2.
History of the Periodic Table <ul><li>Prior to 1860, 60 elements had been discovered, but there was no method for method for determining the atomic mass or the number of atoms of an element in a compound. </li></ul><ul><li>Chemists were using different masses for the same element, which threw off chemical compounds. </li></ul><ul><li>It was virtually impossible to understand each other’s results. </li></ul>
3.
<ul><li>In Sept. 1860, a group of chemists assembled at the First International Congress of Chemists in Germany to discuss these issues. </li></ul><ul><li>At this time, Stanislao Cannizzarro presented a method for measuring the masses of atoms, which allowed all chemists to agree on values and begin to find the relationships between elements. </li></ul>
4.
Mendeleev <ul><li>Dmitri Mendeleev (Russian) heard about the Congress and incorporated the new values in a chemistry text he was writing. </li></ul><ul><li>He wanted to arrange the elements by their properties. </li></ul><ul><li>So, he wrote each element and its atomic mass and observable properties on cards. </li></ul><ul><li>Then, he began to arrange the cards according their properties, looking for patterns. </li></ul>
5.
What he found… <ul><li>He noticed that when elements were arranged in order of increasing atomic masses, certain similarities appeared a regular intervals. </li></ul><ul><li>This is know as periodic. </li></ul>
6.
The first table… <ul><li>His first table was published in 1869. </li></ul><ul><li>The procedure that Mendeleev used to create his table left many blank spaces. </li></ul><ul><li>In 1871, he predicted the existence and properties of the elements that would fill three of the spaces of his table. </li></ul><ul><li>By 1886, all three elements were discovered – scandium, Sc, gallium, Ga, and germanium, Ge. </li></ul><ul><li>Their properties were very similar to what was predicted! </li></ul>
7.
The end of Mendeleev’s story… <ul><li>Because of the success of his predictions, most chemists were persuaded to accept his periodic table. </li></ul><ul><li>He has the credit for discovering the periodic law. </li></ul><ul><li>Still some questions remained… </li></ul>
8.
Moseley <ul><li>In 1911 (40 years later), Henry Moseley (England) was working alongside Rutherford. </li></ul><ul><li>He noticed during some of his research that elements fit into a better pattern when they were arranged in increasing order according to their atomic number. </li></ul><ul><li>His work led us to today’s atomic number as the basis for organizing the periodic table. </li></ul>
9.
Periodic Law <ul><li>The physical and chemical properties of the elements are periodic functions of their atomic numbers. </li></ul><ul><li>In other words, </li></ul><ul><li>When the elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals. </li></ul>
10.
The Modern Periodic Table <ul><li>Since Mendeleev’s periodic table, chemists have: </li></ul><ul><ul><li>Discovered new elements </li></ul></ul><ul><ul><li>Synthesized new elements in the lab </li></ul></ul><ul><ul><li>Placed all in the periodic table, grouping them with elements with similar properties </li></ul></ul>
11.
The Periodic Table <ul><li>An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group. </li></ul>
12.
The Noble Gases <ul><li>John William Strutt (England) and Sir William Ramsay (Scotland) discovered argon, Ar. </li></ul><ul><li>This gas, found in the atmosphere, had been overlooked because of its total lack of chemical reactivity. </li></ul><ul><li>Prior to this, helium, He, had been discovered in the atmosphere, as well. </li></ul><ul><li>Ramsay proposed a new group so that argon and helium would fit into the table, and it fit between Groups 17 and 1. </li></ul><ul><li>Later, he discovered krypton, Kr, and xenon, Xe, and another chemist discovered the final gas, radon, Rn. </li></ul>
13.
The Lanthanides and The Actinides <ul><li>The lanthanides are the 14 elements with atomic numbers 58 (cerium, Ce) to 71 (lutetium, Lu). </li></ul><ul><li>The actinides are the 14 elements with atomic numbers from 90 (thorium, Th) to 103 (lawrencium, Lr). </li></ul><ul><li>The lanthanides and actinides belong to Periods 6 and 7, respectively and between Groups 3 and 4; however, to save space, they are set off below the main portion of the table. </li></ul>
14.
