2. The Periodic Table

The modern Periodic Table arranges the elements into periods (rows) and groups or families (columns), based on atomic number. There are seven periods, representing the principal quantum numbers n = 1 through n = 7 for the s- and p-block elements. Each period is filled sequentially, and each element in a given period has one more proton and one more electron than the element to its left (in their neutral states). Groups contain elements that have the same electronic configuration in their valence shell and share similar chemical properties.

The electrons in the valence shell, known as the valence electrons, are the farthest from the nucleus and have the greatest amount of potential energy. Their higher potential energy and the fact that they are held less tightly by the nucleus allows them to become involved in chemical bonds with the valence electrons of other atoms; thus, the valence shell electrons largely determine the chemical reactivity and properties of the element.

The Roman numeral above each group represents the number of valence electrons elements in that group have in their neutral state. The Roman numeral is combined with the letter A or B to separate the elements intotwo larger classes. The A elements are known as the representative elements and include groups IA through VIIIA.

The elements in these groups have their valence electrons in the orbitals of either s or p subshells. The B elements are known as the nonrepresentative elements and include both the transition elements, which have valence electrons in the s and d subshells, and the lanthanide and actinide series, which have valence electrons in the s and f subshells. For therepresentative elements, the Roman numeral and the letter designation determine the electron configuration.

Elements in Groups IA and IIA (Groups 1 and 2), such as lithium and beryllium, have such low ionization energies that they are called the active metals. The active metals do not exist naturally in their neutral forms; they are always found in ionic compounds, minerals, or ores.

The loss of one electron from the alkali metals (GroupIA) or the loss of two electrons from the alkaline earth metals (Group IIA) results in the formation of a stable, filled valence shell. In contrast, the Group VIIA (Group 17) elements—the halogens—do not typically give up their electrons. In fact, in their ionic form, they are generally anions.

Unlike atomic radii, ionic radii will require some critical thinking and Periodic Table geography to determine. In order to understand ionic radii, we must make two generalizations. One is that metals lose electrons and become positive, while nonmetals gain electrons and become negative.

The other is that metalloids can go in either direction, but tend to follow the trend based on which side of the metalloid line they fall on. Thus, silicon (Si) behaves more like a nonmetal, while germanium (Ge) tends to act like a metal. On the MCAT, these generalizations can also be inferred from information found in passages and questions, such as oxidation states in compounds.

The alkali metals (Group IA or Group 1) possess most of the classic physical properties of metals, except that their densities are lower than those of other metals (as described for lithium earlier in this chapter). The alkali metals have only one loosely bound electron in their outermost shells.

Their Zeff values are very low, giving them the largest atomic radii of all the elements in their respective periods. This low Z value also explains the other trends: low ionization energies, low electron affinities, and low electronegativities. Alkali metals easily lose one electron to form univalent cations, and they react readily with nonmetals—especially the halogens—as in NaCl.

The halogens (Group VIIA or Group 17) are highly reactive nonmetals with seven valence electrons. These elements are desperate to complete their octets by gaining one additional electron. The physical properties of this group are variable. At standard conditions, the halogens range from gaseous (F and Cl ) to liquid (Br ) to solid (I ) forms.

Their chemical reactivity is more uniform, and, due to their very high electronegativities and electron affinities, they are especially reactive toward the alkali and alkaline earth metals. Fluorine (F) has the highest electronegativity of all the elements. The halogens are so reactive that they are not naturally found in their elemental state but rather as ions (called halides) or diatomic molecules.

For nonmetals close to the metalloid line, their group number dictates that they require more electrons than other nonmetals to achieve the electronic configuration seen in Group VIIIA.

These nonmetals gain electrons while their nuclei maintain the same charge. Therefore, these nonmetals close to the metalloid line possess a larger ionic radius than their counterparts closer to Group VIIIA.

Second, as one moves down the elements of a given group, the principal quantum number increases by one each time. This means that the valence electrons are increasingly separated from the nucleus by a greater number of filled principal energy levels, which can also be called inner shells. The result of this increased separation is a reduction in the electrostatic attraction between the valence electrons and the positively charged nucleus.

