1. Orbital filling table

2. (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.

3. *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.

4. He stated that elements should be arranged in order of increasing atomic numbers. So, today’s periodic table was formed.

5. 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.

9. 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

10. GROUPS/FAMILIES: The vertical columns 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)

11. * *For A groups; # of valence electrons= # of the group(except He in 8A) in 7A;number of valence electrons=7 However, we can’t say the same thing for B groups. *Lanthanides&Actinides belong to 3B group(the biggest group including 32 elements)

12. PERIODS: The horizontal rows There are 7 periods
Valence shell determines the period number
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.

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

14. Radius of an atom: Half the distance between two nuclei.
2)ATOMIC RADIUS:
Radius of an atom: Half the distance between two nuclei.
Br Br 2r
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.
in 1920, shortly after it had become possible to determine the sizes of atoms using X-ray crystallography . graphy is a method of determining the arrangement of atoms within a crystal, in which a beam of X-rays strikes a crystal and causes the beam of light to spread into many specific directions. From the angles and intensities of these diffracted beams, a crystallographercan produce a three-dimensional picture of the density of electrons within the crystal. From this electron density, the mean positions of the atoms in the crystal can be determined, as well as their chemical bonds, their disorder and various other information. this method determined the size of atoms, the lengths and types of chemical bon ds.

15. 2)ATOMIC RADIUS:
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)
2)Nuclear charge(as the p+ # increases, atomic volume decreases)

16. Atomic volume decreases
Li Be B C N O F Na
WHY?
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. K

19. IONIC VOLUME: X X+ X+2 X- X-2
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+2 X-
X-2

20. Less e-e repulsion
More e-e repulsion

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

22. !!!It’s not the IE(bec. X is a solid and molecular)
3)IONIZATION ENERGY: The minimum amount of energy required to remove the most losely bound e- from one gaseous atoms is called “the ionization energy(I).” X(g) + I X+(g) + e- WARNING!!! X2(s)+I X+(s) + e-
!!!It’s not the IE(bec. X is a solid and molecular)

23. 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.

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

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

26. Atomic volume decreases& IE generally increases
***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

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

28. The variation of IE within the same group:
***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.

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

30. Outer electron further away outweighs greater nuclear charge
Li has lower IE than H
Outer electron further away
outweighs greater nuclear charge H
First Ionization energy
Outweigh:to exceed in importance Li
Atomic number

31. greater nuclear chargeHe
Be has higher IE than Li
Same shielding
greater nuclear charge
First Ionization energy H Be Li
Atomic number

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

33. First Ionization energyHe
First Ionization energy H C Be B Li
Atomic number

34. First Ionization energyHe N
First Ionization energy H C Be B Li
Atomic number

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

36. First Ionization energyHe F N
First Ionization energy H C O Be B Li
Atomic number

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

38. Na has a lower IE than Li Both are s1 Na has more shieldingHe Ne
Na has a lower IE than Li
Both are s1
Na has more shielding
Greater distance F N
First Ionization energy H C O Be B Li Na
Atomic number

39. First Ionization energy
Atomic number

40. 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. Be
1s22s2
1st I.E. = 900 kJ mol-1 B
1s22s22px1
1st I.E. = 799 kJ mol-1

41. 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.
Offset:counterbalance.

42. The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy.

43. 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.
12Mg
1s22s22p63s2
1st I.E. = 736 kJ mol-1
13Al
1s22s22p63s23px1
1st I.E. = 577 kJ mol-1

44. The 3p electron in Aluminium is slightly more distant from the nucleus than the 3s, and partially screened by the 3s2 electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.

45. 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.

46. 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?
7N: 1s22s22px12py12pz1 1st I.E. = 1400 kJ mol-1
8O: 1s22s22px22py12pz1 1st I.E. = 1310 kJ mol-1

47. 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.

48. The screening is identical (from the 1s2 and, to some extent, from the 2s2 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 2px2 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.

49. Increases
Increases

51. ***We can decide about the group number of A group elements by considering their ionization energy values. Mg(g)(1s22s22p63s2)+I1(176Kcal/mole) Mg+(g) + e-Mg+(g)(1s22s22p63s1)+ I2(348Kcal/mole) Mg+2(g)+ e- Mg+2(g) (1s22s22p6)+ I3(1847Kcal/mole) Mg+3(g)+ e- 1st&2nd IE for Mg atom belong to the removal of valence electrons which are bounded very weakly to the nucleus .However,3rd e- requires considerably more energy than the removal of valence e-s as it will experience higher ENC.

