrons) is the same.(exception:He in 8A group)

12 * *For A groups; # of valence electrons= # of the group(except He in 8A) in 7A;number of valence electrons=7 However, we cant say the same thing for B groups. *Lanthanides&Actinides belong to 3B group(the biggest group including 32 elements)

13 PERIODS: The horizontal rows The horizontal rows There are 7 periods There are 7 periods Valence shell determines the period number Valence shell determines the period number Each of them starts with a metal and ends with a noble gas.(except first and seventh ones) Each of them starts with a metal and ends with a noble gas.(except first and seventh ones) Elements in the same period have the same # of energy levels or shells or principle quantum numbers. Elements in the same period have the same # of energy levels or shells or principle quantum numbers.

14 PERIODIC TRENDS: PERIODIC TRENDS: 1)ATOMIC&MASS NUMBER: AN,MN increases

15 2)ATOMIC RADIUS: Br Br 2r Radius of an atom : Half the distance between two nuclei. Covalent radius: Half the distance between two nuclei in a covalent molecule consisting of identical atoms. Van der Waals radius: This is for 0 group gases. Ionic radius: This for ions in an ionic compound.

16 2)ATOMIC RADIUS: Atomic size (volume,radius) is affected by mainly two factors in the periodic table: Atomic size (volume,radius) is affected by mainly two factors in the periodic table: 1)The # of shells (as it increases, atomic volume also increases) 1)The # of shells (as it increases, atomic volume also increases) 2)Nuclear charge(as the p + # increases, atomic volume decreases) 2)Nuclear charge(as the p + # increases, atomic volume decreases)

17 Li Be B C N O F Na Atomic volume decreases WHY? K 1A 2A3A 4A5A 7A 6A Within the same period;All the elements have the same # of shells but the p + # increases from left to right.Therefore,atomic radius decreases from left to right.

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20 Proceeding down a group; The # of shells of atoms increases but the p + # of atoms also increases.However,the increase in shells makes a bigger effect on the radius than the nuclear charge.Therefore, the atomic volume or radius increases down a group. IONIC VOLUME: X X+X+ X +2 X-X- X -2

21 Less e-e repulsion More e-e repulsion

22 For isoelectronic species; The greater the nuclear charge,the smaller the radius(or volume) 7 N -3, 8 O -2, 11 Na +, 13 Al +3 are isoelectronics.The relationship between their radii; 7 N -3 > 8 O -2 > 11 Na + > 13 Al +3

23 3)IONIZATION ENERGY: The minimum amount of energy required to remove the most losely bound e - s from a mole of gaseous atoms is called the ionization energy(I). X (g) + I 1 X + (g) + e - WARNING!!! X 2 (s) +I 1 X + (s) + e - !!!Its not the IE(bec. X is a solid and molecular)

24 3)IONIZATION ENERGY: The energy required to remove (1 mole of) the first electron from 1 mole of gaseous atom is called the first ionization energy. The energy required to remove (1 mole of) the first electron from 1 mole of gaseous atom is called the first ionization energy.

25 Ionization Energy The second ionization energy is the energy required to remove (1 mole of) the second electron(s). The second ionization energy is the energy required to remove (1 mole of) the second electron(s). Always greater than first IE. Always greater than first IE. The third IE is the energy required to remove a third electron. The third IE is the energy required to remove a third electron. Greater than 1st or 2nd IE. Greater than 1st or 2nd IE.

26 What determines IE The greater the nuclear charge, the greater IE. The greater the nuclear charge, the greater IE. The greater the effective nuclear charge, the greater IE. The greater the effective nuclear charge, the greater IE. Greater distance from nucleus (atomic radius) decreases IE Greater distance from nucleus (atomic radius) decreases IE Shielding of electrons in filled inner orbitals Shielding of electrons in filled inner orbitals

27 *** Because I of d block elements are irregular, rules that we talk about the IE belong to A group elements. The variation of first ionization energies within the same period: As the atomic volume increases, the attraction of the nucleus on the electrons decreases. *Ionization energies in the same period: Noble gases > nonmetals > metals Atomic volume decreases& IE generally increases

28 Ionization energy All the atoms in the same period have the same energy level. All the atoms in the same period have the same energy level. Same shielding. Same shielding. But, increasing nuclear charge and effective nuclear charge, ENC. But, increasing nuclear charge and effective nuclear charge, ENC. So IE generally increases from left to right. So IE generally increases from left to right.

