6.1 - The Periodic Table: A History

6.1 - The Periodic Table: A History

History of the Periodic Table
Twelve elements have been known since ancient times. What do you think they are? (Name them, use your periodic table to help you.) carbon, sulfur, iron, copper, arsenic, silver, tin, antimony, gold, mercury, lead, bismuth

Why do you think these particular elements have been known for so long, while most elements were not discovered until the 1800s and 1900s? All are metals or metalloids, all but one are solids. They are the less reactive metals, therefore were found in their native form. Other common metals, such as sodium and calcium are highly reactive, so are never found in their native form. Gases were not recognized as matter.

Overview of the Periodic Table
Metals Metalloids Nonmetals Noble gases

Overview of the Periodic Table
Metals Metalloids Nonmetals Noble gases 1. excellent heat conductor 2. excellent electrical conductor 3. lustrous (shiny) 4. malleable, ductile 5. silvery-gray, except Cu and Au 6. solids at room T, except Hg Some properties of metals, some properties of nonmetals 1. moderate electrical conductivity 2. appearance – more like metals – lustrous, silvery-gray 3. brittle like nonmetals 4. solids at room T poor heat conductors poor electrical conductors not lustrous brittle variety of colors gases or brittle solids at room T extremely unreactive – “inert” rarely form compounds with other elements colorless, odorless gases at room T

Jöns Jakob Berzelius 1828 Swedish chemist - developed a table of
atomic weights - introduced letters to symbolize elements made the task easier 33 elements known by 1800

Jöns Jakob Berzelius

Johann Döbereiner 1829 German Chemist Triads 53 known elements
Sets of three elements with similar properties: Cl, Br, I…Ca, Ba, Sr…S, Se, Te 8

Döbereiner a) - described triads of elements (e.g. Cl, Br, I; Ca, Ba, Sr; Li, Na, K) - first indication that elements are related to one another - atomic mass is related to chemical properties – the mass of the center element was halfway between the masses of the other two elements, all three have similar properties

elements

(big Chemistry Conference)
1860 Karlsruhe Congress (big Chemistry Conference) Germany

John Newlands 1865 English Chemist Arranged elements by atomic mass
Described the “Rule of octaves” 62 elements Every 8th element has similar properties, Musical reference considered non-scientific 13

Lothar Meyer 1870 German Chemist
Arranged elements based on atomic mass Discovered periodic properties related to atomic volume Established concept of valency Did not know mendeleev, 15

Meyer’s Data

It’s in the Cards Pre-Lab
Ionization energy = the amount of energy, in J or kJ, required to remove 1 electron from an atom in the gaseous state Atomic radius = the distance between the nuclei of two adjacent atoms of the same kind, divided by 2, measured in pm Melting point = the temperature at which a solid becomes a liquid, measured in oC

It’s in the Cards Pre-Lab
Average atomic mass = the weighted average of the masses of all known isotopes of an element, measured in amu (or g) Density = ratio of mass divided by volume, g/mL or g/cm3 Electronegativity = a measure of the relative ability of an atom to attract electrons in the context of a chemical bond, Paulings or none

Dmitri Mendeleev

Dmitri Mendeleev 1869 Russian chemist
Wrote elements and properties on notecards Arranged by atomic mass and properties Noted repetition of properties every 8 or 18 elements Left gaps in the table,. Put elements with similar properties in horizontal rows. 20

Dmitri Mendeleev 1869 Predicted properties of 3 elements!
eka-aluminum, eka-boron, eka-silicon Problems: Ar/K, Te/I, Co/Ni First element of each pair has greater atomic mass Ga, Sc, Ge Daltons model – protons not known!! 21

Properties of Some Elements Predicted by Mendeleev
Predicted Element Element and year discovered Properties Predicted Properties Observed Properties Eka-aluminum Gallium, 1875 Density of metal 6.0 g/mL 5.96 g/mL Melting point Low 30oC Oxide formula Ea2O3 Ga2O3

