Presentation on theme: "TOPIC 4: The Periodic Table and Some Atomic Properties."— Presentation transcript:
TOPIC 4: The Periodic Table and Some Atomic Properties
2 Contents 1.Classifying the Elements:The Periodic Law and the Periodic Table 2.Description of a modern Periodic Table-The Long Form 3.Electron Configurations and The Periodic Table 4.Metals and Nonmetals and Their Ions 5.The Sizes of Atoms and Ions 6.Ionization Energy 7.Electron affinity 8.Magnetic Properties 9.Periodic Properties of The Elements
3 Periodic Table Alkali metalsAlkaline earth metalsTransition ElementsHalogenes Noble Gases Lanthanides and actinides Main Group
4 1s 2s2p 3s3p 4s3d4p 5s4d5p 6s5d6p 7s6d7p 4f 5f Main Group Elements (IA – VIIIA) Transition Elements (IB – VIIIB) Inner Transition Elements
5 METALS, NONMETALS AND THEİR IONS METALS –In the periodic table there are 110 elements, The total number of metal elements present is 85 of which only 15 elements are defined as artificial. -Metals are good conductors of heat and electricity. They easily lose e - and form + ions. -They possess metalic bonds and are lustrous(shiny). They can form alloys among each other. -They are ductile. Most metals can be drawn out into thin wires. -They are present always at solid phase except for Hg, Ga, Cs ve Fr.None of them is volatile except for Hg. They have high melting and boling points -Active metals (s 1 ) react easily with water, half-active metals (s 2 ) react with water rather difficult, inactive (noble) metals do not react with water. The active metals form the bases.
6 METALS, NONMETALS AND THEİR IONS NONMETALS – Except for H, all of them are the elements of the p orbital element are totally in number 16. -They do not conduct heat and electricity. They gain or share e - easily. They are not shiny. - While F -, O = can only (–) charged ions(Anions), the others can have –/+ ions (Anion/Cation) -They build up covalent bonds among each other, and form ionic bonds with metals and metalloids -They have low melting and boiling points, most of them are present at gas state at normal conditions. Br 2 is liquid ; P, S, I 2, As are solid. -They can not combine with the acids without O. Their Hydrides (Compunds with H) are acids.Nonmetaloxides are acid anhydrides (They react with water and form acids).
7 METALLOIDS METALLOIDS are elements that look like metals and in some way behave like metals, but that have also have some nonmetallic properties. They are present at solid form and have high boiling points.
8 Metals lose e- to achieve noble gas electron configuration Nonmetals gain e- to have noble gas electron configuration.
9 Nonmetal Ions: The atoms of Groups 7A and 6A- the most active nonmetals, have one and two electrons fewer than the noble gas at the end of the period. These atoms can acquire the electron configurations of noble gas atoms by gaining the appropriate number of electrons. Cl ([Ne]3s 2 3p 5 ) + e - Cl - ([Ar]) S ([Ne]3s 2 3p 4 ) + 2e - S -2 ([Ar]) In most cases a nonmetal atom will gain a single electron spontaneously, but energy is required to force it to accept additional electrons. Often other processes occuring simultaneously supply the necessary energy (such as an attraction to positive ions).
10 Transition Metal Ions: Only a few transition metal atoms acquire noble-gas electron configuration by losing electrons (such as in the loss of 3 electrons of Scandium, Sc Sc +3 ). Most transition metal atoms, however, do not acquire a noble- gas configuration when they ionize. Furthermore, a common feature of transition metal is the ability to form more than a single type of ion (Example for this situation:Iron). Fe ([Ar]3d 6 4s 2 ) Fe +2 ([Ar]3d 6 ) + 2e - Fe ([Ar]3d 6 4s 2 ) Fe +3 ([Ar]3d 5 ) + 3e-
11 THE SIZES OF ATOMS AND IONS To understand the physical and chemical properties, we need to know something about atomic sizes. Atomic Radius: The probability of finding an electron decreases with increasing distance from the nucleus, but nowhere does the probability fall to zero. For this reason, there is no precise outer boundary to an atom and it is hard to determine the atomic radius. It is the distance between the nuclei of atoms which can be measured effectively. We define atomic radius in term of internuclear distance (covalent radius, ionic radius, metallic radius).