29 *** In the same period from left to right the ionization energies: 1A < 3A < 2A < 4A < 6A < 5A < 7A < 8A irregularities The variation of IE within the same group: Down the group, atomic volumes of elements increase and more shielding effect, same ENC.Therefore, the IE of elements decrease in the same group from top to bottom.

30 First Ionization energy Atomic number He He has a greater IE than H. same shielding greater nuclear charge H

31 First Ionization energy Atomic number H He Li has lower IE than H Outer electron further away l outweighs greater nuclear charge Li

32 First Ionization energy Atomic number H He Be has higher IE than Li Same shielding l greater nuclear charge Li Be

33 First Ionization energy Atomic number H He B has lower IE than Be B has greater shielding greater nuclear charge l p orbital is slightly more diffuse and its electron easier to remove Li Be B

34 First Ionization energy Atomic number H He Li Be B C

35 First Ionization energy Atomic number H He Li Be B C N

36 First Ionization energy Atomic number H He Li Be B C N O Breaks the pattern, because the outer electron is paired in a p orbital and experiences inter- electron repulsion.

37 First Ionization energy Atomic number H He Li Be B C N O F

38 First Ionization energy Atomic number H He Li Be B C N O F Ne Ne has a lower IE than He Ne has a lower IE than He Both are full, Both are full, Ne has more shielding Ne has more shielding Greater distance Greater distance

39 First Ionization energy Atomic number H He Li Be B C N O F Ne Na has a lower IE than Li Both are s 1 Na has more shielding l Greater distance Na

40 First Ionization energy Atomic number

41 Why the drop between groups IIA and IIIA(Be-B)? The explanation lies with the structures of Boron and Aluminium. The outer electron is removed more easily from these atoms than the general trend in their period would suggest. Be1s 2 2s 2 1st I.E. = 900 kJ mol -1 B1s 2 2s 2 2p x 1 1st I.E. = 799 kJ mol -1

42 You might expect the Boron value to be more than the Beryllium value because of the extra proton. Offsetting that is the fact that Boron's outer electron is in a 2p orbital rather than a 2s. 2p orbitals have a slightly higher energy than the 2s orbital, and the electron is, on average, to be found further from the nucleus. This has two effects. The increased distance results in a reduced attraction and so a reduced ionisation energy.

43 The 2p orbital is screened not only by the 1s 2 electrons but, to some extent, by the 2s 2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy.

44 Why the drop between groups IIA and IIIA(Mg-Al)? The explanation for the drop between Magnesium and Aluminium is the same, except that everything is happening at the 3-level rather than the 2-level. 12 Mg 1s 2 2s 2 2p 6 3s 2 1st I.E. = 736 kJ mol Al 1s 2 2s 2 2p 6 3s 2 3p x 1 1st I.E. = 577 kJ mol -1

45 The 3p electron in Aluminium is slightly more distant from the nucleus than the 3s, and partially screened by the 3s 2 electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.

46 If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.

47 Why the drop between groupsVA and VIA (N-O and P-S)? Once again, you might expect the ionisation energy of the group VIA element to be higher than that of group VA because of the extra proton. What is offsetting it this time? 7 N: 1s 2 2s 2 2p x 1 2p y 1 2p z 1 1st I.E. = 1400 kJ mol -1 8 O: 1s 2 2s 2 2p x 2 2p y 1 2p z 1 1st I.E. = 1310 kJ mol-1

48 Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.

49 The screening is identical (from the 1s 2 and, to some extent, from the 2s 2 electrons), and the electron is being removed from an identical orbital. The difference is that in the Oxygen case the electron being removed is one of the 2p x 2 pair. The repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be. The drop in ionisation energy at Sulphur is accounted for in the same way.

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52 ***We can decide about the group number of A group elements by considering their ionization energy values. Mg (g) (1s 2 2s 2 2p 6 3s 2 )+I 1 (176Kcal/mole) Mg + (g) + e - Mg + (g) (1s 2 2s 2 2p 6 3s 1 )+ I 2 (348Kcal/mole) Mg +2 (g) + e - Mg +2 (g) (1s 2 2s 2 2p 6 )+ I 3 (1847Kcal/mole) Mg +3 (g) + e - 1 st &2 nd IE for Mg atom belong to the removal of valence electrons which are bounded very weakly to the nucleus.However,3 rd e - requires considerably more energy than the removal of valence e - s as it will experience higher ENC.