These outermost electrons are held less tightly as the principal quantum number increases. As one goes down in a group, the increased shielding created by the inner shell electrons cancels the increased positivity of the nucleus. Thus, the Zeff is more or less constant among the elements within a given group. Despite this fact, the valence electrons are held less tightly to the nucleus as one moves down a group due to the increased separation between valence electrons and the nucleus. Third, elements can also gain or lose electrons in order to achieve a stable octet formation representative of the noble (inert) gases (Group VIIIA or Group 18). For now, keep in mind that elements, especially the ones that have biological roles, tend to be most stable with eight electrons in their valence shell.

The alkaline earth metals (Group IIA or Group 2), also possess many properties characteristic of metals. They share most of the characteristics of the alkali metals, except that they have slightly higher effective nuclear charges and thus slightly smaller atomic radii.

They have two electrons in their valence shell, both of which are easily removed to form divalent cations. Together, the alkali and alkaline earth metals are called the active metals because they are so reactive that they are not naturally found in their elemental (neutral) state.

Ionization energy (IE), also known as ionization potential, is the energy required to remove an electron from a gaseous species. Removing an electron from an atom always requires an input of heat, which makes it an endothermic process. The greater the atom's Z or the closer the valence electrons are to the nucleus, the more tightly bound they are. This makes it more difficult to remove one or more electrons, increasing the ionization energy.

Thus, ionization energy increases from left to right across a period and from bottom to top in a group. The subsequent removal of a second or third electron requires increasing amounts of energy because the removal of more than one electron means that the electrons are being removed from an increasingly cationic (positive) species. The energy necessary to remove the first electron is called the first ionization energy; the energy necessary to remove the second electron from the univalent cation (X ) to form the divalent cation (X ) is called the second ionization energy, and so on.

periodic law:

the chemical and physical properties of the elements are dependent, in a periodic way, upon their atomic numbers

KEY CONCEPT 3

Cs = largest, least electronegative, lowest ionization energy, least exothermic electron affinity F = smallest, most electronegative, highest ionization energy, most exothermic electron affinity

The chalcogens (Group VIA or Group 16) are an eclectic group of nonmetals and metalloids. While not as reactive as the halogens, they are crucial for normal biological functions. They each have six electrons in their valence electron shell and, due to their proximity to the metalloids, generally have small atomic radii and large ionic radii. Oxygen is the most important element in this group for many reasons; it is one of the primary constituents of water, carbohydrates, and other biological molecules.

Sulfur is also an important component of certain amino acids, and vitamins. Selenium also is an important nutrient for microorganisms and has a role in protection from oxidative stress. The remainder of this group is primarily metallic and generally toxic to living organisms. It is important to note that, at high concentrations, many of these elements—no matter how biologically useful—can be toxic or damaging.

Many of the transition metals (Group B elements) have two or more oxidation states (charges when forming bonds with other atoms). Because the valence electrons of all metals are only loosely held to their atoms, they are free to move, which makes metals good conductors of heat and electricity.

The valence electrons of the active metals are found in the s subshell; those of the transition metals are found in the d subshell; and those of the lanthanide and actinide series elements are in the f subshell. Some transition metals—copper, nickel, silver, gold, palladium, and platinum—are relatively nonreactive, a property that makes them ideal for the production of coins and jewelry.

These complex ions tend to associate in solution either with molecules of water (hydration complexes) or with nonmetals. This ability to form complexes contributes to the variable solubility of certain transition metal-containing compounds. For example, AgCl is insoluble in water but quite soluble in aqueous ammonia due to the formation of the complex ion.

The formation of complexes causes the d-orbitals to split into two energy sublevels. This enables many of the complexes to absorb certain frequencies of light—those containing the precise amount of energy required to raise electrons from the lower- to the higherenergy d-orbitals. The frequencies not absorbed (known as the subtraction frequencies) give the complexes their characteristic colors.

The noble gases (Group VIIIA or Group 18) are also known as inert gases because they have minimal chemical reactivity due to their filled valence shells. They have high ionization energies, little or no tendency to gain or lose electrons, and (for He, Ne, and Ar, at least), no measurable electronegativities.

The noble gases have extremely low boiling points and exist as gases at room temperature. Noble gases have found a commercial niche as lighting sources, due to their lack of reactivity.