12 The covalent radius of Br is 1,14 Å. The length of C – C covalent bond is 1,54 Å. Hence, it can be claimed that C has a radius of 0,77 Å.
13 Variation of Atomic Radii Within a Group of the Periodic Table The more electronic shells in an atom, the larger the atom is. Atomic radius increases from top to bottom through a group of elements. The atomic radius decreases from left to right through a period of elements, whereas atomic radii do not change very much within a transition series.
15 Ionic Radius When a metal atom loses one or more electrons to form a positive ion, there is an excess of nuclear charge over the number of electrons in the resulting cation. The nucleus draws the electrons in closer, and as a consequence Cations are smaller than the atoms from which they are formed For isoelectronic(having equal number of electrons) cations, the more positive the ionic charge is, the smaller is the ionic radius. When a nonmetal atom gains one or more electrons to form a negative ion(anion),the nuclear charge remains constant, but Z eff is reduced because of the additional electrons.
16 Repulsions among the electrons increase and they are not held tightly. Anions are larger than the atoms from which they are formed. For isoelectronic anions, the more negative the ionic charge is, the larger is the ionic radius. Practice: Refer only to the periodic table and arrange the following species in order of increasing size: Ar, K +, Ca +2, Cl -, S -2
17 IONIZATION ENERGY Atoms do not eject electrons spontaneously. Electrons are attracted to the positive charge on the nucleus of an atom, and energy is needed to overcome that attraction. The more easily an atom loses its electrons, the more it tends to have a metallic character. Ionization Energy (I) is the quantity of energy a gaseous atom must absorb so that an electron is stripped from the atom. The electron is the one most loosely held. First ionization energy (I 1 ), is the energy required to strip one electron from a neutral gaseous atom. The second ionization energy, I 2 is the energy to strip an electron from a gaseous ion with a charge of+1. Further ionization energies are (I 3, I 4, …) Na (g) Na + (g) + 1e -
18 Ionization Energy (kJ/mol) ElementI1I1 I2I2 I3I3 I4I4 Na Mg Al Ionization energies decrease as atomic radii increase. The required energy to ionize the remaining electrons from the ion of an atom (after an electron from the neutral atom was stripped) is higher than the first ionization energy. The smaller energy levels, the higher ionization energy required to strip an electron from that level.
19 Example: Arrange the following in order of increasing I 1 (As, Sn, Br, Sr)?
20 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 most of the positively charged ions or neutral atoms gain an electron, heat is given off to the environment (exothermic reaction). The EA is indicated with a – sign in this case. [Cl (g) + e - Cl - (g) EA = kJ/mol ]. The EA of Chlorine is kJ/mol. Li (g) + e - Li - (g) EA = - 59,8 kJ/mol (1s 2 2s 1 1s 2 2s 2 – example for the metal atom)
21 Electron Affinity In order that the anions and some neutral atoms gain an electron, the energy of heat must be absorbed from the environment. This event is an endothermic reaction and has a positive value of EA. Ne (g) + e - Ne - (g) EA = + 29 kJ/mol (1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 1 ) O (g) + e - O - (g) EA = - 141,4 kJ/mol O - (g) + e - O -2 (g) EA = kJ/mol
22 The smaller atoms to the right of the periodic table,namely the halogenes(Group VII A) have large negative electron affinities. For example, in becoming Cl -, a chlorine atom acquires the very stable electron configuration of the noble gas Argon(Ar) For the IIA ve VIIIA group elements, the electron gained must enter the higher energy orbital since all the other s and p orbitals of the atoms are filled. This situation requires an endothermic process.
23 Magnetic Properties The behaviour of atoms and ions in a magnetic field is also helpful in establishing electron configurations. A spinning electron is an electric charge in motion. It induces magnetic field. In a diamagnetic atom or ion all electrons are paired, and these individual magnetic effects cancel out. A diamagnetic species is weakly repelled by a magnetic field. A paramagnetic atom or ion has unpaired electrons, and the individual magnetic effects do not cancel out. The unpaired electrons induce a magnetic field that causes the atom or ion to be attracted into an external magnetic field. The more unpaired electrons present, the stronger the attraction.