53 WARNING!!! There is always needed much more energy to remove inner electrons than the outer electrons since the inner electrons will experience a higher ENC. The sharp jumps between ionization energies help us to find out group numbers of A group elements. We can say that Mg atom is in 2A group by considering the sharp increase between its second and third ionization energies.(because,there is very sharp increase between its 2 nd &3 rd ionization energies)

54 electron being lost: 1st 2nd 3rd 4th 5th 6th 7th (2A) (3A) (4A) (5A) (6A)

55 The rare gases (He, Ne, Ar, Kr, Xe, Rn) appear at peak values of ionization energy, which reflect their chemical inertness, while the alkali metals (Li, Na, K, Rb, Cs) appear at minimum values of ionization energy, in keeping with their reactivity and ease of cation formation.

56 4)ELECTRON AFFINITY: The electron affinity is the energy that is released when an atom in the gas phase gains an electron and is thus converted to an anion, also in the gas phase: The electron affinity is the energy that is released when an atom in the gas phase gains an electron and is thus converted to an anion, also in the gas phase:

57 Electron affinities are difficult to measure and there is no reliable data available for most elements. However, the larger the atom, the lower its electron affinity, as shown with Group VII elements:

58 For reasons outside the scope of this discussion, the electron affinity of Fluorine is an exception to this trend.

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60 5) ELECTRONEGATIVITY: Electronegativity is the power of an atom to attract electron density in a covalent bond (Linus Pauling) Electronegativity is the power of an atom to attract electron density in a covalent bond (Linus Pauling)

61 Electronegativity Electronegativity describes how electrons are shared in a compound Electronegativity describes how electrons are shared in a compound Consider the compound HCl Consider the compound HCl The electron clouds represent where the two electrons in the HCl bond spend their time (sizes of atoms are not being shown) The shared electrons spend more time around Cl than H. In other words Cl is more electronegative than H. HClCl + –

62 Electronegativity + – 0 0 HClHH

63 Electronegativity Paulings electronegativity scale H 2.1 He - Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Ne - Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 Ar -

64 These numbers are derived from several factors including EA, IE, atomic radius These numbers are derived from several factors including EA, IE, atomic radius You do not need to understand where the numbers come from You do not need to understand where the numbers come from You need to know that a high number means the element has a greater pull on electrons You need to know that a high number means the element has a greater pull on electrons

65 Calculating EN differences The first step in defining the polarity of a bond is to calculate electronegativity difference ( EN) The first step in defining the polarity of a bond is to calculate electronegativity difference ( EN) EN = EN large - EN small EN = EN large - EN small E.g. for NaCl, EN = 3.2 –0.9 = 2.3, ionic. E.g. for NaCl, EN = 3.2 –0.9 = 2.3, ionic.

66 Bigger electronegativity, Bigger electronegativity, -Bigger tendency in gaining electrons.

67 Summary Shielding is constant Atomic Radius decreases Ionization energy increases Electronegativity increases Nuclear charge increases Shielding increasesAtomic radius increasesIonic size increasesIonization energy decreasesElectronegativity decreases

68 OXIDATION STATES OF ELEMENTS Group IAIIAIIIAIVAVAVIAVIIAVIII A MAX (+) NA MIN (-) NA NA

69 Important Groups

70 Groups of Elements Alkali Metals(IA)(except Fr): –Group 1A metals –Soft, silvery colored metals that react violently with H 2 O to form basic solutions - They have low melting points and densities due to being the largest atom in their period of the PT. - They have low melting points and densities due to being the largest atom in their period of the PT.

71 Going down the group, the metals get softer and mp decreases due to increase in atomic size. Going down the group, the metals get softer and mp decreases due to increase in atomic size. Alkali Metals(IA)(except Fr):Alkali Metals(IA)(except Fr):

72 - They are the most reactive of all the metals on the periodic table since they can easily lose their one e. - Their rectivity also increases as you move down their family.Therefore; - Their rectivity also increases as you move down their family.Therefore; most reactive ones are: Cesium & Francium

73 Because the elements are very reactive, they easily combine with water and Oxygen.Therefore, most of these elements are not found freely in nature.That is why they are stored in oil in a bottle in the laboratory in order to prevent any reaction with Oxygen. Because the elements are very reactive, they easily combine with water and Oxygen.Therefore, most of these elements are not found freely in nature.That is why they are stored in oil in a bottle in the laboratory in order to prevent any reaction with Oxygen. They tarnish (lose lustre) rapidly when exposed to the air. They tarnish (lose lustre) rapidly when exposed to the air.

