PERIODIC TABLE - Loudoun County Public Schools



PERIODIC TABLE

Development of the Periodic Table (pp. 348 - 352)

1. In the late 1790s Lavoisier compiled a list of 23 known elements.

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2. During the 1800s many more were discovered as scientists used electricity to separate compounds and newly developed instruments, such as the spectrometer, to identify the elements. By 1870, there were approximately 70 known elements.

3. Chemists were overwhelmed by the volume of new information about the elements. A tool for organizing all this information was needed.

4. A big step came in 1860 when chemists agreed upon a method for accurately determining atomic masses of the elements.

5. The first attempt at organizing the elements was made by John Newlands. He noticed that properties of the elements were repeated every eighth element. He called this the _________________________. This law was not widely accepted because it did not work for all the known elements.

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6. Even so, Newlands was correct that the properties of the elements do repeat in a periodic way.

7. Dimitri Mendeleev also arranged the elements by increasing atomic mass. He did not limit the length of his rows. He noticed that the elements fell into columns of elements with similar properties.

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8. Mendeleev also left empty spaces to account for elements that had not yet been discovered. He was able to correctly predict the undiscovered element’s properties.

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9. There were problems with Mendeleev’s table because as new elements were discovered and atomic masses better determined, some of the elements were placed out of order.

10. Henry Moseley, who had discovered that the atoms of each element contain a unique number of protons in their nuclei, arranged the elements by their atomic number. This adjustment fixed the periodic table.

11. The modern periodic table is based on the _________________________, which states that the physical and chemical properties of the elements tend to change with increasing atomic number in a periodic way.

Modern Periodic Table

1. Each element has its own box containing information about the element.

2. The boxes are arranged in order of increasing atomic number into a series of columns called ____________________, or families, and rows called ____________________.

Regions of the Periodic Table

1. There are five major regions of the periodic table:

2. The metals are the largest region of the periodic table. Elements within this region are solids, generally lustrous, ductile and malleable. They are also excellent conductors of heat and electricity.

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3. Nonmetals make up the second largest region of the table. These elements have a variety of physical states and properties. In general they are poor conductors of heat and electricity, brittle and not lustrous.

4. Metalloids are elements that have properties of both metals and nonmetals. These elements are sandwiched between the metals and nonmetals along the “stair step”. There are seven of them:

5. The noble gases are the gases located in Group 18. These gases do not have a tendency to react with anything.

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6. Hydrogen, because of its unique properties, is also its own region.

FAMILY CHARACTERISTICS

Group 18: Noble Gases

1. The noble gases were once called inert gases because they were thought to be unreactive.

2. No stable compounds of helium, neon and argon have ever been formed.

3. The other noble gases – xenon, krypton, and radon – have very low reactivity. They have been forced to form compounds.

4. Noble gases have full orbitals in the highest energy level, called an ____________________.

5. From this low reactivity we can infer that the noble gas electron configuration is very ____________________.

6. Atoms from the other 17 groups either gain or lose electrons to form compounds. In doing so, the atoms achieve an electron configuration like that of the noble gases.

Group 1: Alkali Metals

1. These elements have metallic properties: soft, shiny, highly reactive, can be cut with a knife, react with atmospheric oxygen, good conductors of heat and electricity.

2. Reactivity comes from their characteristic electron configuration – a single electron in the highest energy level.

3. By losing an electron, an alkali metal atom achieves the stable, nonreactive electron configuration of the noble gas.

Group 2: Alkaline Earth Metals

1. Harder, denser, stronger and have higher melting points than Group 1 elements.

2. Group 2 elements are less reactive than Group 1 elements.

3. This is because they must lose 2 electrons to achieve a noble gas electron configuration.

Groups 3 - 12: Transition Elements

1. These are all metals, but not as reactive as the elements in Groups 1 and 2.

2. They are harder, denser and have higher melting points.

3. There are many irregularities in the electron configurations of the transition elements. These variations occur in the way some of the s-orbitals fill. The transition elements are where the d-orbitals begin to fill.

4. The bottom two rows are known as the _________________________. Each also has its own name. The _________________________, elements 58 – 71, and the _________________________,

elements 90 – 103.

5. The lanthanides are shiny, reactive metals.

6. The actinides have an unstable arrangement of protons and neutrons in the nucleus. They are usually radioactive.

Groups 13 – 18: Main Block Elements

1. These are also known as the _________________________ because they represent a wide range of chemical and physical properties.

2. Within each group the properties may vary systematically.

Group 17: Halogens

1. The halogens combine easily with metals, especially the alkali metals, to form compounds known as ____________________.

2. The halogens are the most reactive nonmetal elements. Their electron configuration is one electron short of the noble gas configuration.

