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Chapter 4: The Periodic Table

Section 1: How Are Elements Organized?

A. History of the Periodic Table

                             

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So what's the origin of this universally recognized shape or did Mendeleev (father of the periodic table) come down a mountain carrying the famous outline carved in rectangular stone?

The quest for a systematic arrangement of the elements started with the discovery of individual elements. By 1860 about 60 elements were known and a method was needed for organization.  In fact many scientists made significant contributions that eventually enabled Mendeleev to construct his table. The periodic table did not end with Mendeleev but continued to take shape for the next 75 years.

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The development of the periodic table begins with German chemist Johann Dobereiner (1780-1849) who grouped elements based on similarities.  Calcium (atomic weight 40), strontium (atomic weight 88), and barium (atomic weight 137) possess similar chemical prepares. Dobereiner noticed the atomic weight of strontium fell midway between the weights of calcium and barium:

     Ca   Sr   Ba      (40 + 137) ÷ 2 = 88

      40     88     137

Was this merely a coincidence or did some pattern to the arrangement of the elements exist? Dobereiner noticed the same pattern for the alkali metal triad (Li/Na/K) and the halogen triad (Cl/Br/I).

   Li   Na  K         Cl   Br   I

      7     23     39           35    80   127

In 1829 Dobereiner proposed the Law of Triads: Middle element in the triad had atomic weight that was the average of the other two members. Soon other scientists found chemical relationships extended beyond triads. Fluorine was added to Cl/Br/I group; sulfur, oxygen, selenium and tellurium were grouped into a family; nitrogen, phosphorus, arsenic, antimony, and bismuth were classified as another group.

 

First Periodic Table

It was a 19th century geologist who first recognized periodicity in the physical properties of the elements. Alexandre Beguyer de Chancourtois (1820-1886), professor of geology at the School of Mines in Paris, published in 1862 a list of all the known elements. The list was constructed as a helical graph wrapped around a cylinder--elements with similar properties occupied positions on the same vertical line of cylinder (the list also included some ions and compounds).   Using geological terms and published without the diagram, de Chancourtois ideas were completely ignored until the work of Mendeleev.

 

Law of Octaves[pic][pic][pic]

English chemist John Newlands (1837-1898), having arranged the 62 known elements in order of increasing atomic weights, noted that after interval of eight elements similar physical/chemical properties reappeared.  Newlands was the first to formulate the concept of periodicity in the properties of the chemical elements. In 1863 he wrote a paper proposing the Law of Octaves: Elements exhibit similar behavior to the eighth element following it in the table.

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Mendeleev's Periodic Table[pic]

Then in 1869, Russian chemist Dmitri Mendeleev (1834-1907) proposed arranging elements by atomic weights and properties (Lothar Meyer independently reached similar conclusion but published results after Mendeleev).  Mendeleev's periodic table of 1869 contained 17 columns with two partial periods of seven elements each (Li-F & Na-Cl) followed by two nearly complete periods (K-Br & Rb-I).

In 1871 Mendeleev revised the 17-group table with eight columns (the eighth group consisted of transition elements). This table exhibited similarities not only in small units such as the triads, but showed similarities in an entire network of vertical, horizontal, and diagonal relationships. The table contained gaps but Mendeleev predicted the discovery of new elements.  In 1906, Mendeleev came within one vote of receiving the Nobel Prize in chemistry.

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Noble Gases[pic]

Lord Rayleigh (1842-1919) and William Ramsey (1852-1916) greatly enhanced the periodic table by discovering the "inert gases."  In 1895 Rayleigh reported the discovery of a new gaseous element named argon. This element was chemically inert and did not fit any of the known periodic groups. Ramsey followed by discovering the remainder of the inert gases and positioning them in the periodic table. So by 1900, the periodic table was taking shape with elements were arranged by atomic weight.  For example, 16g oxygen reacts with 40g calcium, 88g strontium, or 137g barium. If oxygen used as the reference, then Ca/Sr/Ba assigned atomic weights of 40, 88, and 137 respectively.

Rayleigh (physics) and Ramsey (chemistry) were awarded Nobel prizes in 1904.  The first inert gas compound was made in 1962 (xenon tetrafluoride) and numerous compounds have followed; today the group is more appropriately called the noble gases.

