6.1 - The Periodic Table: A History

http://www.woodrow.org/teachers/ci/1992/MENDELEEV.GIF

6.1 - The Periodic Table: A History

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

http://www.glogster.com/media/1/6/49/85/6498532.jpg

History of the Periodic Table

I. 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



1848 57 elements

1860

Karlsruhe Congress

(big Chemistry Conference)

Germany

John Newlands 1865• English Chemist

• Arranged elements by atomic mass

• Described the “Rule of octaves”

• 62 elements

http://www.rsc.org/education/teachers/learnnet/periodictable/scientists/newlands.jpg

Lothar Meyer 1870

• German Chemist

• Arranged elements based on atomic mass

• Discovered periodic properties related to atomic volume

• Established concept of valency

http://www.chemistrydaily.com/chemistry/Lothar_Meyer

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

http://anhso.net/data/69/X_kun/571478/mendeleev18371.jpg

• 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

Dmitri Mendeleev 1869

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



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

dissolves in acid


Predicted Element



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

– Arranged by atomic mass

– Triads: [Cl Br I], [Ca Ba Sr], [Li Na K]

• Newlands 1865 – Arranged by atomic mass

– Rule of Octaves

• Meyer 1870– Arranged by atomic mass, periodic trend with atomic volume

– Established concept of valency

• Mendeleev– Arranged by atomic mass

– Repetition every 8 or 18 elements

– Predicted 3 elements not yet discovered: eka-aluminium - gallium, eka-silicon - germanium and eka-boron - scandium

Discovery of the Noble Gases1890s

• Lord Rayleigh (physicist) and Sir William Ramsay (chemist)

• 1894 - Argon “the lazy one”, discovered when Ramsay was trying to isolate nitrogen

• 1895 - Helium – found on earth in uranium minerals (found in the sun in 1868)

• 1898 - Neon “the new one”Krypton “the hidden one”Xenon “the alien one”

• 1910 – Radon

Properties:Largely unreactive8 electrons in valence shellLow boiling and melting points

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

http://www.explicatorium.com/images/Personalidades/Henry_Moseley.jpg


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

Letter to Seaborg

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 of the Periodic Table “periodic” = repeating pattern

Electron configuration

( reactivity and bonding)

1. Atomic radius

2. Ionization energy

3. Electronegativity

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

1. 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

Atomic Radius

Atomic Radius

Atomic Radius1. 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 Radius1. 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. 10-12 m 3. What are the values of the smallest and largest atomic radii shown?

What elements have these atomic radii? 31 pm – helium; 265 pm - cesium

4. 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.

5. 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 1.7 Xb) Divide the atomic radius of Cs by the atomic radius of Rn and round to 2 significant figures. Cs:Rn 1.9 Xc) Summarize your findings about a) and b) here:

2. 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

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?


1. What happens to the values of the successive ionization energies of an element?

The values of the successive ionization energies increase.

2. 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 Energies1. What is meant by first ionization energy? 2. Which element has the smallest first ionization energy?

The largest? What are their values? 3. What generally happens to the first ionization energy of

the elements within a period as the atomic number of the elements increases?

4. What accounts for the general trend in the first ionization energy of the elements within a period?

5. Based on the graph, rank the group 2A elements in periods 2-5 in decreasing order of first ionization energy.

8. What generally happens to the first ionization energy of the elements within a group as the atomic number of the elements increases?

9. What accounts for the general trend in the first ionization energy of the elements within a group?

Ionization Energies1. What is meant by first ionization energy? First ionization energy is the energy

required to remove the first electron from a gaseous atom.2. Which element has the smallest first ionization energy? The largest? What are their

values? rubidium – about 400 kJ/mol; helium – about 2375 kJ/mol3. 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.

4. 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.

5. 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

3. 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

ElectronegativityIncreases

Decreases


Twelve elements have been known since ancient times.

What do you think they are?

(Name them, use your periodic table to help you.)


Why do you think these particular elements have been known for so long, while most elements were not discovered until the 1800s and 1900s?

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

1. poor heat conductors

2. poor electrical conductors

3. not lustrous

4. brittle

5. variety of colors

6. gases or brittle solids at room T

1. extremely unreactive – “inert”

2. rarely form compounds with other elements

3. colorless, odorless gases at room T

The Periodic Table

The Periodic Table1. How many elements are listed in the periodic table? (the one Dr. Hart gave you…)

__________

2. What is the atomic number of selenium? _________

3. What is the symbol for palladium? _________

4. What is the atomic mass of strontium? ________

5. How are elements that are gases at room temperature designated in the periodic table? _________________

6. How many columns of elements does the periodic table contain? ______

7. 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 Table1. How many elements are listed in the periodic table? (the one Dr. Hart gave you…)

___118_______

2. What is the atomic number of selenium? __34____

3. What is the symbol for palladium? ___Pd______

4. What is the atomic mass of strontium? ___87.62 amu or g_____

5. How are elements that are gases at room temperature designated in the periodic table? ___their boxes contain a red balloon______________

6. How many columns of elements does the periodic table contain? ___18___

7. 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 Table9. How many rows of elements does the periodic table contain? ___

10. What is another name for a row of elements? _____________

11. Which period contains the least number of elements? ______

12. 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 Table9. How many rows of elements does the periodic table contain? _7_

10. What is another name for a row of elements? period

11. Which period contains the least number of elements? Period 1

12. What element is found in period 4, group 7B? manganese

13. How are metals designated in this periodic table? Boxes are tinted blue

14. How are metalloids designated in this periodic table? Boxes are tinted green

15. 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? __________________________


1. What are the four sections, or blocks, of the periodic table? s-, p-, d- and f-blocks

2. What does each block represent? The energy sublevel being filled by valence electrons

3. 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

6. Why does the s-block span two groups of elements? The single s orbital can hold a maximum of two valence electrons

7. 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.

8. 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?

___________16. 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 412. 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 616. 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___

Warmup

• Name the four scientists and the scientific meeting we talked about Wednesday

• 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

Electron Configuration

Compare the charges on the ion list with the positionof 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

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.


Element Electron config

Ion Ion

# e-’s

Ion

# p+

O

F

Ne

Na

Mg

Prediction: smallest to largest:


Element Electron config

Ion Ion

# e-’s

Ion

# p+

O1s22s22p4

1s22s22p6 O2- 10 e- 8 p+

F1s22s22p5

1s22s22p6 F- 10 e- 9 p+

Ne1s22s22p6

1s22s22p6 Ne 10 e- 10 p+

Na1s22s22p63s1

1s22s22p6 Na+ 10 e- 11 p+

Mg1s22s22p63s2

1s22s22p6 Mg2+ 10 e- 12 p+

Prediction: smallest to largest: Mg2+ < Na+ < Ne < F-< O2-

Atomic Radius½ the distance between nuclei in a diatomic molecule

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

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 Radius1. In this table of ionic radii, how is the charge of the ions of elements in groups

1A-4A related to the group number?

2. 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)

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6.1 - The Periodic Table: A History