Periodic Table Study Guide

1.5 Periodic Table (Chem.1.5)

Explore this Phenomenon

This is a picture of about 3 pounds of sodium reacting with water. Sodium is an alkali metal.

1. Would you expect the other alkali metals to react the same way?
2. Would you expect the elements in the same row to react the same way?

 

Standard Chem.1.5

Use the periodic table as a mode l to predict the relative properties of elements based on the patterns of electrons in the outermost energy level of atoms. Emphasize conceptual understanding of trends and patterns. Examples could include trends in ionization energy, atomic radius, or electronegativity. Examples of properties for main group elements could include general reactivity, bonding type, or ion formation. (PS1.A)

As you read the following section, identify patterns of the periodic table such as atomic radius, amount of protons, neutrons, and electrons, and/or physical properties.

 

The Periodic Law

The first periodic table was developed by Dmitri Mendeleev, organized by atomic weight. When Mendeleev put his periodic table together, nobody knew about the existence of the nucleus . It was not until 1911 that Rutherford conducted his gold foil experiment that demonstrated the presence of the nucleus in the atom . Just two years later, in 1913, English physicist Henry Moseley (1887-1915) examined x-ray spectra of a number of chemical elements. He would shoot X-rays through crystals of the element and study the wavelengths of the radiation he detected. Moseley found that there was a relationship between wavelength and atomic number . His results led to the definition of atomic number as the number of protons contained in the nucleus of each atom. He then realized that the elements of the periodic table should be arranged in order of increasing atomic number rather than increasing atomic mass.

When ordered by atomic number , the discrepancies within Mendeleev’s table disappeared. Tellurium has an atomic number of 52, while iodine has an atomic number of 53. So even though tellurium does indeed have a greater atomic mass than iodine, it is properly placed before iodine in the periodic table. Mendeleev and Moseley are credited with being most responsible for the modern periodic law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their chemical and physical properties . The result is the periodic table as we know it today. Each new horizontal row of the periodic table corresponds to the beginning of a new period because a new principal energy level is being filled with electrons. Elements with similar chemical properties appear at regular intervals, within the vertical columns called groups.

 

The Modern Periodic Table

The periodic table has undergone extensive changes in the time since it was originally developed by Mendeleev and Moseley. Many new elements have been discovered, while others have been artificially synthesized. Each fits properly into a group of elements with similar properties. The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties appear in the same vertical column or group.

The figure below shows the most commonly used form of the periodic table. Each square shows the chemical symbol of the element along with its name. Notice that several of the symbols seem to be unrelated to the name of the element: Fe for iron, Pb for lead, etc. Most of these are the elements that have been known since ancient times and have symbols based on their Latin names. The atomic number of each element is written above the symbol.

A period is a horizontal row of the periodic table. There are seven periods in the periodic table, with each one beginning at the far left. A new period begins when a new principal energy level begins filling with electrons. Period 1 has only two elements (hydrogen and helium), while periods 2 and 3 have 8 elements. Periods 4 and 5 have 18 elements. Periods 6 and 7 have 32 elements because the two bottom rows that are separated from the rest of the table belong to those periods. They are pulled out in order to make the table itself fit more easily onto a single page.

A group is a vertical column of the periodic table, based on the organization of the outer shell electrons. There are a total of 18 groups. There are two different numbering systems that are commonly used to designate groups and you should be familiar with both. The traditional system used in the United States involves the use of the letters A and B. The first two groups are 1A and 2A, while the last six groups are 3A through 8A. The middle groups use B in their titles. Unfortunately, there was a slightly different system in place in Europe. To eliminate confusion the International Union of Pure and Applied Chemistry (IUPAC) decided that the official system for numbering groups would be a simple 1 through 18 from left to right. Many periodic tables show both systems simultaneously.

 

Can you guess what types of metal screws are made of?