74 Chemical Reactions of Alkali Metals Reaction with water ( Their rectivity increases as you move down the group) Reaction with water ( Their rectivity increases as you move down the group) 2Na(s) + H 2 O(l) 2Na + (aq) + 2OH - (aq) + H 2 (g) 2Na(s) + H 2 O(l) 2Na + (aq) + 2OH - (aq) + H 2 (g) With lithium, the reaction occurs slowly and steadily With lithium, the reaction occurs slowly and steadily In the case of sodium, the reaction is vigorous, producing enough heat to melt the sodium which fizzes around the surface quite vigorously In the case of sodium, the reaction is vigorous, producing enough heat to melt the sodium which fizzes around the surface quite vigorously With potassium the reaction is violent and the heat produced is enough to ignite the hydrogen gas evolved, which burns with a purple flame. With potassium the reaction is violent and the heat produced is enough to ignite the hydrogen gas evolved, which burns with a purple flame. They produce alkaline solution and hydrogen gas as a result of the rxn w/ water. They produce alkaline solution and hydrogen gas as a result of the rxn w/ water.

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76 Li K Na Alkali Metal Family They give caharacteristic colour in the flame(chemical peoperty).

77 Chemical Reactions of Alkali Metals Reaction with oxygen, Reaction with oxygen, React with oxygen to produces oxides. React with oxygen to produces oxides. 4 Li(s) + O 2 (g) 2 Li 2 O(s) 4 Li(s) + O 2 (g) 2 Li 2 O(s)

78 Chemical Reactions of Alkali Metals Reaction with halogens Reaction with halogens Reaction with halogens produces salts Reaction with halogens produces salts 2Na(s) + Cl 2 (g)2NaCl(s) 2Na(s) + Cl 2 (g)2NaCl(s) 2K(s) + Br 2 (g) 2KBr(s) 2K(s) + Br 2 (g) 2KBr(s) 2Cs(s) + l 2 (g)2CsI(s) 2Cs(s) + l 2 (g)2CsI(s)

79 Exist as diatomic molecules in which atoms are joined by a single cov. bond. Exist as diatomic molecules in which atoms are joined by a single cov. bond. The elements in this group are referred to as Halogens because they produce salts when combined with alkali metals (e.g.NaCl).these salts are usually white & soluble in water. The elements in this group are referred to as Halogens because they produce salts when combined with alkali metals (e.g.NaCl).these salts are usually white & soluble in water. HALOGENS(VIIA) except At:HALOGENS(VIIA) except At:

80 They are all very reactive and quite electronegative nonmetals. The ease w/ which they gain electrons decreases going down the group. They are all very reactive and quite electronegative nonmetals. The ease w/ which they gain electrons decreases going down the group. Halogens tend to be less reactive as you move down the group.Flourine is the most reactive Halogen and combines with other elements very readily. Halogens tend to be less reactive as you move down the group.Flourine is the most reactive Halogen and combines with other elements very readily. Oxidizing power decreases down the group. Oxidizing power decreases down the group. HALOGENS(VIIA) except At:HALOGENS(VIIA) except At:

81 Most are Poisonous. Most are Poisonous. When Fluorine combines with Na to form NaF,it is an effective cavity fighter that is added to toothpastes. When Fluorine combines with Na to form NaF,it is an effective cavity fighter that is added to toothpastes. Chlorine is a great bacteria fighter so it is used in swimming pools and household cleaning agents. Chlorine is a great bacteria fighter so it is used in swimming pools and household cleaning agents. Iodine is also useful for eliminating bacteria.Since it is not as reactive as Chlorine, it can be used on humans. Iodine is also useful for eliminating bacteria.Since it is not as reactive as Chlorine, it can be used on humans.

