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The periodic law Chapter 5 1 Why do we need a table? To organize the elements To show trends 2 Periodic A repeating pattern 3 Mendeleev’s table 1869 – Dmitri Mendeleev – Russian Arranged the elements in order of increasing mass and noticed that chemical properties were periodic Put the elements into groups according to properties 4 Mendeleev vs. Meyer 1860s Mendeleev and German Lothar Meyer each made an eight column table. Mendeleev left some blanks in his table in order for all the columns to have similar properties – he predicted elements that hadn’t been discovered yet. 5 Why similar properties? Why did they group according to properties and mass and not atomic number or number of outer level electrons? 6 Germanium Mendeleev’s predictions Atomic mass = 72 Atomic mass = 72.60 High melting point Melting point = 958 °C Density = 5.5 g/cm3 Dark gray metal Actual element Density = 5.36 g/cm3 Gray metal Mendeleev’s blank spots and his ability to predict future elements helped his table win acceptance. 7 Mendeleev’s table Elements arranged in order of increasing mass. Properties are repeated in an orderly, periodic, fashion. Mendeleev’s periodic law – the properties of the elements are a periodic function of their masses. 8 9 Mass mistakes? In order for Mendeleev to arrange his elements by properties, he had to put tellurium and iodine in the wrong order. He explained this by assuming that their masses hadn’t been measured very accurately. 10 More mass mistakes? Nickel and cobalt Argon and potassium Better mass measurements just confirmed the discrepancy 11 Explanation 1913 – Henry Moseley X-ray experiments revealed the atomic number was the number of protons Modern periodic law – the properties of the elements are a periodic function of their atomic numbers 12 Modern periodic table An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group. 13 Noble gases Not discovered on Earth until 1894 1900. Group 18 was added to the table 14 Lanthanides Hard to separate All have similar properties Added to the table in the early 1900s 15 Actinides Discovered later Also all have similar properties 16 Periodicity Elements in the same group (column) have similar properties. 17 Chemical properties of an element Are governed by the electron configuration of an atom’s highest energy level 18 Period length Determined by the number of electrons than can occupy the sublevels being filled in that period. Table 5-1 19 Full periodic table Table with f-block in place 20 1st period 1s sublevel being filled 1s can hold 2 electrons, so there are 2 elements in the 1st period. 21 2nd and 3rd periods 2s and 2p or 3s and 3p being filled s and p sublevels can hold 8 total, so there are eight elements in these periods 22 4th and 5th periods Add d sublevels, which can hold 10 electrons Need to fill 4s, 3d, and 4p – 18 electrons 18 elements in each period 23 6th and 7th periods Add f-block, which holds 14 electrons Fill 6s, 5d, 4f, 6p Need 32 electrons 32 elements in each period 24 Figure 5-5 Shows blocks 25 Electron configurations Elements in columns 1, 2, and 1318 have their last electron added in an s or p orbital. Elements in columns 3-12 have their last electron added in a d level. 26 The s-block elements: Groups 1 and 2 Chemically reactive metals Group 1 Have 1 electron in outer s orbital Coefficient represents period Row 2: 2s1, Row 3: 3s1, etc. (ns1) Group 2 Have 2 electrons in outer s orbital Coefficient represents period Row 2: 2s2, Row 3: 3s2, etc. (ns2) 27 Alkali metals Metals in group 1 Have silvery appearance Soft enough to cut with a knife Not found alone in nature React violently with nonmetals Melting point decreases as you go down the table 28 Alkaline-earth metals Group 2 Harder, denser, and stronger than alkali metals Higher melting points than alkalis Less reactive Not found alone in nature 29 Hydrogen and helium Hydrogen Located above group 1 because of its electron configuration Not really in group 1, because its properties don’t match Helium Has an electron configuration like group 2 elements In group 18 because it is unreactive 30 Discuss Page 133 Sample problem 5-1 and practice problems 31 Discuss Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [Rn] 7s1 is located. Group 1, 7th period, s block Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [He]2s2 is located. Group 2, second period, s block 32 d-block elements: Groups 3-12 End in d1 to d10. Coefficients are one less than the period Example: Fe is in the 6th column of transition elements in the 4th period, ends in 3d6 33 Transition elements Groups 3-12 Typical metallic properties Good conductors High luster Less reactive than alkalis and alkaline-earths Some are unreactive enough to appear in nature 34 p-block elements: groups 13-18 End in p1 to p6. Coefficients are the same as the period ns2np1 Always have a full s-sublevel 35 p-block elements Properties vary greatly Includes all nonmetals except hydrogen and helium Includes all the metalloids Solids, liquids and gases Between metals and nonmetals Brittle solids Semiconductors – can conduct under certain conditions Includes some metals Less reactive than alkalis and alkaline-earths 36 Halogens Group 17 Most reactive nonmetals Form compounds called salts 37 f-block elements Lanthanides and actinides Endings are f1 to f14 Coefficients are two less than the period All actinides are radioactive Those after neptunium are synthetic 38 Discuss Sample problems and practice problems on pages 136, 138, and 139 With your group first, then join with another group. Do you have any questions? 