Screws come in all sizes and shapes. They are all (well, almost all) made of some kind of metal. But they have differences in size, shape, and type of metal. Physical characteristics also differ. Some screws are long, and others are short. One screw may have a flat-head slot while another screw may have a Phillips-head. Some of the screws in the picture below are used to fasten things together, and others are used to hang heavy objects on a wall.

Chemists classify materials in many ways. We can sort elements on the basis of their electron arrangements. The way the electrons are distributed determines the chemical properties of the element . Another way is to classify elements based on physical properties . Some common physical properties are color , volume, and density. Other properties that allow us to sort on the basis of behavior are conduction of heat and electricity, malleability (the ability to be hammered into very thin sheets), ductility (the ability to be pulled into thin wires), melting point, and boiling point. Three broad classes of elements based on physical properties are metals, nonmetals, and metalloids.

 

Metals

A metal is an element that is a good conductor of heat and electricity. Metals are also malleable, which means that they can be hammered into very thin sheets without breaking. They are ductile, which means that they can be drawn into wires. When a fresh surface of any metal is exposed, it will be very shiny because it reflects light well. This is called luster. All metals are solid at room temperature with the exception of mercury (Hg), which is a liquid . Melting points of metals display a very wide variance. The melting point of mercury is -39°C, while the highest melting metal is tungsten (W), with a melting point of 3422°C. The elements in blue in the periodic table below are metals. About 80 percent of the elements are metals.

Gold has been used by many civilizations for making jewelry (see Figure below ). This metal is soft and easily shaped into a variety of items. Since gold is very valuable and often used as currency, gold jewelry has also often represented wealth.

 

 

Gold jewelry

Copper is a good conductor of electricity and is very flexible and ductile. This metal is widely used to conduct electric current in a variety of appliances, from lamps to stereo systems to complex electronic devices (see Figure below ).

 

 

Copper wire exposed.

Mercury is the only metal to exist as a liquid at room temperature (see Figure below ). This metal was extensively used in thermometers for decades until information about its toxicity became known. Mercury switches were once common, but are no longer used. However, new federally-mandated energy-efficient light bulbs that are now used contain trace amounts of mercury and represent hazardous waste.

 

 

Pouring Mercury.

When we sort parts in our shop or garage, we often classify them in terms of common properties. One container might hold all the screws (possibly sub-divided by size and type). Another container would be for nails. Maybe there is a set of drawers for plumbing parts.

When you get finished, you could also have a collection of things that don’t nicely fit a category. You define them in terms of what they are not. They are not electrical components, or sprinkler heads for the yard, or parts for the car. These parts may have some common properties, but are a variety of items.

 

Nonmetals

In the chemical world, these “spare parts” would be considered nonmetals, loosely defined as not having the properties of metals . A nonmetal is an element that is generally a poor conductor of heat and electricity. Most properties of nonmetals are the opposite of metals . There is a wider variation in properties among the nonmetals than among the metals. Nonmetals exist in all three states of matter . The majority are gases, such as nitrogen and oxygen. Bromine is a liquid . A few are solids, such as carbon and sulfur. In the solid state, nonmetals are brittle, meaning that they will shatter if struck with a hammer. The solids are not lustrous. Melting points are generally much lower than those of metals. The green elements in the table below are nonmetals.

Nonmetals have a wide variety of uses. Sulfur can be employed in gunpowder, fireworks, and matches to facilitate ignition (see Figure below ). This element is also widely used as an insecticide, a fumigant, or a means of eliminating certain types of fungus. An important role for sulfur is the manufacture of rubber for tires and other materials. First discovered in 1839 by Charles Goodyear, the process of vulcanization makes rubber more flexible and elastic as well as being more resistant to changes in temperature . A major use of sulfur is for the preparation of sulfur-containing compounds such as sulfuric acid.