24 Periodic Properties of the elements: Physical properties (melting points, boiling points, etc. ) Chemical properties (reduction,oxidation) Metal ElementsNonmetal Elements LusterDull appearance Malleable Brittle, high degree of hardness Good conductors of heat and electricity Poor conductors of heat and electricity Their ionic compounds with the oxides have base character Their oxides have acid character They are the cations in the aqueous solutions They are the anions or oxi- anions in the aqueous solutions.
26 1A 3 Li (Lithium) 11 Na (Sodium) 19 K (Potassium) 37 Rb (Rubidium) 55 Cs (Cesium) 87 Fr (Francium) - Each of them has 1 electron in the s orbital. - FROM TOP TO DOWN, the melting point decreases. - The density rises. - Atomic radius rises - First ionization energy lowers ( 1 ) - They are the group having the smallest 1 among the other groups of elements in the periodic table. - They are extremely reactant since their ability to gain or lose electron and to build up ions is high. Therefore, they are present in nature in the form of compounds. (M M + + e - ) Group 1A (Alkali Metals):
27 2M (s) + H 2(g) 2MH (s) They react with hydrogen to form hydrides. They react with sulphur to form sulphides: 2M (s) + S (s) M 2 S (s) They react with chlorine to form chlorides: 2M (s) + Cl 2(g) 2MCl (s) When they react with water, hydrogen gas and the hydroxide of the alkali metal are produced (exothermic reaction): 2M (k) + 2H 2 O (s) 2MOH (aq) + H 2(g) Their reactions with oxygen is more complex: 4Li (s) + O 2(g) 2Li 2 O (s) (lithium oxide) 2Na (s) + O 2(g) Na 2 O 2(s) (sodium peroxide) K, Rb ve Cs may form superoxides K (s) + O 2(g) KO 2(l) (potassium superoxide)
28 - Compared to the Group 1A metals - they have greater hardness - higher densities - higher melting points - higher 1 values and that’s why they are less reactant. - Be and Mg are the least active earth metals. - Calcium and the other elements placed below it react with water at room temperature : Ca (s) + 2H 2 O (l) Ca(OH) 2(aq) + H 2(g) Group 2 A: Alkaline Earth Metals 2A 4 Be (Beryllium) 12 Mg (Magnesium) 20 Ca (Calcium) 38 Sr (Strontium) 56 Ba (Barium) 88 Ra (Radium)
29 GROUP 3A METALS Aluminium is an excellent reducing agent since it is easily oxidized to +3 ions. It’s main use is light- weight alloys Aluminium is such a good reducing agent that it will extract oxygen from metal oxides to produce Aluminiumoxide Fe 2 O 3 (s) + 2 Al (s) Al 2 O 3 (s) + 2 Fe (l) 13 Al (Aluminium) 31 Ga (Gallium) 49 In (Indium) 81 Tl (Thallium)
30 It reacts with acids to produce H 2 (g): 2 Al (s) + 6 H + (aq) 2 Al 3+ (aq) + 3 H 2 (g) Powered Aluminium is easily oxidized by air or other oxidants in highly exothermic reactions used in rocket fuels and explosives Galium is used to make GaAs that can convert light into electricity(photoconduction). This semi-conducting material is also used in light emitting diodes(LED) and in solid state devices such as transistors Indium is a soft silvery metal used to make low-melting point alloys. Thallium and its compounds are extremely toxic. It’s use is in high temperature superconductors
31 GROUP 6A: THE OXYGEN FAMILY 6A 8 O (Oxygen) 16 S (Sulphur) 34 Se (Selenium) 52 Te (Tellurium) 84 Po (Polonium) - From top to bottom the elements tend to show more metalic character - Oxygen is in gaseous form, the other elements are at solid state. - Nonmetals: Oxygen, sulphur and selenium; - Tellurium is metalloide, polonium is metal. - Oxygen can be found in two different molecular forms: O 2 and O 3 (ozon) (allotropic form). 3O 2(g) 2O 3(g) (+ 284,6 kJ) -
32 Oxygen tends to strip electrons from the other elements, and therefore it is a good oxidizing agent. - Oxygen generally is in the form of O -2 when it forms compounds with metals - The other two forms of oxygen anions are O 2 2- and O Sulphur has plenty of allotropic forms of which the most stable one is found as S 8 in yellow color and at solid phase. - It makes up sulphides by gaining electron from other elements (S -2 ).16Na (s) + S 8(s) 8Na 2 S (s) - It is naturally abundant in the form of metal-sulphur compound.
