Chapter 5: The Periodic Law 5.1 History of the Periodic Table I. M
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Chapter 5: The Periodic Law 5.1 History of the Periodic Table I. M

Chapter 5: The Periodic Law 5.1 History of the Periodic Table I. Mendeleev, Moseley, and the Modern Periodic Table A. Mendeleev: 1. Created a table in which elements with similar properties were grouped together when arranged in order of increasing atomic mass—a periodic table of the elements. 2. Predicted the existence and properties of elements yet to be discovered (gave more credence to his periodic table) B. Moseley: 1. There were problems with Mendeleev’s table that Moseley discovered could be fixed by arranging elements according to atomic number rather than atomic mass. 2. Elements are arranged in order of increasing number of protons (atomic number) C. The Periodic Law 1. The physical and chemical properties of the elements are periodic functions of their atomic numbers 2. Elements on the table are arranged in order of increasing atomic number (number of protons) 5.2 Electron Configuration and the Periodic Table I. Metals, Nonmetals, and Metalloids A. Metals 1. Good conductors of heat and electricity 2. Lustrous (shiny) 3. Solids (except mercury) 4. Ductile (can be drawn into wire) 5. Malleable (can be hammered into thin sheets) B. Nonmetals 1. Poor conductors of heat and electricity 2. Most are gaseous 3. Solids tend to be brittle C. Metalloids 1. Some properties of metals, some of nonmetals II. III. Periods and the Blocks of the Periodic Table A. Periods 1. Horizontal rows on the periodic table 2. Period number corresponds to the highest energy level of the elements in the period (the energy level with the valence electrons) B. Sublevel Blocks 1. Periodic table can be broken into blocks corresponding to s, p, d, f sublevels Blocks and Groups A. s-­‐Block, Groups 1 and 2 1. Group 1 -­‐ The alkali metals

a. One s electron in outer shell b. Soft, silvery metals of low density and low melting points c. Highly reactive, never found pure in nature 2. Group 2 -­‐ The alkaline earth metals a. Two s electrons in outer shell b. Denser, harder, stronger, less reactive than Group 1 c. Too reactive to be found pure in nature B. d-­‐Block, Groups 3 -­‐ 12 1. Metals with typical metallic properties 2. Referred to as "transition" metals 3. Group number = sum of outermost s and d electrons C. p-­‐Block elements, Groups 13 -­‐ 18 1. Properties vary greatly a. Metals (1) softer and less dense than d-­‐block metals (2) harder and more dense than s-­‐block metals b. Metalloids (1) Brittle solids with some metallic and some nonmetallic properties (2) Semiconductors c. Nonmetals (1) Halogens (Group 17) are most reactive of the nonmetals D. f-­‐Block, Lanthanides and Actinides (inner transition metals) 1. In the periodic table, the f-­‐block elements are wedged between Groups 3 and 4 in the sixth and seventh periods. 2. Lanthanides are shiny metals similar in reactivity to the Group 2 metals 3. Actinides a. All are radioactive b. Transuranium: elements after uranium-­‐ all man-­‐made IV. Periodic Facts: 1. 92 naturally occurring elements 2. Elements 93 and above (transurnaium) are all man-­‐made 3. Elements 43 and 61 (technetium and promethium) are also synthetic (man-­‐made) 4. All elements after #83 (84 polonium and above) are radioactive 5.3 Periodic Trends I. Atomic Radii A. Atomic Radius 1. One half the distance between nuclei of identical atoms that are bonded together B. Trends 1. Atomic radius tends to decrease across a period due to increasing positive nuclear charge 2. Atomic radii tend to increase down a group due to increasing number energy levels (outer electrons are farther from the nucleus) II. Ionization Energy A. Ion 1. An atom or a group of atoms that has a positive or negative charge 2. Positive charge: cation (loss of electrons)

3. Negative charge: anion (gain of electrons) B. Ionization 1. Any process that results in the formation of an ion C. Ionization Energy (IE) 1. The energy required to remove one electron from a neutral atom of an element, measured in kilojoules/mole (kJ/mol) ! + !"!#$% → ! ! + ! ! D. Trends 1. Ionization energy of main-­‐group elements tends to increase across each period a. Atoms are getting smaller, electrons are closer to the nucleus (increased nuclear charge = stronger attraction for electrons in the same energy level) 2. Ionization energy of main-­‐group elements tends to decrease as atomic number increases in a group (down the group) a. Atoms are getting larger; electrons are farther from the nucleus b. Outer electrons become increasingly more shielded from the nucleus by inner electrons = easier to remove 3. Metals have a characteristic low ionization energy 4. Nonmetals have a high ionization energy 5. Noble gases have a very high ionization energy E. Removing Additional Electrons !" + !"#$% !"# → !" ! + ! ! !" ! + !"#$%& !"# → !" !! + ! ! !" !! + !"#$%& !"# → !" !!! + ! _ 1. Ionization energy increases for each successive electron 2. Each electron removed experiences a stronger effective nuclear charge (more positive charges than negative charges) 3. The greatest increase in ionization energy comes when trying to remove an electron from a stable, noble gas configuration (complete octet in the outer energy level) III. IV. Trends in Ionic Size A. Cations 1. Positive ions 2. Always Smaller than the corresponding atom a. Protons outnumber electrons b. Less shielding of electrons B. Anions 1. Negative ions 2. Always Larger than the corresponding atoms a. Electrons outnumber protons b. Greater electron-­‐electron repulsion C. Trends 1. Ion size tends to increase downward within a group 2. Cationic and anionic radii decrease across the period Trends in Electronegativity