Predicted Element Element and year discovered Properties Predicted Properties Observed Properties Eka-boron Scandium, 1877 Density of metal 3.5 g/mL 3.86 g/mL Oxide formula Eb2O3 Sc2O3 Solubility of oxide dissolves in acid

Predicted Element Element and year discovered Properties Predicted Properties Observed Properties Eka-silicon Germanium, 1886 Melting point High 900oC Density of metal 5.5 g/mL 5.47 g/mL Color of metal Dark gray Grayish white Oxide formula EsO2 GeO2 Density of oxide 4.7 g/mL 4.70 g/mL Chloride formula EsCl4 GeCl4

Review Döbereiner 1829 Newlands 1865 Meyer 1870 Mendeleev
Arranged by atomic mass Triads: [Cl Br I], [Ca Ba Sr], [Li Na K] Newlands 1865 Rule of Octaves Meyer 1870 Arranged by atomic mass, periodic trend with atomic volume Established concept of valency Mendeleev Repetition every 8 or 18 elements Predicted 3 elements not yet discovered: eka-aluminium - gallium, eka-silicon - germanium and eka-boron - scandium Triads Published a table a few years prior to Mendeleev Octaves Periodic trends with gaps, made accurate predictions 28

Discovery of the Noble Gases 1890s
Lord Rayleigh (physicist) and Sir William Ramsay (chemist) Argon “the lazy one”, discovered when Ramsay was trying to isolate nitrogen Helium – found on earth in uranium minerals (found in the sun in 1868) Neon “the new one” Krypton “the hidden one” Xenon “the alien one” 1910 – Radon Properties: Largely unreactive 8 electrons in valence shell Low boiling and melting points Used as cryogenic refrigerents Xe short arc lamps used in IMAX projectors 29

Nucleus discovered – 1910 Rutherford predicted that the charge of an atom is proportional to its mass

Henry Moseley 1913 English Physicist
worked with Rutherford – was given the task of testing his prediction about charge vs. mass Periodic Law: Properties of elements are periodic functions of their atomic numbers 31

Ön of emitted X-rays corresponded to # protons  atomic number “Do other properties match atomic numbers?” Yes! arranged the periodic table by atomic #’s, not mass

Law of Atomic Numbers - the properties of elements are periodic functions of their atomic numbers (not atomic mass)  corrected incorrect placement of cobalt and nickel, and iodine and tellurium

Glenn Seaborg 1940’s American Scientist at UC Berkeley
Nobel Prize in Physics, 1951 Discovered 7 elements beyond U Developed actinide series and added it to PT Seaborgium the only element publicly named after a living person plutonium, americium, curium, berkelium, californium, mendelevium, nobelium Seaborgium named 1974, he Died 1999 Nobel prize winner Advisor to 10 presidents From truman to clinton Developed 100 isotopes including Iodine-131, used to treat thyroid disease Joined the Manhattan Project - advocated nuclear disarmament 34

Seaborgium Lawrencium Berkelium, Californium Americium
Letter to Seaborg Seaborgium Lawrencium Berkelium, Californium Americium 36

The Newest Elements Riken Institute, Japan (1st time for Asia)
113 Zinc-70 + Bismuth (2012) US-Russia collaborative (Dubna, Lawrence Livermore): 115 Calcium-48 + Americium (2003) 117 Calcium-48 + Berkelium (2009) 118 Calcium-38 + Californium (2002)

Warmup Name the four scientists and the important scientific meeting we talked about previously Write them down in chronological order, clearly indicating who came before and who came after the scientific meeting Use a couple of words or a phrase to remind yourself of their contribution to the history of the periodic table, to make a connection you will remember

Trends of the Periodic Table “periodic” = repeating pattern
Overall theme = electrons’ positions relative to each other and the nucleus determine the following properties.