 

Moseley's Periodic Law[pic]

Soon after Rutherford's landmark experiment of discovering the proton in 1911, Henry Moseley (1887-1915) subjected known elements to x-rays. He was able to derive the relationship between x-ray frequency and number of protons. When Moseley arranged the elements according to increasing atomic numbers and not atomic masses, some of the inconsistencies associated with Mendeleev's table were eliminated. The modern periodic table is based on Moseley's Periodic Law (atomic numbers). At age 28, Moseley was killed in action during World War I and as a direct result Britain adopted the policy of exempting scientists from fighting in wars.  Shown below is a periodic table from 1930:

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Modern Periodic Table[pic]

The last major change to the periodic table resulted from Glenn Seaborg's work in the middle of the 20th century. Starting with plutonium in 1940, Seaborg discovered transuranium elements 94 to 102 and reconfigured the periodic table by placing the lanthanide/actinide series at the bottom of the table. In 1951 Seaborg was awarded the Nobel Prize in chemistry and element 106 was later named seaborgium (Sg) in his honor.

Section2 : Reading the Periodic Table

The Periodic table is an arrangement of elements in the order of their increasing atomic numbers to show that elements have related properties.

1. Classification elements into groups with similar properties.

2. To predict the possibilities of new elements based on their properties.

The modern form of the periodic law states that properties of the elements are the periodic function of their atomic numbers and the properties of the elements depend on their electronic configuration.

Main Group Elements ( also known as the Representative Elements

← Vertical columns 1-2, 13-18

← s & p blocks of table

Groups

• The vertical columns are called Groups or Families

• Eighteen columns

• Elements in each group have the same number of electrons in the outer shell and hence the same valency.

• Identical Chemical behavior for each group elements

• Lower down the group, number of energy levels increase by 1; hence atomic radius also increases.

Group 1 or IA: Alkali Metals

Group 2 or IIA: Alkaline Earth Metals

Groups 3-12: Transition Elements (Non-representative elements)

Group 13 or IIIA: Boron Family

Group14 or IVA: Carbon Family

Group15 or VA: Nitrogen Family

Group 16 or VIA: Oxygen Family or Chalcogens

Group 17 or VIIA: Halogens

Group 18 or VIIIA: Noble Gases or Inert Gases

Periods

• The horizontal rows are called Periods or Series

• Seven rows

• No further division

• 6 th period includes Lanthanide series

• 7 th period includes Actinide series

• Each period begins with an alkali metal.

• Number of electrons increase by one, across the period.

• Elements change from metals (Na) through semi-metals (Si-Silicon) to non-metals (Argon-Ar).

• Elements change from reducing agents (Na) to oxidizing agents (Cl).

• Number of shells remain the same; atom size (radii)decreases.

The modern periodic table consists of arrangements of elements in three broad categories.

Metals

Found to the left of the staircase, exception Hydrogen.

All are solids with the exception of mercury

Properties include

Lustrous

Conductors of heat

Conductors of electricity

Ductile

Malleable

Nonmetals

Found to the right of the staircase

They are solids, liquids and mostly gases.

Properties are opposite to that of the metals.

Semi-Metals (Metalloids)

Elements that border the staircase, B, Ge, Si, Sb, As, Te, At

Properties include both properties of metals and nonmetals.

natural elements : occur naturally, elements 1-92

transuranium elements : elements beyond uranium, artificial elements prepared from

other elements during synthesis procedures

Nomenclature system for elements beyond 100

IUPAC (International Union of Pure and Applied Chemistry) devised a naming system for elements beyond 100

____ ____ ____ ium

0 - nil 4 - quad 8 - oct ex) 210

1 - un 5 - pent 9 - enil

2 - bi 6 - hex

3 - tri 7 - sep

Section: Tour of the Periodic Table

1. The _________ have a single electron in the highest occupied energy level.

2. The __________ are in the s- and p-blocks of the periodic table.

3. All the __________have two valence electrons and get to a stable electron configuration by losing two electrons.

4. Unlike the main-group elements, each group of the _________ does not have the identical outer electron configuration.

5. The ____________, the most reactive group of non-metals, achieve stable electron configurations by gaining one electron.

6. The ___________ have a full set of electrons in their outermost energy level.

7. The _____________are very stable and have low reactivity.

8. The ___________ are highly reactive and readily form salts with metals.

9. In general, the __________ are metals that are less reactive than the alkali metals and the alkaline earth metals.

10. The ____________ are metals that lose one electron when they react with water to form alkaline solutions.

11. Most elements are ____________.

12. With its one valence electron, ______________ reacts with many other elements.

13. Groups ________ and _________ compose the main-group elements.

14. The ________________include all members of Groups 1 through 12, as well as some of the elements of Groups ______________ through _______________.