A. Electronegativity 1. A measure of the ability of an atom in a chemical compound to attract electrons 2. Elements that do not form compounds are not assigned electronegativities B. Trends 1. Nonmetals have characteristically high electronegativity a. Highest in the upper right corner (highest = Fluorine 4.0) 2. Metals have characteristically low electronegativity a. Lowest in the lower left corner of the table (Fr and Cs 0.7) 3. Electronegativity tends to increase across a period 4. Electronegativity tends to decrease down a group of main-­‐group elements V. Valence Electrons A. Electrons in the outer energy level (s & p blocks) 1. Electrons available to be lost, gained, or shared in the formation of chemical compounds 2. Chemical compounds form because electrons are lost, gained, or shared between atoms 3. Electrons that interact are those in the highest energy levels (s & p blocks) VI. Sample Problems 1.) Identify the group, period, and block in which the element that has theelectron configuration [Xe]6s 2 is located.2.) Write the electron configuration for the Group 1 element in the third period.Is this element likely to be more reactive or less reactive than the elementdescribed in (a)?3.) An element has the electron configuration [Kr]4d 5 5s 1 . Identify the period,block, and group in which this element is located. Then, consult the periodictable to identify this element and the others in its group.4.) Without looking at the periodic table, write the outer electron configuration forthe Group 14 element in the second period. Then, name the element, andidentify it as a metal, nonmetal, or metalloid.5.) Name the block and group in which each of the following elements is locatedin the periodic table. Then, use the periodic table to name each element.Identify each element as a metal, nonmetal, or metalloid. Finally, describewhether each element has high reactivity or low reactivity.a. [Xe]4f 14 5d 9 6s 1 c. [Ne]3s 2 3p 6b. [Ne]3s 2 3p 5 d. [Xe]4f 6 6s 26.) Of the elements magnesium, Mg, chlorine, Cl, sodium, Na, and phosphorus,P, which has the largest atomic radius? Explain your answer in terms oftrends of the periodic table.7.) Consider two main-group elements, A and B. Element A has a first ionizationenergy of 419 kJ/mol. Element B has a first ionization energy of 1000 kJ/mol.

Decide if each element is more likely to be in the s block or p block. Whichelement is more likely to form a positive ion?8.) Of the elements gallium, Ga, bromine, Br, and calcium, Ca, which has thehighest electronegativity? Explain your answer in terms of periodic trends.Solutions to notes problems above: 1.) The element is in Group 2, as indicated by the group configuration of ns 2 .It is in the sixth period, as indicated by the highest principal quantum number in itsconfiguration, 6.The element is in the s block.2,) In a third-period element, the highest occupied energy level is the third main energylevel, n = 3. The 1s, 2s, and 2p sublevels are completely filled.This element has the following configuration:1s 2 2s 2 2p 6 3s 1 or [Ne]3s 1Because it is in Group 1, this element is likely to be more reactive than the elementdescribed in (a), which is in Group 2.3.) The number of the highest occupied energy level is 5, so the element is in the fifthperiod.There are five electrons in the d sublevel, which means that it is incompletely filled.The d sublevel can hold 10 electrons. Therefore, the element is in the d block.For d-block elements, the number of electrons in the ns sublevel (1) plus the number ofelectrons in the (n - 1)d sublevel (5) equals the group number, 6.This Group 6 element is molybdenum. The others in Group 6 are chromium, tungsten,and seaborgium.4.) The group number is higher than 12, so the element is in the p block.The total number of electrons in the highest occupied s and p sublevels is thereforeequal to the group number minus 10 (14 - 10 = 4).Two electrons are in the s sublevel, so two electrons must also be present in the 2psublevel.The outer electron configuration is 2s 2 2p 2 .The element is carbon, C, which is a nonmetal.5.)a. The 4f sublevel is filled with 14 electrons. The 5d sublevel is partially filled with nineelectrons. Therefore, this element is in the d block. The element is the transition metalplatinum, Pt, which is in Group 10 and has a low reactivity.b. The incompletely filled p sublevel shows that this element is in the p block. A total ofseven electrons are in the s and p sublevels, so this element is in Group 17, thehalogens. The element is chlorine, Cl, and is highly reactive.c. This element has a noble-gas configuration and thus is in Group 18 in the p block.The element is argon, Ar, which is an unreactive nonmetal and a noble gas.d. The incomplete 4f sublevel shows that the element is in the f block and is alanthanide. Group numbers are not assigned to the f block. The element is samarium,Sm. All of the lanthanides are reactive metals.

6.) Sodium has the largest atomic radius. All of the elements are in the third period. Ofthe four, sodium has the lowest atomic number and is the first element in the period.Atomic radii decrease across a period.7.) Element A has a very low ionization energy, which means that atoms of A loseelectrons easily.Element A is most likely to be an s-block metal because ionization energies increaseacross the periods.Element B has a very high ionization energy which means that atoms of B havedifficulty losing electrons.Element B would most likely lie at the end of a period in the p block.Element A is more likely to form a positive ion because it hasa much lower ionization energy than element B does.8.) All of these elements are in the fourth period. Bromine has the highest atomicnumber and is farthest to the right in the period. Bromine should have the highestelectronegativity because electronegativity increases across the periods.

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