6.3 - Periodic Trends Trends are most pronounced with the representative elements 40

Trends of the Periodic Table “periodic” = repeating pattern
1. Electron configuration (  reactivity and bonding) Atomic radius Ionic radius 4. Ionization energy 5. Electronegativity

1. Electron Configuration
Compare the charges on the ion list with the position of the element in the periodic table

Electron Configuration
Noble gas configuration = [core] e-’s ‘Outer’ electrons = valence e-’s Elements of groups 1A-8A have valence e-’s in s and p orbitals

Periodic Trends The position of a valence electron and the ability to remove it from an atom are related to the number of protons in the nucleus the extent to which the valence electron is shielded from the positively-charged nucleus by the negatively-charged core electrons

Isoelectronic Series = a group of ions and atoms that have the same electron configuration 1. Draw the electron configuration of each of the following elements. 2. What ions will they form? 3. When ions, how many electrons does each have? How many protons? 4. Predict the relative diameters of the members of this isoelectronic series.

Isoelectronic Series O F Ne Na Mg Element Electron config Ion # e-’s
# p+ O F Ne Na Mg Prediction: smallest to largest:

Isoelectronic Series O F Ne Na Mg Element Electron config Ion # e-’s
# p+ O 1s22s22p4  1s22s22p6 O2- 10 e- 8 p+ F 1s22s22p5 F- 9 p+ Ne 1s22s22p6 10 p+ Na 1s22s22p63s1 Na+ 11 p+ Mg 1s22s22p63s2 Mg2+ 12 p+ Prediction: smallest to largest: Mg2+ < Na+ < Ne < F-< O2-

The Periodic Table

The Periodic Table How many elements are listed in the periodic table? (the one Dr. Hart gave you…) __________ What is the atomic number of selenium? _________ What is the symbol for palladium? _________ What is the atomic mass of strontium? ________ How are elements that are gases at room temperature designated in the periodic table? _________________ How many columns of elements does the periodic table contain? ______ What is another name for a column of elements? __________ 8. What two group numbers can be used to designate elements in the second column of the periodic table? _________

The Periodic Table How many elements are listed in the periodic table? (the one Dr. Hart gave you…) ___118_______ What is the atomic number of selenium? __34____ What is the symbol for palladium? ___Pd______ What is the atomic mass of strontium? ___87.62 amu or g_____ How are elements that are gases at room temperature designated in the periodic table? ___their boxes contain a red balloon______________ How many columns of elements does the periodic table contain? ___18___ What is another name for a column of elements? ___group or family_______ 8. What two group numbers can be used to designate elements in the second column of the periodic table? __group 2A or group 2_______

The Periodic Table How many rows of elements does the periodic table contain? ___ What is another name for a row of elements? _____________ Which period contains the least number of elements? ______ What element is found in period 4, group 7B? __________ 13. How are metals designated in this periodic table? __________________________________ 14. How are metalloids designated in this periodic table? _______________________________ 15. How are nonmetals designated in this periodic table? _______________________________ 16. What can be said about the electron configurations of all the elements in a group? _________

The Periodic Table How many rows of elements does the periodic table contain? _7_ What is another name for a row of elements? period Which period contains the least number of elements? Period 1 What element is found in period 4, group 7B? manganese How are metals designated in this periodic table? Boxes are tinted blue How are metalloids designated in this periodic table? Boxes are tinted green How are nonmetals designated in this periodic table? Boxes are tinted yellow 16. What can be said about the electron configurations of all the elements in a group? Their valence electron configurations are identical

The s-, p-, d-, and f-Block Elements

1. What are the four sections, or blocks, of the periodic table? _____________ 2. What does each block represent? _________________________________ 3. What do elements in the s-block have in common? ________________ 4. What is the valence electron configuration of each element in group 1A? ______ 5. What is the valence electron configuration of each element in group 2A? ______ 6. Why does the s-block span two groups of elements? ______________________ 7. Why does the p-block span six groups of elements? _______________________ 8. Why are there no p-block elements in period 1? __________________________

What are the four sections, or blocks, of the periodic table? s-, p-, d- and f-blocks What does each block represent? The energy sublevel being filled by valence electrons What do elements in the s-block have in common? Valence electrons only in the s orbitals 4. What is the valence electron configuration of each element in group 1A? s1 5. What is the valence electron configuration of each element in group 2A? s2 Why does the s-block span two groups of elements? The single s orbital can hold a maximum of two valence electrons Why does the p-block span six groups of elements? The three p orbitals can each hold a maximum of two electrons, resulting in a maximum of six valence electrons, which corresponds to the six columns spanned by the p-block. Why are there no p-block elements in period 1? The p sublevel does not exist for the first principal energy level.