Hydrogen

1. Most common element in the universe.

2. It has only one electron and reacts very rapidly with most other elements.

***** Identify the region from which each of the following elements comes:

Ba I

Sb Rn

***** For each of the given elements, list two other elements with similar chemical properties:

Fe Br

Se Rb

CLASSIFICATION OF THE ELEMENTS

Organizing by Electron Configuration

1. Elements within the same group have the same electron configuration for their outermost energy level.

2. These electrons are the valence electrons.

3. Atoms in the same group have similar chemical properties because they have the same number of valence electrons.

4. The energy level of an element’s valence electrons indicates the period in the periodic table in which it is found.

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

1. The s-block consists of groups 1 and 2 and hydrogen and helium. In this block, the valence electrons occupy only s-orbitals.

2. The p-block elements consist of groups 13 – 18. As one progresses from left to right, one more electron is added to the p-orbital until it is filled with six electrons. The noble gases exhibit great stability because both their s- and p-orbitals are filled.

3. The d-block contains the transition elements and it is the largest of the blocks. These elements are characterized by a filled outermost s-orbital and a filled, or partially filled, d-orbital.

4. The f-block elements contain the inner transition elements from the lanthanide and actinide series.

***** Strontium has an abbreviated electron configuration of [Kr]5s2. Without using the periodic table, determine the group, period and block in which strontium is located.

***** Write the abbreviated electron configuration of the element in Group 12, period 4.

PERIODIC TRENDS

1. The _________________________ states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

2. Because the elements are arranged side by side in order of increasing atomic number, one can view certain important vertical and horizontal trends.

Atomic Radius

1. Recall from Rutherford’s experiments that the nucleus was found to occupy a small fraction of the atom’s entire volume.

2. It is the electron cloud surrounding the nucleus that determines the boundaries of the atom.

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3. Because the electrons travel in this “cloud” region, it is difficult to measure the size of an atom.

4. The atomic radius is determined in two primary ways. One is to measure the distance between centers of like atoms joined together in a diatomic molecule. The other is to measure the bond lengths of atoms in compounds.

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5. Looking at the periodic table one can notice things that affect the size of an atom.

6. As one moves down through a group, a new principal energy level is added. Each new level is farther from the nucleus.

7. As each energy level is added, the energy levels and electrons closer to the nucleus ____________________ it from the effects of the positive charge. This _________________________ allows the outer electrons to be farther away.

8. When crossing a period from left to right each atom gains one more proton and one more electron.

9. No principal energy level is added so the electrons enter the same energy level. The additional protons in the nucleus provide more pull on the electrons bringing them closer to the nucleus.

10. Trend:

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Ionic Radius

1. Atoms can gain or lose electrons to form charged particles.

2. These particles are called _______________.

3. When atoms lose an electron they form a positively charged ion and become smaller in size. Positive ions are smaller than their parent atom.

4. When atoms gain an electron they form a negatively charged ion and become larger in size. Negative ions are larger than their parent atom.

Ionization Energy

1. To form a positive ion, an electron must be removed from a neutral atom.

2. The electron is removed from the outermost energy level, known as the _________________________.

3. The electrons in that “shell” are known as _________________________________________________.

4. This requires energy to overcome the attraction between the positive charge in the nucleus and the negative charge of the electron.

6. This energy is known as the _________________________.

7. Ionization energy is defined as the energy required to remove an electron from a gaseous atom.

8. The energy required to remove the first electron from an atom is called the first ionization energy.

9. The process is:

10. As one moves down a group (column) it becomes easier to remove electrons because they are farther away from the nucleus and the shielding effect increases.

11. As one moves left to right across a period it becomes harder to remove electrons because there is no increase in the shielding effect, the positive charge in the nucleus increases, and the electrons are closer to the nucleus.

12. Trend:

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Electronegativity

1. The ____________________ of an element indicates the relative ability of an atom to attract electrons to itself when the atom is involved in a chemical bond.

2. These values cannot be measured. In fact, they were developed by a chemist, Linus Pauling, as a way to explain chemical bonding in molecules. They are based on a fluorine being the element with the strongest attraction for electrons, so it is a relative scale.

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3. The noble gases are ignored since they are basically inert.

4. Fluorine is the most electronegative element with a value of 3.98 while cesium and francium are the least electronegative.

5. In a chemical bond, the atom with the greater electronegativity more strongly attracts the electrons in that bond.

6. Trend:

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***** In each of the following sets of elements, which element would be expected to have the highest ionization energy?

a. Cs, K, Li b. Ba, Sr, Ca c. I, Br, Cl d. Mg, Si, S

***** Arrange the following sets of elements in order of increasing atomic size.

a. Sn, Xe, Rb, Sr

b. Rn, He, Xe, Kr

c. Pb, Ba, Cs, At

***** In each of the following sets of elements, indicate which element has the smallest electronegativity value.

a. Na, K, Rb b. S, Na, Si c. P, N, As d. O, N, F

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