 

Bromine is a versatile element, used mainly in the manufacture of flame-retardant materials, especially important for children’s clothing (see Figure below ). For treatment of water in swimming pools and hot tubs, bromine is beginning to replace chlorine as a disinfectant because of its higher effectiveness. When incorporated into compounds, bromine atoms play important roles in pharmaceuticals for treatment of pain, cancer, and Alzheimer’s disease.

 

Helium is one of the many nonmetals that is a gas . Other nonmetal gases include hydrogen, fluorine, chlorine, and all the group eighteen noble (or inert) gases. Helium is chemically non-reactive, so it is useful for applications such as balloons (see Figure below) and lasers , where non-flammability is extremely important. Liquid helium exists at an extremely low temperature and can be used to cool superconducting magnets for imaging studies (MRI, magnetic resonance imaging). Leaks in vessels and many types of high-vacuum apparatus can be detected using helium. Inhaling helium changes the speed of sound , producing a higher pitch in your voice. This is definitely an unsafe practice and can lead to physical harm and death.

 

Metalloids

Some elements are “none of the above.” They don’t fit neatly into the categories of metal or non-metal because of their characteristics. A metalloid is an element that has properties that are intermediate between those of metals and nonmetals . Metalloids can also be called semimetals. On the periodic table, the elements colored yellow, which generally border the stair-step line, are considered to be metalloids. Notice that aluminum borders the line, but it is considered to be a metal since all of its properties are like those of metals.

 

Examples of Metalloids

Silicon is a typical metalloid. It has a luster like metal, but is brittle like a nonmetal. Silicon is used extensively in computer chips and other electronics because its electrical conductivity is in between that of a metal and a nonmetal.

 Boron is a versatile element that can be incorporated into a number of compounds. Borosilicate glass is extremely resistant to thermal shock. Extreme changes in the temperature of objects containing borosilicates will not create any damage to the material, unlike other glass compositions, which would crack or shatter. Because of their strength, boron filaments are used as light, high-strength materials for airplanes, golf clubs, and fishing rods. Sodium tetraborate is widely used in fiberglass as insulation and also is employed in many detergents and cleaners.

 

Arsenic has long played a role in murder mysteries, being used to commit the foul deed (see Figure below). This use of the material is not very smart since arsenic can be easily detected on autopsy. We find arsenic in pesticides, herbicides, and insecticides, but the use of arsenic for these applications is decreasing due to the toxicity of the metal. Its effectiveness as an insecticide has led arsenic to be used as a wood preservation.

 

Antimony is a brittle, bluish-white metallic material that is a poor conductor of electricity (see Figure below ). Used with lead, antimony increases the hardness and strength of the mixture . This material plays an important role in the fabrication of electronic and semiconductor devices. About half of the antimony used industrially is employed in the production of batteries , bullets, and alloys.

 

 

Atomic Radius

The size of atoms is important when trying to explain the behavior of atoms or compounds. One of the ways we can express the size of atoms is with the atomic radius. This data helps us understand why some molecules fit together and why other molecules have parts that get too crowded under certain conditions.

The size of an atom is defined by the edge of its orbital . However, orbital boundaries are fuzzy and in fact are variable under different conditions. In order to standardize the measurement of atomic radii, the distance between the nuclei of two identical atoms bonded together is measured. The atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together.

The atomic radius (r) of an atom can be defined as one half the distance (d) between two nuclei in a diatomic molecule.

Atomic radii have been measured for elements. The units for atomic radii are picometers, equal to 10 -12 meters. As an example, the internuclear distance between the two hydrogen atoms in an H 2 molecule is measured to be 74 pm. Therefore, the atomic radius of a hydrogen atom is 742=37 pm.

Atomic radii of the representative elements measured in picometers.

 

Periodic Trend

The atomic radius of atoms generally decreases from left to right across a period. There are some small exceptions, such as the oxygen radius being slightly greater than the nitrogen radius. Within a period, protons are added to the nucleus as electrons are being added to the same principal energy level . These electrons are gradually pulled closer to the nucleus because of its increased positive charge. Since the force of attraction between nuclei and electrons increases, the size of the atoms decreases. The effect lessens as one moves further to the right in a period because of electron- electron repulsions that would otherwise cause the atom’s size to increase.