52. 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 2nd&3rd ionization energies)

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

54.                                                                                                                                                                                                         
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.

55. 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:

56. 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:
                                                                                                                                          

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

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

60. Electronegativity + – H Cl
08/10/99 + –
Electronegativity describes how electrons are shared in a compound
Consider the compound HCl H Cl
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.

61. Electronegativity + – 0 0 H Cl H H

62. Electronegativity H 2.1 He - Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F
Pauling’s 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 Si
1.8 P S Cl Ar

63. These numbers are derived from several factors including EA, IE, atomic radius
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

64. Calculating EN differences
08/10/99
The first step in defining the polarity of a bond is to calculate electronegativity difference ( EN)
 EN = EN large - EN small
E.g. for NaCl,  EN = 3.2 –0.9 = 2.3, ionic.

65. 6)Metallic and Nonmetallic activity
Metallic activity: tendency to lose electron/become oxidized. As the atomic radius increases, it becomes easier to lose electron.
Nonmetallic activity: tendency to gain electron/become reduced.It becomes difficult to gain electron as the atomic radius increases.

66. 6)METALLIC AND NONMETALLIC CHARACTER:
Metals always lose electrons in compounds.The ease with losing an electron forms the metallic character of elements.Therefore, the elements with low ionization energies reflect the metallic properties well.
Tendecy in gaining electrons forms the nonmetallic character.

67. Increasing nonmetallic character(Except 8A)

68. 6)Metallic and Nonmetallic activity
-Metallic activity decreases
-Nonmetallic activity increases
-Metallic activity increases
-Nonmetallic activity decreases

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

70. Which elements are the most reactive metals and nonmetals in the periodic table?

71. Francium(Fr) and Cesium(Cs) in 1A are the most reactive metals.
Fluorine(F) in 7A is the most reactive nonmetal.

72. Summary Shielding is constant Ionization energy decreases
Electronegativity decreases
Nuclear charge increases
Atomic radius increases
Shielding increases
Ionic size increases
Shielding is constant
Atomic Radius decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases

73. Oxides of period 3 elements
Metallic Oxides in Period 3 Sodium oxide: Na2O ionic Magnesium oxide: MgO ionic Aluminum oxide: Al2O3 ionic Metalloid oxide in Period 3 Silicon dioxide: SiO2 covalent Nonmetallic oxides in Period 3 Tetraphosphorus decoxide: P4O10 covalent Sulfur trioxide: SO3 covalent Dichlorine heptoxide: Cl2O7 covalent

74. 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: Na2O + H2O  2 NaOH basic
Magnesium oxide: MgO + H2O  Mg(OH)2 basic
Aluminum oxide: Al2O3 + H2O  2 Al(OH)3 amphoteric
Metalloid oxide in Period 3 is acidic
Silicon dioxide: SiO2 + H2O  H2SiO3 acidic
Nonmetallic oxides in Period 3 are acidic
Tetraphosphorus decoxide: P4O H2O  4H3PO4 acidic
Sulfur trioxide: SO3 + H2O  H2SO4 acidic
Dichlorine heptoxide: Cl2O7 + H2O  2HClO4 acidic
Argon does not form an oxide

75. OXIDATION STATES OF ELEMENTS
Group IA
IIA
IIIA
IVA VA
VIA
VIIA
VIIIA
MAX(+) +1 +2 +3 +4 +5 +6 +7 NA
MIN(-) -4 -3 -2 -1

76. Important Groups

77. Alkali Metals(IA)(except Fr):
Groups of Elements
Alkali Metals(IA)(except Fr):
Group 1A metals
Soft, silvery colored metals that react violently with H2O to form basic solutions
- They have low melting points and densities due to being the largest atom in their period of the PT.
Because of Fr’s short of supply and nuclear instability.

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

79. 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; most reactive ones are: Cesium & Francium

80. 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.

81. Chemical Reactions of Alkali Metals
Reaction with water (Their rectivity increases as you move down the group)
2Na(s) + H2O(l) Na+(aq) + 2OH- (aq) + H2(g)
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
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.
Fizz:give off bubbles of gas.