82 Going down the group, their physical state varies at room temp & pressure depending on the Van Der Waals force strength present between molecules since the molecules have different molecular mass values. Going down the group, their physical state varies at room temp & pressure depending on the Van Der Waals force strength present between molecules since the molecules have different molecular mass values. F 2, Cl 2 -- gas F 2, Cl 2 -- gas Br liquid Br liquid I 2 -- solid, forms a purple gas on heating. I 2 -- solid, forms a purple gas on heating. They all need an electron to become stable, thus form negative ions They all need an electron to become stable, thus form negative ions HALOGENS(VIIA) except At:HALOGENS(VIIA) except At:

83 Are slightly soluble in water as they are non-polar molecules. Are slightly soluble in water as they are non-polar molecules. Concentrated solutions of chlorine- green tinge Concentrated solutions of chlorine- green tinge Solutions of bromine- darken from yellow through orange to brown as the concentration increases. Solutions of bromine- darken from yellow through orange to brown as the concentration increases. Iodine dissolved in non-polar solvents like hexane-- violet solution. Iodine dissolved in non-polar solvents like hexane-- violet solution. Iodine dissolved in polar solvents like water & ethanol-- brown solution. Iodine dissolved in polar solvents like water & ethanol-- brown solution. HALOGENS(VIIA) except At:HALOGENS(VIIA) except At:

84 Halogens dissociate slightly in aqueous solutions, forming an acidic solution: Halogens dissociate slightly in aqueous solutions, forming an acidic solution: Cl 2 (aq) + H 2 O(l) H + (aq)+ Cl - (aq) + HOCl(aq) Cl 2 (aq) + H 2 O(l) H + (aq)+ Cl - (aq) + HOCl(aq) HOCl(hypochlorous acid): weak acid, reacts as an oxidant since it donates its one oxygen.It oxidises colored dyes to colorless products. HOCl(hypochlorous acid): weak acid, reacts as an oxidant since it donates its one oxygen.It oxidises colored dyes to colorless products. HOCl turns blue litmus paper into red, and then make it colorless. Therefore, HOCl and OCl - are used in bleaches.They are also toxic to microbes. HOCl turns blue litmus paper into red, and then make it colorless. Therefore, HOCl and OCl - are used in bleaches.They are also toxic to microbes. HALOGENS(VIIA) except At:HALOGENS(VIIA) except At:

85 Displacement reactions of Halogens A more reactive halogen is capable of replacing less reactive one from its solution Cl 2 reacts with Br - and I - Cl 2 (aq) + 2Br - (aq) 2Cl - (aq) + Br 2 (l) Cl 2 (aq) + 2I - (aq) 2Cl - (aq) + I 2 (s) Br 2 reacts with I- Br 2 (aq) + 2I - (aq) 2Br - (aq) + I 2 (s) I 2 non-reactive with halide ions 84

86 The common insoluble halides (ions of halogens) are those of Pb and Ag. The common insoluble halides (ions of halogens) are those of Pb and Ag. PbI 2 -- is a bright yellow colored, can be used as a test for iodide ion. PbI 2 -- is a bright yellow colored, can be used as a test for iodide ion. Look at P. 77 Look at P. 77 & 78 for the colors of halogens and tests of halide ions. HALOGENS(VIIA) except At:HALOGENS(VIIA) except At:

87 Oxides of period 3 elements Metallic Oxides in Period 3 Sodium oxide: Na 2 Oionic Magnesium oxide: MgOionic Aluminum oxide: Al 2 O 3 ionic Metalloid oxide in Period 3 Silicon dioxide: SiO 2 covalent Nonmetallic oxides in Period 3 Tetraphosphorus decoxide: P 4 O 10 covalent Sulfur trioxide: SO 3 covalent Dichlorine heptoxide: Cl 2 O 7 covalent 86

88 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3. Acidic/Basic Metallic oxides in Period 3 are basic Sodium oxide: Na 2 O + H 2 O 2 NaOH basic Magnesium oxide: MgO + H 2 O Mg(OH) 2 basic Net ionic eqn: O 2- + H 2 O 2 OH - Aluminum oxide: Al 2 O 3 + 6HCl 2 AlCl 3 + 3H 2 O amphoteric Al 2 O 3 +2 NaOH + 3H 2 O 2NaAl(OH) 4 Al 2 O 3 +2 NaOH + 3H 2 O 2NaAl(OH) 4 87

89 Oxides of period 3 elements Metalloid oxide in Period 3 is acidic Silicon dioxide:SiO 2 + H 2 O H 2 SiO 3 acidic (silicic acid) Nonmetallic oxides in Period 3 are acidic Tetraphosphorus decoxide: P 4 O H 2 O 4H 3 PO 4 acidic Sulfur trioxide: SO 3 + H 2 O H 2 SO 4 acidic Dichlorine heptoxide: Cl 2 O 7 + H 2 O 2HClO 4 acidic Argon does not form an oxide

90 Look at!! P.80 P.80


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