Periodic_trends - Section 12.1 Classification of the...

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Unformatted text preview: Section 12.1 Classification of the Elements OBJECTIVES: Explain why you can infer the properties of an element based on those of other elements in the periodic table. Section 12.1 Classification of the Elements OBJECTIVES: Use electron configurations to classify elements as noble gases, representative elements, transition metals, or inner transition metals. 12.1 The Development of the Periodic Table Dmitri Mendeleev took the 70 known elements and listed them in several columns based on their physical and chemical properties. The main property he used was atomic mass, but this led to a periodic repetition of other properties, thus creating the first periodic table. In 1913 Henry Moseley determined the nuclear charge(number of protons), also called the atomic number, of the known elements. Moseley arranged these elements based on their atomic number creating the modern periodic table. 12.2 The Modern Periodic Table The modern periodic table is arranged in 7 rows of elements, with increasing atomic number, and 18 columns or families. The 7 horizontal rows are called periods. The families are identified as 2 specific groups, A or B. The "A" group of elements are also known as representative elements. These elements exhibit a wide variety of chemical and physical properties. The "B" group of elements are also know as transition elements and often have more than one combining charge. The sequence of charge is the same in all the periods and this led to the creation of the Periodic Law which states that: "When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties." 12.3 Electron Configuration and Periodicity Of the 3 subatomic particles the electron plays the greatest role in the physical and chemical properties of the elements. There is a direct relationship between the similarity of properties of the elements and their electron configuration(the placement of electrons around the nucleus). How to get the electron configuration of an element The electron structure around an atom is based on the fact that in nature everything seeks the lowest possible energy because this is the most stable of states. In the world of the atom the electrons and nucleus interact to make the most stable electron arrangement possible. This is called the electron configuration of an atom. ) There are 7 main energy (quantum) levels around the nucleus of an atom numbered 1 to 7 ) Within each quantum level there are anywhere from 1 to 4 sublevels labelled s, p, d and f ) Within each sublevel there could be 1, 3, 5 or 7 orbitals ) Within each orbital you will find 0, 1, or 2 electrons Energy Quantum energy level Sublevel - spin electron orbital + spin electron Three rules govern the electron configuration of an atom: ) the Aufbau principle Electrons enter the orbitals of lowest energy(ground state) first. ) The Pauli Exclusion Principle Any orbital may contain at most 2 electrons, but they must be of opposite spin ) Hund's Rule When electrons enter orbitals of equal energy, one electron will enter each orbita; until all orbitals Writing Electron Configurations Order of Filling Sublevels with Electrons the energy sublevels are filled in a specific order that is shown by the arrow diagram seen below: or you can use the block organization of the periodic table to write electron configurations H Li 1 3 1s 1 Na 11 K 19 Rb 37 Cs 55 Fr 87 1s22s 1 1s22s22p63s 1 1s He 2 Ne 2 10 Ar 18 1s 2s 2p 2 2 Kr 36 Xe 6 54 Rn 86 1s22s22p63s 3p Elements can be classified into 4 categories based on their electron configuration. ) The Noble Gases (inert gases which do not have any chemical reactivity) elements with outermost s and p sublevels which are full. ) Representative Elements (columns 1,2,13-17) elements that partially filled outermost s and p sublevels 3) Transition Elements (all metals except columns 1 and 2) elements whose outermost s and nearby d sublevels contain electrons 4) Inner Transition Elements (elements at the bottom of the periodic table) elements whose outermost s and nearby f sublevel generally contain electrons Writing electron configurations the easy way Yes there is a shorthand Electron Configurations repeat The shape of the periodic table is a representation of this repetition. When we get to the end of the column the outermost energy level is full. This is the basis for our shorthand. Write symbol of the noble gas before the element, in [ ]. Then, the rest of the electrons. Aluminum's full configuration: The Shorthand previous noble gas Ne is: 1s22s22p6 so, Al is: [Ne] 3s23p1 1s22s22p63s23p1 More examples Ge = 1s22s22p63s23p64s23d104p2 Hf = 1s22s22p63s23p64s23d104p65s2 4d105p66s24f145d2 Thus, Hf = [Xe]6s24f145d2 Thus, Ge = [Ar] 4s23d104p2 The Shorthand Again Sn- 50 electrons The noble gas before it is Kr Takes care of 36 Next 5s2 Then 4d10 Finally 5p2 [ Kr ] 5s2 4d10 5p2 Trends in Atomic Size First problem: Where do you start measuring from? The electron cloud doesn't have a definite edge. They get around this by measuring more than 1 atom at a time. Atomic Size Atomic Radius = half the distance between two nuclei of a diatomic molecule. Radius } Influenced by three factors: 1. Energy Level 2. Charge on nucleus Trends in Atomic Size Higher energy level is further away. More charge pulls electrons in closer. 