33 GROUP 7A: HALOGENES 7A 9 F (Fluorine) 17 Cl (Chlorine) 35 Br (Bromine) 53 I (Iodine) 85 At (Astatine) - Astatin is a radioactive element is rarely found, so it’s properties are not known. - All halogenes are nonmetal. - Under normal conditions each element is composed of di-atomic molecules - Fluorine: Pale yellow, chlorine: Yellowish green, bromine: Reddish brown, iodine: Violet - Halogenes are the group with greatest electron affinity, When they strip an electron from other elements, this reaction is exothermic.
34 Fluorine and Chlorine are the most reactive ones, fluorine can strip electrons from nearly all of the other elements. - Chlorine react with water and hydrochloric acid and hypochlorous acid are produced : Cl 2 (g) + H 2 O (l) HCl (aq) + HOCl (aq) - Hypochlorous acid is an important desinfectant and used in swimming pools(Chlorine is added into the water). - Halogenes react with plenty of metals to form metallic halogenes Cl 2 (g) + 2Na (s) 2NaCl (s)
35 Grup 8A: THE NOBLE GASES (INERT GASES) 8A 2 He (Helium) 10 Ne (Neon) 18 Ar (Argon) 36 Kr (Krypton) 54 Xe (Xenon) 86 Rn (Radon) - Nonmetal. - At gaseous state at room temperature. - Composed of monoatomic molecules - Their “s” and “p” orbitals are totally filled in - I 1 is extremely high, falls down in the group from top to bottom. - Rn is pretty much reactive, some of it’s properties are not known - Very low reactivity rate - In order to react with, they need the elements with the highest electron affinity Their compounds which we know until now: XeF 2,XeF 4,XeF 6,KrF 2 He, Ne, ve Ar do not have any compounds and are totally inert.
36 THE TRANSITION ELEMENTS d-block and f-block elements of the periodic table High melting points,good electrical conductivity, moderate to extreme hardness Little variation among the atomic radii across the first transition series(Exceptions:Sc and Ti).In the series of elements in which the 4f subshell is filled,atomic sizes actually decreases somewhat. This phenomenon is called lanthanide contraction(lanthanide series) Both ionic and covalent character( ionic character in the compounds with the transition metal in lower oxidation states; covalent character in higher oxidation states
37 The Transition Elements
38 THE TRANSITION ELEMENTS Higher oxidation state in transition metals are stabilized when oxide and/or fluoride ions are bound to the metal Increasing stability of higher oxidation states for the elements further down a group. Ionization energies are fairly constant across the first transition series. Standard electrode potentials increase in value across the series.All these elements are more readily oxidized than hydrogen(Exception: Cu) High catalytic activitity(ability to adsorb gaseous species, some important catalysts: Ni,Fe,Pt,Rh,V 2 O 5,Cr 2 O 3,MnO 2,TiCl 4
39 Different aqueous solutions of the transition elements in various colors
40 THE TRANSITION ELEMENTS Most transition elements are paramagnetic. Fe, Co and Ni are ferromagnetic which means that the magnetic moments are aligned into domains.The ordering of domain persists when the object is removed from the magnetic field and thus permanent magnetism results Group 2B(Zn, Cd and Hg) is sometimes classified as post-transition metals since these elements have the electron configuration with no incomplete d-shell (d 10 s 2 ). They have lower melting points because their full d-subshells prevent d-d bonding.