9. What is the ending of the electron configuration of each element in group 4A? _____ 10. What is the electron configuration of neon? __________ 11. In what period does the first d-energy sublevel appear? __________ 12. Why does the d-block span ten groups of elements? _________________________ 13. What is the ending of the electron configuration of each element in group 3B? _____ 14. What is the electron configuration of titanium? _______________ 15. In what period does the first f-energy sublevel appear? ___________ Determine the group, period, and block for the element having the electron configuration [Xe]4f145d106s26p3. a. group_____ b. period ______ c. block _____

9. What is the ending of the electron configuration of each element in group 4A? p2 10. What is the noble gas electron configuration of neon? [He]2s22p6 11. In what period does the first d-energy sublevel appear? Period 4 Why does the d-block span ten groups of elements? The five d orbitals can each hold a maximum of two electrons, resulting in a total of ten possible valence electrons. 13. What is the ending of the electron configuration of each element in group 3B? d1 14. What is the noble gas electron configuration of titanium? [Ar]4s23d2 15. In what period does the first f-energy sublevel appear? Period 6 Determine the group, period, and block for the element having the electron configuration [Xe]4f145d106s26p3. a. group__5A or 15___ b. period __6____ c. block __p___

½ the distance between nuclei in a diatomic molecule
Atomic Radius ½ the distance between nuclei in a diatomic molecule What about non-diatomic atoms? Metals? 58

Atomic Radius Trend down a group: larger
Valence e-’s farther from nucleus Shielding effect (#e-’s between nucleus and valence electrons) decreases pull of nucleus on valence electrons What about non-diatomic atoms? Metals? 59

2. Atomic Radius Trend across a period: smaller
Add e- to valence shell, add p+, stronger pull from nucleus draws e-’s closer. Shielding effect is constant across period Not as noticeable with heavier elements What about non-diatomic atoms? Metals? 60

Atomic Radius

Atomic Radius 1. Which groups and periods of elements are shown in the table of atomic radii? 2. In what unit is atomic radius measured? Express this unit in m. 3. What are the values of the smallest and largest atomic radii shown? What elements have these atomic radii? 4. What happens to atomic radii within a period as the atomic number increases? 5. What accounts for the trend in atomic radii within a period? 6. What happens to atomic radii within a group? 7. What accounts for the trend in atomic radii within a group? 8. a) Divide the atomic radius of Cs by the atomic radius of Li and round to 2 significant figures. Cs:Li b) Divide the atomic radius of Cs by the atomic radius of Rn and round to 2 significant figures. Cs:Rn c) Summarize your findings about a) and b) here:

Atomic Radius Which groups and periods of elements are shown in the table of atomic radii? groups 1A-8A; periods 1-6 2. In what unit is atomic radius measured? pm Express this unit in m m What are the values of the smallest and largest atomic radii shown? What elements have these atomic radii? 31 pm – helium; 265 pm - cesium What happens to atomic radii within a period as the atomic number increases? The atomic radius of the elements within a period generally decreases as the atomic number of the elements increases. What accounts for the trend in atomic radii within a period? With increasing atomic number, the increased positive charge of the nucleus pulls more strongly on the outermost electrons, pulling them closer to the nucleus. The size of the shield stays the same, so becomes less effective. Consequently, the atomic radius decreases. 6. What happens to atomic radii within a group? The atomic radius within a group generally increases as the atomic number of the elements increases. 7. What accounts for the trend in atomic radii within a group? With increasing atomic number, the increased pull by the larger positive charge of the nucleus is offset by the outer electrons’ larger orbitals and by shielding by inner electrons. Consequently, the atomic radius increases. 8. a) Divide the atomic radius of Cs by the atomic radius of Li and round to 2 significant figures. Cs:Li X b) Divide the atomic radius of Cs by the atomic radius of Rn and round to 2 significant figures. Cs:Rn X c) Summarize your findings about a) and b) here:

3. Ionic Radius Cations (+) smaller than original atom
remove e-’s  greater pull from nucleus Anions (-) larger than original atom Increased repulsion swells the shell

Ionic Radius

Ionic Radius 1. In this table of ionic radii, how is the charge of the ions of elements in groups 1A-4A related to the group number? a) Divide the radius of Cs with the radius of its ion: b) Divide the radius of Li with the radius of its ion: c) Divide the radius of Be with the radius of its ion: d) Divide the radius of B with the radius of its ion: e) Summarize your findings about a)-d) here: 3. a) Divide the radius of the F ion with the radius of the neutral F atom: b) Divide the radius of the O ion with the radius of the neutral O atom: c) Divide the radius of the N ion with the radius of the neutral N atom: d) Summarize your findings about a)-c) here: e) Compare and contrast 2 e) and 3 d)

4. Ionization Energy the energy required to remove an electron from an atom in the gas phase (in J or kJ) there is a series of ionization energies for each atom (since > 1 electron can be removed) removing each subsequent electron requires more energy

Diagram from Document Camera

Successive ionization energies (kJ/mol)
Ionization Energy Successive ionization energies (kJ/mol) Element First Second Third Fourth Na 496 4,562 6,912 9,543 Mg 738 1,451 7,733 10,540 Al 578 1,817 2,745 11,577

Successive Ionization Energies

1. What happens to the values of the successive ionization energies of an element? 2. How is a jump in ionization energy related to the valence electrons of the element?

What happens to the values of the successive ionization energies of an element? The values of the successive ionization energies increase. How is a jump in ionization energy related to the valence electrons of the element? The jump occurs after the valence electrons have been removed.

First Ionization Energy

Ionization Energies 1. What is meant by first ionization energy?
Which element has the smallest first ionization energy? The largest? What are their values? What generally happens to the first ionization energy of the elements within a period as the atomic number of the elements increases? What accounts for the general trend in the first ionization energy of the elements within a period? Based on the graph, rank the group 2A elements in periods 2-5 in decreasing order of first ionization energy. What generally happens to the first ionization energy of the elements within a group as the atomic number of the elements increases? What accounts for the general trend in the first ionization energy of the elements within a group?

Ionization Energies 1. What is meant by first ionization energy? First ionization energy is the energy required to remove the first electron from a gaseous atom. Which element has the smallest first ionization energy? The largest? What are their values? rubidium – about 400 kJ/mol; helium – about 2375 kJ/mol What generally happens to the first ionization energy of the elements within a period as the atomic number of the elements increases? The first ionization energy of the elements within a period generally increases as the atomic number of the elements increases. What accounts for the general trend in the first ionization energy of the elements within a period? With increasing atomic number, the increased positive charge of the nucleus produces an increased hold on the valence electrons. Consequently, the first ionization energy increases. Based on the graph, rank the group 2A elements in periods 1-5 in decreasing order of first ionization energy. beryllium, magnesium, calcium, strontium 6. What generally happens to the first ionization energy of the elements within a group as the atomic number of the elements increases? The first ionization energy of the elements within a group generally decreases as the atomic number of the elements increase. 7. What accounts for the general trend in the first ionization energy of the elements within a group?

Summary of Trends in First Ionization Energy
Trend across a period: increases Trend down a group: decreases

5. Electronegativity How much one atom pulls on another atom’s electrons in a bond Only refers to atoms involved in a bond (molecule or compound). Trend across a period: Increases Trend down a group: Decreases

Electronegativity

Electronegativity Increases Decreases 80

6.1 - The Periodic Table: A History

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