15. Elements in Groups ____________ through ___________, including the two long rows below the table, are called transition elements.

16. In the transition elements, electrons are usually added to the __________ orbital, which is why these elements are also known as the ______________.

17. The __________ include all of the elements in Groups 17 and 18 as well as some members of Groups ___________ through _____________.

18. In the ______________, electrons are being added to the 4f orbitals.

19. In the ______________, electrons are being added to the 5f orbitals.

20. The ________________are unique in that all are unstable and radioactive.

Quiz on Sections 1 and 2

Section 3: Periodic Trends

Trends of the Periodic Table

Note: These are general periodic trends of elements. There are many exceptions to these general rules.

Review

Period - a row of elements on the periodic table. Remember that sentences are written in rows and end with a period.

Group - a column of elements on the periodic table. Remember that group is spelled group and groups go up and down.

Atomic Radius - Atomic radius is simply the radius of the atom, an indication of the atom's volume.

Period - atomic radius decreases as you go from left to right across a period.

Why? Stronger attractive forces in atoms (as you go from left to right) between the opposite charges in the nucleus and electron cloud cause the atom to be 'sucked' together a little tighter.

Group - atomic radius increases as you go down a group.

Why? There is a significant jump in the size of the nucleus (protons + neutrons) each time you move from period to period down a group. Additionally, new energy levels of elections clouds are added to the atom as you move from period to period down a group, making the each atom significantly more massive, both is mass and volume.

Electronegativity - Electronegativity is an atom's 'desire' to grab another atom's electrons.

Period - electronegativity increases as you go from left to right across a period.

Why? Elements on the left of the period table have 1 -2 valence electrons and would rather give those few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's electrons. As a result, they have low electronegativity. Elements on the right side of the period table only need a few electrons to complete the octet, so they have strong desire to grab another atom's electrons.

Group - electronegativity decreases as you go down a group.

Why? Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have so many electrons that loosing or acquiring an electron is not as big a deal. This is due to the shielding affect where electrons in lower energy levels shield the positive charge of the nucleus from outer electrons resulting in those outer electrons not being as tightly bound to the atom.

Ionization Energy - Ionization energy is the amount of energy required to remove the outmost electron. It is closely related to electronegativity.

Period - ionization energy increases as you go from left to right across a period.

Why? Elements on the right of the chart want to take others atom's electron (not given them up) because they are close to achieving the octet. The means it will require more energy to remove the outer most electron. Elements on the left of the chart would prefer to give up their electrons so it is easy to remove them, requiring less energy (low ionization energy).

Group - ionization energy decreases as you go down a group.

Why? The shielding affect makes it easier to remove the outer most electrons from those atoms that have many electrons (those near the bottom of the chart).

Reactivity - Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom's electrons (electronegativity) because it is the transfer/interaction of electrons that is the basis of chemical reactions.

Metals

Period - reactivity decreases as you go from left to right across a period.

Group - reactivity increases as you go down a group

Why? The farther to the left and down the periodic chart you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity.

Non-metals

Period - reactivity increases as you go from the left to the right across a period.

Group - reactivity decreases as you go down the group.

Why? The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electron.

Ionic Radius vs. Atomic Radius

Metals - the atomic radius of a metal is generally larger than the ionic radius of the same element.

Why? Generally, metals loose electrons to achieve the octet. This creates a larger positive charge in the nucleus than the negative charge in the electron cloud, causing the electron cloud to be drawn a little closer to the nucleus as an ion.

Non-metals - the atomic radius of a non-metal is generally smaller than the ionic radius of the same element.

Why? Generally, non-metals loose electrons to achieve the octet. This creates a larger negative charge in the electron cloud than positive charge in the nucleus, causing the electron cloud to 'puff out' a little bit as an ion.

Melting Point

Metals - the melting point for metals generally decreases as you go down a group.

Non-metals - the melting point for non-metals generally increases as you go down a group.

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Summary of Periodic Table Trends

Moving Left --> Right

• Atomic Radius Decreases

• Ionization Energy Increases

• Electronegativity Increases

Moving Top --> Bottom

• Atomic Radius Increases

• Ionization Energy Decreases

• Electronegativity Decreases

Chapter Test

Experiment: Periodic Trends

Objective: To discover the periodic trends of certain physical properties of elements related to their position on the Periodic Table of Elements.