Nonmetals are found predominantly on the upper right side of the Periodic Table. Nonmetals are generally brittle in the solid state and show little or no metallic luster. They have high ionization energies, electron affinities, and electronegativities, as well as small atomic radii and large ionic radii.

They are usually poor conductors of heat and electricity. All of these characteristics are manifestations of the inability of nonmetals to easily give up electrons. Nonmetals are less unified in their chemical and physical properties than the metals.

Before exploring the periodic trends, let's take stock of three key rules that control how valence electrons work in an atom. First, as we've already mentioned, as one moves from left to right across a period, electrons and protons are added one at a time. As the positivity of the nucleus increases, the electrons surrounding the nucleus, including those in the valence shell, experience a stronger electrostatic pull toward the center of the atom.

This causes the electron cloud, which is the outer boundary defined by the valence shell electrons, to move closer and bind more tightly to the nucleus. This electrostatic attraction between the valence shell electrons and the nucleus is known as the effective nuclear charge (Zeff ), a measure of the net positive charge experienced by the outermost electrons. This pull is somewhat mitigated by nonvalence electrons that reside closer to the nucleus. For elements in the same period, Zeff increases from left to right.

KEY CONCEPT 1

Alkali and alkaline earth metals are both metallic in nature because they easily lose electrons from the s subshell of their valence shells.

As we move across a period from left to right, protons and electrons are added one at a time to the atoms. Because the electrons are being added only to the outermost shell and the number of innershell electrons remains constant, the increasing positive charge of the nucleus pulls the outer electrons more closely inward and holds them more tightly. The Zeff increases left to right across a period, and as a result, atomic radius decreases from left to right across a period.

As we move down a group, the increasing principal quantum number implies that the valence electrons will be found farther away from the nucleus because the number of inner shells is increasing, separating the valence shell from the nucleus. Although the Z remains essentially constant, the atomic radius increases down a group. Within each group, the largest atom will be at the bottom, and within each period, the largest atom will be in Group IA (Group 1). For reference, the largest atomic radius in the Periodic Table belongs to cesium (Cs, 260 pm), and the smallest belongs to helium (He, 25 pm). Francium is typically not considered because it is exceptionally rare in nature.

KEY CONCEPT 2

Atomic radius refers to the size of a neutral element, while an ionic radius is dependent on how the element ionizes based on its element type and group number.

Separating the metals and nonmetals are a stair-step group of elements called the metalloids. The metalloids are also called semimetals because they share some characteristics with both metals and nonmetals. The electronegativities and ionization energies of the metalloids lie between those of metals and nonmetals. Their physical properties—densities, melting points, and boiling points—vary widely and can be combinations of metallic and nonmetallic characteristics.

For example, silicon (Si) has a metallic luster but is brittle and a poor conductor. The reactivities of the metalloids are dependent on the elements with which they are reacting. Boron (B), for example, behaves like a nonmetal when reacting with sodium (Na) and like a metal when reacting with fluorine (F). The elements classified as metalloids form a "staircase" on the Periodic Table and include boron, silicon, germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te), polonium (Po), and astatine (At).

The values for second ionization energies are disproportionally larger for Group IA monovalent cations (like Na ) but generally not that much larger for Group IIA or subsequent monovalent cations (like Mg ). This is because removing one electron from a Group IA metal results in a noble gas-like electron configuration.

Group VIIIA (Group 18) elements, or noble or inert gases, are the least likely to give up electrons. They already have a stable electron configuration and are unwilling to disrupt that stability by giving up an electron. Therefore, noble gases are among the elements with the highest ionization energies.

Halogens are the most "greedy" group of elements on the Periodic Table when it comes to electrons. By acquiring one additional electron, a halogen is able to complete its octet and achieve a noble gas configuration. This exothermic process expels energy in the form of heat. Electron affinity refers to the energy dissipated by a gaseous species when it gains an electron. Note the electron affinity is essentially the opposite concept from ionization energy. Because this is an exothermic process, ΔH has a negative sign; however, the electron affinity is reported as a positive number. This is because electron affinity refers to the energy dissipated: if of energy is released, and the electron affinity is The stronger the electrostatic pull (the higher the Zeff) between the nucleus and the valence shell electrons, the greater the energy release will be when the atom gains the electron. Thus, electron affinity increases across a period from left to right. Because the valence shell is farther away from the nucleus as the principal quantum number increases, electron affinity decreases in a group from top to bottom.