 

Group Trend

The atomic radius of atoms generally increases from top to bottom within a group. As the atomic number increases down a group, there is again an increase in the positive nuclear charge. However, there is also an increase in the number of occupied principle energy levels. Higher principal energy levels consist of orbitals which are larger in size than the orbitals from lower energy levels. The effect of the greater number of principal energy levels outweighs the increase in nuclear charge and so atomic radius increases down a group.

A graph of atomic radius plotted versus atomic number. Each successive period is shown in a different color. As the atomic number increases within a period, the atomic radius decreases.

 

Atoms to Ions

Atoms cannot only gain extra electrons. They can also lose electrons. In either case, they become ions. Ions are atoms that have a positive or negative charge because they have unequal numbers of protons and electrons. If atoms lose electrons, they become positive ions, or cations. If atoms gain electrons, they become negative ions, or anions. Consider the example of fluorine (see Figure below ). A fluorine atom has nine protons and nine electrons, so it is electrically neutral. If a fluorine atom gains an electron , it becomes a fluoride ion with an electric charge of -1.

 

How Ions Form

The process in which an atom becomes an ion is called ionization. It may occur when atoms are exposed to high levels of radiation. The radiation may give their outer electrons enough energy to escape from the attraction of the positive nucleus . However, most ions form when atoms transfer electrons to or from other atoms or molecules. For example, sodium atoms may transfer electrons to chlorine atoms. This forms positive sodium ions (Na + ) and negative chloride ions (Cl - ).

Atoms form ions by losing or gaining electrons because it makes the atom more stable and this state takes less energy to maintain. The most stable state for an atom is to have its outermost energy level filled with the maximum possible number of electrons. In the case of metals such as lithium, with just one electron in the outermost energy level, a more stable state can be achieved by losing that one outer electron. In the case of nonmetals such as fluorine, which has seven electrons in the outermost energy level, a more stable state can be achieved by gaining one electron and filling up the outer energy level.

 

Properties of Ions

Ions are highly reactive, especially as gases. They usually react with ions of opposite charge to form neutral compounds. For example, positive sodium ions and negative chloride ions react to form the neutral compound sodium chloride, commonly known as table salt. This occurs because oppositely charged ions attract each other. Ions with the same charge, on the other hand, repel each other. Ions are also deflected by a magnetic field , as you saw in the opening image of the northern lights.

 

Ionization Energy

Ionization energy is the energy required to remove an electron from a specific atom . It is measured in kJ/mol, which is an energy unit, much like calories. The ionization energies associated with some elements are described in Table below . For any given atom , the outermost valence electrons will have lower ionization energies than the inner-shell kernel electrons. As more electrons are added to a nucleus , the outer electrons become shielded from the nucleus by the inner shell electrons. This is called electron shielding.

 

Element	IE1	IE2	IE3	IE4	IE5	IE6
H	1312
He	2373	5251
Li	620	7300	11,815
Be	899	1757	14,850	21,005
B	801	2430	3660	25,000	32,820
C	1086	2350	4620	6220	38,000	47,261
N	1400	2860	4580	7500	9400	53,000
O	1314	3390	5300	7470	11,000	13,000

 

If we plot the first ionization energies vs. atomic number for the main group elements, we would have the following trend:

Ionization energy and atomic number.

Moving from left to right across the periodic table, the ionization energy for an atom increases. We can explain this by considering the nuclear charge of the atom. The more protons in the nucleus , the stronger the attraction of the nucleus to electrons. This stronger attraction makes it more difficult to remove electrons.

Within a group, the ionization energy decreases as the size of the atom gets larger. On the graph, we see that the ionization energy increases as we go up the group to smaller atoms. In this situation, the first electron removed is farther from the nucleus as the atomic number (number of protons) increases. Being farther away from the positive attraction makes it easier for that electron to be pulled off.