Periodicity <ul><li>Periodicity, or a repeating pattern, with respect to atomic number can be observed in any group of elements in the periodic table. </li></ul><ul><li>Figure 4 (p. 136) </li></ul>
15.
Electron Configurations and the Periodic Table Generally, the electron configuration of an atom’s highest occupied energy level governs that atom’s chemical properties!
16.
Periods and Blocks of the Periodic Table <ul><li>Elements are arranged not only arranged in groups, but also in horizontal rows, called periods. </li></ul><ul><li>The length of a period depends on the number of electrons that can occupy the sublevels being filled in that period. </li></ul>
17.
<ul><li>In the 1 st period, the 1s sublevel is being filled, which can hold a total of 2 electrons. </li></ul><ul><li>In the 2 nd period, the 2s sublevel, which can hold 2 electrons, AND the 2p sublevel, which can hold 6 electrons, are being filled. </li></ul><ul><li>In the 3 rd period, the 3s, 3p and 3d sublevels are filled, and this pattern continues. </li></ul><ul><li>Basically, the period of an element can be determined from the element’s electron configuration. </li></ul>
18.
Example <ul><li>Arsenic </li></ul><ul><li>[Ar]3d 10 4s 2 4p 3 </li></ul><ul><li>The 4 in 4p 3 indicates that arsenic’s highest occupied energy level is the 4 th level. </li></ul><ul><li>Therefore, Arsenic is found in the 4 th period!! </li></ul>
19.
<ul><li>Based on the electron configurations, the periodic table can be divided into 4 blocks, the s, p, d, and f blocks. </li></ul><ul><li>This is determined by whether an s, p, d, or f sublevel is being filled in elements of that block. </li></ul>
20.
The s-Block Elements: Groups 1 and 2 <ul><li>The elements of the s block are chemically reactive metals. </li></ul><ul><ul><li>Group 1 is more reactive than Group 2. </li></ul></ul><ul><ul><li>The outermost energy level of Group 1 elements has 1 electron and Group 2 has 2 electrons. </li></ul></ul>
21.
Alkali Metals <ul><li>The elements of Group 1 </li></ul><ul><li>Lithium </li></ul><ul><li>Sodium </li></ul><ul><li>Potassium </li></ul><ul><li>Rubidium </li></ul><ul><li>Cesium </li></ul><ul><li>Francium </li></ul>
22.
Alkali Metals <ul><li>Silvery appearance </li></ul><ul><li>Soft (can be cut w/ knife) </li></ul><ul><li>Not found in nature as a free element </li></ul><ul><li>Combine with most nonmetals </li></ul>
23.
Alkaline-Earth Metals <ul><li>The elements of Group 2 </li></ul><ul><li>Beryllium </li></ul><ul><li>Magnesium </li></ul><ul><li>Calcium </li></ul><ul><li>Strontium </li></ul><ul><li>Barium </li></ul><ul><li>Radium </li></ul>
24.
Alkaline-Earth Metals <ul><li>Harder, denser, stronger than alkali metals </li></ul><ul><li>Higher melting points </li></ul><ul><li>Also, too reactive to be found in nature </li></ul>
25.
Hydrogen and Helium <ul><li>Hydrogen </li></ul><ul><ul><li>Has a configuration of 1s 1 </li></ul></ul><ul><ul><li>Doesn’t share properties of Group 1 elements </li></ul></ul><ul><ul><li>Unique element </li></ul></ul><ul><ul><li>Doesn’t resemble any group </li></ul></ul><ul><li>Helium </li></ul><ul><ul><li>Has a configuration of 1s 2 </li></ul></ul><ul><ul><li>Doesn’t have 8 electrons in outer level </li></ul></ul><ul><ul><li>However, it is stable and unreactive </li></ul></ul><ul><ul><li>Placed in Group 18 </li></ul></ul>
26.