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

84. Chemical Reactions of Alkali Metals
Reaction with oxygen,
React with oxygen to produces oxides.
4Li(s) + O2(g) Li2O(s)

85. Chemical Reactions of Alkali Metals
Reaction with halogens
Reaction with halogens produces salts
2Na(s) + Cl2(g) 2NaCl(s)
2K(s) + Br2(g) 2KBr(s)
2Cs(s) + l2(g) 2CsI(s)

86. ALKALINE EARTH METALS(IIA):
Alkaline earth metals share many characteristics with the alkali metals but there are some differences.The group 2 elements are much denser and harder with higher melting points.
While tehese metals are still reactive, they are not as reactive as the alkali metals.

87. They are silvery white metals.
Many are found in rocks in the earth’s crust.

88. HALOGENS(VIIA) except At:
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.
Because of At’s rarity and nuclear instability

89. HALOGENS(VIIA) except At:
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.
Oxidizing power decreases down the group.

90. Most are Poisonous .
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.
Iodine is also useful for eliminating bacteria.Since it is not as reactive as Chlorine, it can be used on humans.

91. HALOGENS(VIIA) except At:
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.
F2, Cl2 -- gas
Br2 --- liquid
I2 -- solid, forms a purple gas on heating.
They all need an electron to become stable, thus form negative ions

92. HALOGENS(VIIA) except At:
Are slightly soluble in water as they are non-polar molecules.
Concentrated solutions of chlorine- green tinge
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 polar solvents like water & ethanol--brown solution.
Tinge:slight degree.

93. HALOGENS(VIIA) except At:
Halogens dissociate slightly in aqueous solutions, forming an acidic solution:
Cl2(aq) + H2O(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 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.

94. Displacement reactions of Halogens
A more reactive halogen is capable of replacing less reactive one from its solution
Cl2 reacts with Br- and I-
Cl2(aq) + 2Br-(aq)  2Cl-(aq) + Br2(l)
Cl2(aq) + 2I-(aq)  2Cl-(aq) + I2(s)
Br2 reacts with I-
Br2(aq) + 2I-(aq)  2Br-(aq) + I2(s)
I2 non-reactive with halide ions

95. HALOGENS(VIIA) except At:
The common insoluble halides (ions of halogens) are those of Pb and Ag.
PbI2 -- is a bright yellow colored, can be used as a test for iodide ion.

96. Oxides of period 3 elements
Metallic Oxides in Period 3 Sodium oxide: Na2O ionic Magnesium oxide: MgO ionic Aluminum oxide: Al2O3 ionic Metalloid oxide in Period 3 Silicon dioxide: SiO2 covalent Nonmetallic oxides in Period 3 Tetraphosphorus decoxide: P4O10 covalent Sulfur trioxide: SO3 covalent Dichlorine heptoxide: Cl2O7 covalent

97. 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: Na2O + H2O  2 NaOH basic
Magnesium oxide: MgO + H2O  Mg(OH)2 basic
Net ionic eqn: O2- + H2O  2 OH-
Aluminum oxide: Al2O3 + 6HCl  2 AlCl3+ 3H2O amphoteric
Al2O3 +2 NaOH + 3H2O  2NaAl(OH)4

98. Oxides of period 3 elements
Metalloid oxide in Period 3 is acidic
Silicon dioxide: SiO2 + H2O  H2SiO3 acidic
(silicic acid)
Nonmetallic oxides in Period 3 are acidic
Tetraphosphorus decoxide: P4O H2O  4H3PO4 acidic
Sulfur trioxide: SO3 + H2O  H2SO4 acidic
Dichlorine heptoxide: Cl2O7 + H2O  2HClO4 acidic
Argon does not form an oxide
Silicic:[suh-lis-ik]

99. NOBEL (Inert)GASES(VIIIA):
They are stable and will rarely react with any other elements.
They are gaseous at room conditions.
Each nobel gas has its valence orbitals full-filled and 8 valence electrons,except He.It has 2 valence electrons.

100. The Nobel gases are used in neon signs
The Nobel gases are used in neon signs.Each nobel gas glows a different colour when electricity is passed through it.For instance, Helium glows pink,neon glows orange-red, and argon glows purple.It isn’t just neon in those signs!

101. Jellyfish lamps made with noble gases