Groups IA and IIA (Groups 1 and 2) have very low electron affinities, preferring to give up electrons to achieve the octet configuration of the noble gas in the previous period. Conversely, Group VIIA (Group 17) elements have very high electron affinities because they need to gain only one electron to achieve the octet configuration of the noble gases (Group VIIIA or Group 18) in the same period. Although the noble gases would be predicted to have the highest electron affinities according to the trend, they actually have electron affinities on the order of zero because they already possess a stable octet and cannot readily accept an electron. Most metals also have low electron affinity values

For metals, the trend is similar but opposite. Metals closer to the metalloid line have more electrons to lose to achieve the electronic configuration seen in Group VIIIA. Because of this, the ionic radius of metals near the metalloid line is dramatically smaller than that of other metals.

Metals closer to Group IA have fewer electrons to lose and therefore experience a less drastic reduction in radius during ionization. Note that tellurium (Te) behaves as a nonmetal and boron (B) behaves as a metal; under varying conditions, these metalloids can have opposite behavior.

Metals are found on the left side and in the middle of the Periodic Table. They include the active metals, the transition metals, and the lanthanide and actinide series of elements. Metals are lustrous (shiny) solids, except for mercury, which is a liquid under standard conditions. They generally have high melting points and densities, but there are exceptions, such as lithium, which has a density about half that of water.

Metals have the ability to be deformed without breaking; the ability of metal to be hammered into shapes is called malleability, and its ability to be pulled or drawn into wires is called ductility. At the atomic level, a metal is defined by a low effective nuclear charge, low electronegativity (high electropositivity), large atomic radius, small ionic radius, and low ionization energy. All of these characteristics are manifestations of the ability of metals to easily give up electrons.

The transition elements (Groups IB to VIIIB or Groups 3 to 12) are considered to be metals and as such have low electron affinities, low ionization energies, and low electronegativities. These metals are very hard and have high melting and boiling points. They tend to be quite malleable and are good conductors due to the loosely held electrons that progressively fill the d-orbitals in their valence shells.

One of the unique properties of the transition metals is that many of them can have different possible charged forms, or oxidation states because they are capable of losing different numbers of electrons from the s- and d-orbitals in their valence shells. For instance, copper (Cu) can exist in either the +1 or the +2 oxidation state, and manganese (Mn) can exist in the +2, +3, +4, +6, or +7 oxidation state. Because of this ability to attain different positive oxidation states, transition metals form many different ionic compounds. These different oxidation states often correspond to different colors; solutions with transition metal-containing complexes are often vibrant,

KEY CONCEPT 4

Periodic Trends Left → Right Atomic radius ↓ Ionization energy ↑ Electron affinity ↑ Electronegativity ↑ Top → Bottom Atomic radius ↑ Ionization energy ↓ Electron affinity ↓ Electronegativity ↓ Note: Atomic radius is always opposite the other trends. Ionic radius is variable.

Think of an atom as a cloud of electrons surrounding a dense core of protons and neutrons. The atomic radius of an element is thus equal to one-half of the distance between the centers of two atoms of an element that are briefly in contact with each other.

The distance between two centers of circles in contact is akin to a diameter, making this radius calculation simple. The atomic radius cannot be measured by examining a single atom because the electrons are constantly moving around, making it impossible to mark the outer boundary of the electron cloud.

Electronegativity is a measure of the attractive force that an atom will exert on an electron in a chemical bond. The greater the electronegativity of an atom, the more it attracts electrons within a bond. Electronegativity values are related to ionization energies: the lower the ionization energy, the lower the electronegativity; the higher the ionization energy, the higher the electronegativity. The first three noble gases are exceptions: despite their high ionization energies, these elements have negligible electro-negativity because they do not often form bonds.

The electronegativity value is a relative measure, and there are different scales used to express it. The most common scale is the Pauling electronegativity scale, which ranges from 0.7 for cesium, the least electronegative (most electropositive) element, to 4.0 for fluorine, the most electronegative element. Electronegativity increases across a period from left to right and decreases in a group from top to bottom.