The d Block Elements: Groups 3-12 <ul><li>d orbitals do not appear until the 3 rd energy level. Therefore, d Block elements don’t fit in until the 3 rd period! </li></ul><ul><li>Elements in the d-block Groups 4-11 have some deviations and do not necessarily have identical outer electron configurations. </li></ul>
27.
Example <ul><li>Group 10 </li></ul><ul><ul><li>Ni = [Ar]3d 8 4s 2 </li></ul></ul><ul><ul><li>Pd = [Kr]4d 10 5s 0 </li></ul></ul><ul><ul><li>Pt = [Xe]4f 14 5d 9 6s 1 </li></ul></ul><ul><ul><li>Notice, though, that in each case the sum of the outer s and d electrons is equal to the group number! </li></ul></ul>
28.
Transition Elements <ul><li>d-Block elements are metals with typical metallic properties </li></ul>
29.
Transition Elements <ul><li>Good conductors of electricity </li></ul><ul><li>High luster </li></ul><ul><li>Less reactive </li></ul><ul><li>Some don’t form compounds easily </li></ul><ul><li>Some exist in nature </li></ul>
30.
The p-Block Elements: Groups 13-18 <ul><li>The p-block elements consist of all the elements in Groups 13-18, except helium. </li></ul><ul><li>All elements in these groups will have a full s orbital and a p orbital. </li></ul><ul><ul><li>Group 13 = s 2 p 1 </li></ul></ul><ul><ul><li>Group 14 = s 2 p 2 </li></ul></ul><ul><ul><li>And so on… </li></ul></ul><ul><ul><li>Group 18 = s 2 p 6 </li></ul></ul>
31.
<ul><li>For p-block atoms, the total number of electrons in the highest occupied level is equal to the group number minus 10. </li></ul><ul><li>Example </li></ul><ul><ul><li>Bromine Group 17 </li></ul></ul><ul><ul><li>17-10 = 7 electrons in its highest energy level </li></ul></ul><ul><ul><li>[Ar]3d 10 4s 2 4p 5 </li></ul></ul>
32.
<ul><li>p-Block includes: </li></ul><ul><ul><li>all of the nonmetals, except hydrogen and helium </li></ul></ul><ul><ul><li>All six of the metalloids (boron, silicon, germanium, arsenic, antimony, and tellurium) </li></ul></ul><ul><ul><li>Eight metals </li></ul></ul><ul><li>Group 17 elements are the halogens and are the most reactive nonmetals. </li></ul>
33.
The f-Block Elements: Lanthanides and Actinides <ul><li>f-Block elements are wedged between Groups 3 and 4 in the 6 th and 7 th periods. </li></ul>
34.
<ul><li>Lanthanides </li></ul><ul><ul><li>Shiny metals </li></ul></ul><ul><ul><li>Similar in reactivity to Group 2 </li></ul></ul><ul><ul><li>4f orbitals are being filled </li></ul></ul><ul><li>Actinides </li></ul><ul><ul><li>All radioactive </li></ul></ul><ul><ul><li>5f orbitals are being filled </li></ul></ul>
35.
Atomic Radii May be defined as ½ the distance between the nuclei of identical atoms that are bonded together.
36.
Period Trends <ul><li>The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus. </li></ul><ul><li>Electrons add to s and p sublevels in the same level, and they are gradually pulled closer to the nucleus. </li></ul><ul><li>Increase pull = Decrease in radii </li></ul><ul><li>The difference in radii between neighboring atoms in each period grows smaller. </li></ul>
37.
Group Trends <ul><li>The radii of the elements increase as you go down the group. </li></ul><ul><li>As electrons occupy sublevels in successively higher main energy levels located farther from the nucleus, the sizes of the atoms increase. </li></ul><ul><li>In general, the atomic radii of the main-group elements increase down a group. </li></ul>
38.
Ionization Energy, IE The energy required to remove one electron from a neutral atom of an element.
39.
Ionization Energy <ul><li>An electron can be removed from an atom if enough energy is supplied. </li></ul><ul><li>A + energy A + + e - </li></ul><ul><li>A + represents an ion of element A with a single positive charge, which is referred to as a 1 + ion. </li></ul>
40.