 

Ionic Radius

The ionic radius for an atom is measured in a crystal lattice, requiring a solid form for the compound . These radii will differ somewhat depending upon the technique used. Usually X-ray crystallography is employed to determine the radius for an ion.

 

Comparison of ion sizes to atom sizes for Groups 1, 2, 13, 16 and 17. The atoms are shown in gray. Groups 1, 2, and 13 are metals and form cations, shown in red. Groups 16 and 17 are nonmetals and form anions, shown in blue.

The removal of electrons always results in a cation that is considerably smaller than the parent atom . When the valence electron(s) are removed, the resulting ion has one fewer occupied principal energy level, so the electron cloud that remains is smaller. Another reason is that the remaining electrons are drawn closer to the nucleus because the protons now outnumber the electrons. One other factor is the number of electrons removed. The potassium atom has one electron removed to for the corresponding ion, while calcium loses two electrons.

The addition of electrons always results in an anion that is larger than the parent atom. When the electrons outnumber the protons, the overall attractive force that the protons have for the electrons is decreased. The electron cloud also spreads out because more electrons results in greater electron-electron repulsions. Notice that the group 16 ions are larger than the group 17 ions. The group 16 elements each add two electrons while the group 17 elements add one electron per atom to form the anions.

 

Electronegativity

Valence electrons of both atoms are always involved when those two atoms come together to form a chemical bond . Chemical bonds are the basis for how elements combine with one another to form compounds. When these chemical bonds form, atoms of some elements have a greater ability to attract the valence electrons involved in the bond than other elements. Electronegativity is a measure of the ability of an atom to attract electrons when the atom is part of a compound . Electronegativity differs from electron affinity because electron affinity is the actual energy released when an atom gains an electron.

Electronegativity is not measured in energy units, but is rather a relative scale. All elements are compared to one another, with the most electronegative element , fluorine, being assigned an electronegativity value of 3.98. Fluorine attracts electrons better than any other element. The table below shows the electronegativity values for the elements.

The electronegativity scale was developed by Nobel Prize winning American chemist Linus Pauling. The largest electronegativity (3.98) is assigned to fluorine and all other electronegativities measurements are on a relative scale.

Since metals have few valence electrons , they tend to increase their stability by losing electrons to become cations. Consequently, the electronegativities of metals are generally low. Nonmetals have more valence electrons and increase their stability by gaining electrons to become anions. The electronegativities of nonmetals are generally high.

 

Trends

Electronegativities generally increase from left to right across a period. This is due to an increase in nuclear charge. Alkali metals have the lowest electronegativities, while halogens have the highest. Because most noble gases do not form compounds, they do not have electronegativities. Note that there is little variation among the transition metals . Electronegativities generally decrease from top to bottom within a group due to the larger atomic size.

Of the main group elements, fluorine has the highest electronegativity (EN = 4.0) and cesium the lowest (EN = 0.79). This indicates that fluorine has a high tendency to gain electrons from other elements with lower electronegativities. We can use these values to predict what happens when certain elements combine. The following video shows this.

When the difference between electronegativities is greater than ~1.7, then a complete exchange of electrons occurs. Typically this exchange is between a metal and a nonmetal. For instance, sodium and chlorine will typically combine to form a new compound and each ion becomes isoelectronic with its nearest noble gas . When we compare the EN values, we see that the electronegativity for Na is 0.93 and the value for Cl is 3.2. The absolute difference between ENs is |0.93 - 3.2| = 2.27. This value is greater than 1.7, and therefore indicates a complete electron exchange occurs.

 

Putting It Together

Let us revisit the phenomenon:

Remember the Alkali metals from the beginning of the chapter? You learned that alkali metals react quickly with air and water.

1. Do elements in the same group have the same properties? Why or why not?
2. Do elements in the same period have the same properties? Why or why not?