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Background: The Periodic Table is arranged according to the Periodic Law. The Periodic Law states that when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern. Students can discover these patterns by examining the changes in properties of elements on the Periodic Table. The properties that will be examined in this lesson are: atomic radius, first ionization energy, and electronegativity.

Atomic radius => distance from the center of an atom's nucleus to its outer most electron

First ionization energy => the amount of energy needed to remove one (the outermost) electron from an atom

Electronegativity => measure of an atoms attraction for electrons in a chemical bond

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Lesson: In this exercise you will look at a few physical properties of elements and how those properties are related to their position on the Periodic Table. Analyze the data found on the Periodic Table sites to answer the questions listed below.

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Activities:[pic]

1. Explore some on-line periodic tables such as Web Elements, Chemicool, Chem4Kids, Chemical Elements, Periodic Table from Los Alamos, Dynamic Periodic Table

2. Using the Chemicool or Dynamic Periodic Table

a. Find the atomic radius for each of the first five elements in Period 2 (Li, Be, B, C & N) (click on the element symbol).

b. Find the atomic radius for each of the first five elements in Period 3 (Na, Mg, Al, Si & P) (click on the element symbol).

c. Find the atomic radius for these elements in Period 4 (K, Ca, Mn, Fe, Zn, Ga, Br & Kr) (click on the element symbol).

d. What appears to be the trend in atomic radius as you move from left to right in a row?

e. What appears to be the trend in atomic radius as you move down a column?

f. Is the pattern of atomic radius absolute or general (always true or generally true)?

3. Using the Chemicool or Dynamic Periodic Table

a. Find the 1st ionization energy for each of the first five elements in Period 2 (Li, Be, B, C & N) (click on the element symbol).

b. Find the 1st ionization energy for each of the first five elements in Period 3 (Na, Mg, Al, Si & P) (click on the element symbol).

c. Find the 1st ionization energy for these elements in Period 4 (K, Ca, Mn, Fe, Zn, Ga, Br & Kr) (click on the element symbol).

d. What appears to be the trend in 1st ionization energy as you move from left to right in a row?

e. What appears to be the trend in 1st ionization energy as you move down a column?

f. Is the pattern of atomic radius absolute or general (always true or generally true)?

4. Using the Chemicool or Dynamic Periodic Table

a. Find the electronegativity for each of the first five elements in Period 2 (Li, Be, B, C & N) (click on the element symbol).

b. Find the electronegativity for each of the first five elements in Period 3 (Na, Mg, Al, Si & P) (click on the element symbol).

c. Find the electronegativity for these elements in Period 4 (K, Ca, Mn, Fe, Zn, Ga, Br & Kr) (click on the element symbol).

d. What appears to be the trend in electronegativity as you move from left to right in a row?

e. What appears to be the trend in electronegativity as you move down a column?

f. Is the pattern of atomic radius absolute or general (always true or generally true)?

5. Consider all three of the properties that you have examined.

a. State the general trend for each property if you move from left to right on the Periodic Table. Now, state the general trend from top to bottom.

b. How do these properties show periodicity (periodic trends)?

6. Using the link examine and view the line graph of atomic radius.

a. What do the different colors show?

b. Can you see a pattern in the second period that is repeated in the third period?

c. How does this graph agree with your observations of atomic radius made earlier?

d. Why do the fourth and fifth periods have more dots and different patterns?

7. Using the link examine and view the line graph of 1st ionization energy (enthalpy).

a. What do the different colors show?

b. Can you see a pattern in the second period that is repeated in the third period?

c. How does this graph agree with your observations of 1st ionization energy made earlier?

d. Why do the fourth and fifth periods have more dots and different patterns?

8. Using the link examine and view the line graph of electronegativity.

a. What do the different colors show?

b. Can you see a pattern in the second period that is repeated in the third period?

c. How does this graph agree with your observations of electronegativity made earlier?

d. Why do the fourth and fifth periods have more dots and different patterns?

9. Using the color-coded tables, atomic radius, 1st ionization energy and electronegativity, answer the questions below.

a. How does this show periodic trends of the selected property?

b. Which method did you find most informative?

c. Which method was easiest to see the general pattern and not get confused by exceptions in that pattern?

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