3. Shielding effect (blocking effect?) Group trends As we go down a group... each atom has another energy level, so the atoms get bigger. H Li Na K Rb As you go across a period, the radius gets smaller. Electrons are in same energy level. More nuclear charge. Outermost electrons are closer. Periodic Trends Na Mg Al Si P S Cl Ar Rb K Atomic Radius (nm) Overall Na Li Ar H Ne Kr 10 Atomic Number Trends in Ionization Energy The amount of energy required to completely remove an electron from a gaseous atom. Removing one electron makes a 1+ ion. The energy required to remove the first electron is called the first ionization energy. Ionization Energy The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron. Greater than 1st or 2nd IE. Ionization Energies Table 12.1, p. 281 kJ/ mol Symbol First H He Li Be B C N O F Ne Atom 1312 2731 520 900 800 1086 1402 1314 1681 2080 Second 5247 7297 1757 2430 2352 2857 3391 3375 3963 Third Ion Formed 11810 Be 14840 3569 2+ 4619 4577 5301 6045 6276 What determines IE The greater the nuclear charge, the greater IE. Greater distance from nucleus decreases IE Filled and halffilled orbitals have lower energy, so achieving them is easier, lower IE. Shielding effect Shielding The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. Second electron has same shielding, if it is in the same period + Group trends As you go down a group, first IE decreases because... The electron is further away. More shielding. Periodic trends All the atoms in the same period have the same energy level. Same shielding. But, increasing nuclear charge So IE generally increases from left to right. Exceptions at full and 1/2 full orbitals. He First Ionization energy H He has a greater IE than H. same shielding greater nuclear charge Atomic number He Ne N F q Na First Ionization energy H Be B C O Li Na has a lower IE than Li q Both are s1 q Na has more shielding q Greater distance Atomic number H to Br First Ionization energy kJ/mol Atomic number Driving Force Full Energy Levels require lots of energy to remove their electrons. Noble Gases have full orbitals. Atoms behave in ways to achieve noble gas configuration. 2nd Ionization Energy For elements that reach a filled or halffilled orbital by removing 2 electrons, 2nd IE is lower than expected. True for s2 Alkaline earth metals form 2+ ions. 3rd IE Using the same logic s2p1 atoms have an low 3rd IE. Atoms in the aluminum family form 3+ ions. 2nd IE and 3rd IE are always higher than 1st IE!!! Trends in Electron Affinity The energy change associated with adding an electron to a gaseous atom. Easiest to add to group 7A. Gets them to full energy level. Increase from left to right: atoms become smaller, with greater nuclear charge. Decrease as we go down a group. Trends in Ionic Size Cations form by losing electrons. Cations are smaller than the atom they come from. Metals form cations. Cations of representative elements have noble gas configuration. Ionic size Anions form by gaining electrons. Anions are bigger that the atom they come from. Nonmetals form anions. Anions of representative elements have noble gas configuration. Configuration of Ions Ions always have noble gas configuration. Na is: 1s22s22p63s1 Forms a 1+ ion: 1s22s22p6 Same configuration as neon. Metals form ions with the configuration of the noble gas before them they lose electrons. Configuration of Ions Nonmetals form ions by gaining electrons to achieve noble gas configuration. They end up with the configuration of the noble gas after them. Group trends Adding energy level Ions get bigger as you go down. Li1+ Na1+ K1+ Rb1+ Cs1+ Periodic Trends Across the period, nuclear charge increases so they get smaller. Energy level changes between anions and cations. Li1+ Be B 2+ 3+ N3- O2- F1- C4+ Size of Isoelectronic ions Iso means the same Iso electronic ions have the same # of electrons Al3+ Mg2+ Na1+ Ne F1 O2 and N3 all have 10 electrons all have the configuration: 1s22s22p6 Size of Isoelectronic ions Positive ions that have more protons would be smaller. Ne F1O2N3- Al3+ Na1+ Mg2+ Electronegativity The tendency for an atom to attract electrons to itself when it is chemically combined with another element. How fair is the sharing? Big electronegativity means it pulls the electron toward it. Atoms with large negative electron affinity have larger electronegativity. Electronegativity Video Group Trend The further down a group, the farther the electron is away, and the more electrons an atom has. More willing to share. Low electronegativity. Metals are at the left of the table. They let their electrons go easily Low electronegativity At the right end are the nonmetals. They want more electrons. Try to take them away from others High electronegativity. Periodic Trend Summary: Fig. 12.10, p.285 Look at the Families Group IA - The Alkali Metals (Li, Na, K, Rb, Cs, Fr) Highly colored in flames = fireworks Group IIA - The Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra) Lose 2 valence electrons Also react with H2O to form an alkaline solution (basic), and hydrogen gas, but less violently Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g Group IIA - The Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra) Various forms of CaCO3 Group IIA - The Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra) Strong reaction of magnesium with oxygen to produce magnesium oxide Mg(s) + O2(g) MgO(s) flashbulbs Boron is mined in the form of Borax, and is used in laundry soap Laboratory glassware contains borosilicates Group IIIA - The Boron Family (B, Al, Ga, In, Tl) Group IIIA - The Boron Family (B, Al, Ga, In, Tl) Aluminum metal is the most abundant metal in the earth's crust and has many uses Group IIIA - The Boron Family (B, Al, Ga, In, Tl) Gallium Arsenide is used in the manufacture of computer chips Group IVA - The Carbon Family (C, Si, Ge, Sn, Pb) Carbon is essential for life and is found in all organic molecules Group IVA - The Carbon Family (C, Si, Ge, Sn, Pb) Carbon is found in different structures or ALLOTROPES - graphite, one of the softest substances known, and diamond, the hardest Group IVA - The Carbon Family (C, Si, Ge, Sn, Pb) Quartz or SiO2 Elemental Si is used in the semiconductor industry Group VA - The Nitrogen Family (N, P, As, Sb, Bi) There are two varieties of P, red and white. Group VIA - The Oxygen Family (O, S, Se, Te, Po) Stratospheric ozone shields us from harmful UV The ozone radiation. Ozone is destroyed by Cl-containing "hole" over molecules used in refrigeration Antarctica Group VIIA - The Halogens (F, Cl, Br, I, At) Br2 and I2 Halogen mean "salt-former". Here sodium metal reacts vigorously with Cl2(g) Group VIIIA - The Noble (Inert) Gases He, Ne, Ar, Kr, Xe The Lights of Las Vegas Helium-Neon lasers The Transition Elements An important use of transition elements is as pigments in paints and glasses ...
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This note was uploaded on 07/06/2011 for the course SCH 4U taught by Professor White during the Spring '10 term at Beacon FL.

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