39 Atomic radius Ideally, the distance from the center of the atom to the edge of it’s orbital. But, atoms are “fuzzy”, not clearly defined. Defined as one-half the distance between the nuclei of identical atoms that are bonded together. 40 Period trends – see figure 5-13 As we move from left to right across the table, we gain protons. There is a greater positive charge on the nucleus. This greater charge pulls harder on the outer electrons, pulling them in closer. The atom gets smaller. 41 Group trends As we move down the table, the principle quantum number increases. When the principle quantum number increases, the electron cloud gets bigger. The size of the atoms gets bigger. 42 Discuss Which of the elements Li, Rb, K, and Na has the smallest atomic radius? Why? Li, it is highest on the table Which of the elements Zr, Rb, Mo, and Ru has the largest atomic radius? Why? Rb, it is farthest to the left on the table 43 Ion An atom or group of bonded atoms that has a positive or negative charge 44 Ionization Any process that makes ions 45 Ionization energy (IE) First ionization energy (IE1) – the energy required to remove the most loosely held electron. Measured in kJ/mol 46 Ionization energy – see figure 5-15 Experimentally determined. Tends to increase as you move across a row from left to right From isolated atoms in the gas phase Why group 1 is most reactive Caused by higher charge Tends to decrease as you move down a column Electrons are farther from nucleus Shielding from inner electrons 47 Other Ionization Energies – see Table 5-3 Energy required to remove other electrons from positive ions. IE2, IE3, etc Get higher as you remove more electrons Less shielding 48 Noble Gases Have High ionization energies When a positive ion of another element reaches a noble gas configuration, its ionization energy goes up. Example: When K loses one electron, it has Ar’s electron configuration This makes it stable Its IE2 is much higher than its IE1 49 Discuss State in words the general trends in ionization energies down a group and across a period of the periodic table. 50 Electron affinity The energy change that occurs when an electron is gained by a neutral atom Most atoms release energy Represented by a negative number Some atoms gain energy Represented by a positive number These ions will be unstable KJ/mol 51 Period trends – see figure 5-17 Group 17 has most negative electron affinity. Tends to get more negative (release more energy) as we move to the right Exceptions: groups with full or half-full sublevels are more stable 52 Group trends Not as regular Usually, electrons add with greater difficulty as we move down 53 Adding additional electrons Second electron affinities are all positive because it is more difficult to add electrons to a negative ion. If a noble gas configuration has been reached, it is even more difficult. 54 Discuss State in words the general trends in electron affinities down a group and across a period of the periodic table. 55 Ionic Radii Cation – a positive ion Ionic radius smaller than atomic radius Anion – a negative ion Ionic radius is larger 56 Period Trends – see figure 5-19 Metals form cations by losing electrons Ions are smaller Radius decreases as we move across Nonmetals form anions by gaining electrons Ions are larger Radius decreases as we move across 57 Group trends Ionic radius increases as you go down the table 58 Valence electrons Available to be lost, gained or shared in the formation of chemical compounds In highest energy levels For s-block, the group number is the same as the number of valence electrons For the p-block, the group number is 10 more than the number of valence electrons 59 Electronegativity The measure of the ability of an atom in a compound to attract electrons The atom with higher electronegativity pulls the electrons closer to itself 60 Electronegativity trends (figure 5-20) Increases left to right across the rows Decreases down the columns 61 Discuss Explain why elements with high (more negative) electron affinities are also the most electronegative. 62 d- and f-block elements Properties vary less and with less regularity than others Atomic radii d-block Usual patterns f-block (unusual) Increase across periods Decrease down groups 63 d- and f-block elements Ionization energy Ionic radii Increase across periods d-block increases down groups (unusual) f-block decreases down groups Cations have smaller radii Electronegativity d-block follows normal rules f-block all have similar electronegativities 64 Discuss Among the main-group elements, what is the relationship between group number and the number of valence electrons? In general, how do the periodic properties of the d-block elements compare with those of the maingroup elements? 65 Prelab notes Precipitate – solid that falls out of a solution The formation of a precipitate indicates there has been a chemical change. This means that there were ions present that were free to react. 66