Period Trends <ul><li>In general, ionization energies of the main-group elements increase across each period. </li></ul><ul><ul><li>Group 1 has the lowest ionization energies in their respective periods; so, they lose electrons most easily, which is the reason for their high reactivity. </li></ul></ul><ul><ul><li>Group 18 has the highest ionization energies, and they do not lose electrons easily, which is the reason for their low reactivity. </li></ul></ul>
41.
<ul><li>The increase is caused by increasing nuclear charge. A higher charge more strongly attracts electrons. </li></ul><ul><li>Increasing nuclear charge is responsible for both increasing ionization energy and decreasing radii across the periods. </li></ul>
42.
Group Trends <ul><li>Among the main-group elements, ionization energies generally decrease down the group. </li></ul><ul><li>Electrons are in higher energy levels and are removed more easily. </li></ul>
43.
Electron Affinity The energy change that occurs when an electron is acquired by a neutral atom.
44.
Electron Affinity <ul><li>Neutral atoms can acquire electrons. </li></ul><ul><li>Most atoms release energy when they acquire an electron </li></ul><ul><li>A + e - A - + energy </li></ul><ul><li>Some atoms must be “forced” to gain an electron by the addition of energy. </li></ul><ul><li>A + e - + energy A - </li></ul><ul><li>An ion produced in this way will be unstable and will lose the added electron spontaneously. </li></ul>
45.
Period Trends <ul><li>Group 17 elements gain electrons most readily, which is indicated by the large negative values of halogens’ electron affinities. This accounts for their high reactivities. </li></ul><ul><li>In general, as electrons add to the same sublevel of atoms, electron affinities become more negative. </li></ul>
46.
Group Trends <ul><li>Electrons add with greater difficulty down a group because </li></ul><ul><ul><li>There is a slight increase in effective nuclear charge down a group, which decreases electron affinities. </li></ul></ul><ul><ul><li>An increase in atomic radius down a group, which decreases electron affinities. </li></ul></ul>
48.
Ionic Radii <ul><li>A positive ion is known as a cation. The formation always leads to an increase in atomic radius. </li></ul><ul><li>A negative ion is known as a anion. The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius. </li></ul>
49.
Period Trends <ul><li>The metals at the left tend to form cations and the nonmetals at the upper right tend to form anions. </li></ul><ul><li>Cationic radii decrease across a period because the electron cloud shrinks. </li></ul><ul><li>Anionic radii decrease across each period for the elements in Groups 15-18. </li></ul>
50.
Group Trends <ul><li>The outer electrons in both cations and anions are in higher energy levels as one reads down a group. </li></ul><ul><li>Therefore, just as there is a gradual increase of atomic radii down a group, there is also a gradual increase of ionic radii. </li></ul>
51.
Valence Electrons <ul><li>The electrons available to be lost, gained, or shared in the formation of chemical compounds. </li></ul><ul><li>Often located in incompletely filled main-energy levels. </li></ul><ul><li>Example </li></ul><ul><ul><li>Electron in the 3s sublevel of Na is lost to form Na + </li></ul></ul>
52.
<ul><li>Group 1 has 1 valence electron and Group 2 has 2 valence electrons. </li></ul><ul><li>Groups 13-18 – To find the valence electrons, Group # - 10 = valence electrons. </li></ul>
53.
Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound
54.
Period Trends <ul><li>Electronegativities tend to increase across each period, although there are exceptions. </li></ul><ul><li>The alkali and alkaline-earth metals are the lease electronegative elements. </li></ul><ul><li>Nitrogen, oxygen, and the halogens are the most electronegative elements – their atoms attract electrons strongly in compounds. </li></ul>
55.
Group Trends <ul><li>Electronegativities tend to either decrease down a group or remain the same. </li></ul><ul><li>The combination of the period and group trends in electronegativity results in the highest values belonging to the elements in the upper right of the periodic table and the lowest belong to the elements in the lower left. </li></ul>