MCAT General Chemistry Review
Part I Review
Chapter 2: The Periodic Table
The pharmacological history of lithium is an interesting window into the scientific and medical communities’ attempt to take advantage of the chemical and physical properties of an element for human benefit. By the mid-1800s, the medical community was showing great interest in theories that linked uric acid to a myriad of maladies. When it was discovered that solutions of lithium carbonate dissolved uric acid, therapeutic preparations containing the lithium carbonate salt became popular. Even nonmedical companies tried to profit from lithium’s reputation as a cure-all by adding it to their soft drinks.
Eventually, fascination with theories of uric acid wore off, and lithium’s time in the spotlight seemed to be coming to an end. Then, in the 1940s, doctors began to recommend salt-restricted diets for cardiac patients. Lithium chloride was made commercially available as a salt (sodium chloride) substitute. Unfortunately, lithium is quite toxic at fairly low concentrations, and when medical literature in the late 1940s reported several incidents of severe poisonings and multiple deaths—some associated with only minor lithium overdosing—U.S. companies voluntarily withdrew all lithium salts from the market. Right around this time, Australian psychiatrist John Cade proposed the use of lithium salts for the treatment of mania. Cade’s clinical trials were quite successful. In fact, his use of lithium salts to control mania was the first instance of successful medical treatment of a mental illness—and lithium carbonate became commonly prescribed in Europe for manic behavior. Not until 1970 did the U.S. Food and Drug Administration finally approve the use of lithium carbonate for manic illnesses.
Lithium (Li) is a chemical element with atomic number 3. It is an alkali metal, very soft, and under standard conditions it is the least-dense solid element, with a specific gravity of 0.53. Lithium is so reactive that it does not naturally occur on earth in its elemental form, being found only in various salt compounds.
Why would medical scientists pay attention to this particular element? What would make medical scientists believe that lithium chloride would be a good substitute for sodium chloride for patients on salt-restricted diets? The answers lie in the periodic table.
The Periodic Table
In 1869, Russian chemist Dmitri Mendeleev published the first version of his periodic table, in which he showed that ordering the known elements according to atomic weight produced a pattern of periodically recurring physical and chemical properties. Since then, the periodic table of the elements has been revised, using the work of physicist Henry Moseley, to organize the elements on the basis of increasing atomic number rather than atomic weight. Using this revised table, the properties of certain elements that had not yet been discovered were predicted. Experimentation later confirmed a number of these predictions. The periodic table puts into visual representation the principle of the periodic law: The chemical and physical properties of the elements are dependent, in a periodic way, upon their atomic numbers.
The modern periodic table arranges the elements into periods (rows) and groups (columns), also known as families. There are seven periods, representing the principal quantum numbers n = 1 through n = 7. Each period is filled sequentially, and each element in a given period has one more proton and one more electron (in the neutral state) than the element to its left. Groups or families include elements that have the same electronic configuration in their valence shell, which is the outermost 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 of all the electrons in the atom. 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 other elements (by way of the valence shells of the other elements); 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. The Roman numeral is combined with the letter A or B to separate the elements into two larger classes. The A elements are known as the representative elements and include groups IA, IIA, IIIA, IVA, VA, VIA, VIIA, and 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 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, d, and f subshells. For the representative elements, the Roman numeral and the letter designation determine the electron configuration. For example, an element in Group VA will have five valence electrons and a valence electron configuration of s2p3. The MCAT does not require you to know the corresponding association between Roman numerals and valence electron configuration for the nonrepresentative elements, however. The use of Roman numerals and letters to identify a particular family is confusing, because European and North American scientists traditionally have used the Roman numeral–letter system in different ways. In light of this, IUPAC developed and recommends a group identification system using Arabic numbers, 1–18, starting with the alkali metals on the left and ending with the noble gases on the right.
Periodic Properties of the Elements
We hope that it goes without saying that the MCAT will not expect you to have memorized the entire periodic table. Those of you with biology backgrounds may need the services of a Sherpa to find any element beyond the fourth period. The even less adventurous among you may never have ventured past chlorine! Fortunately, the periodic table is a guide unto itself, sort of a self-referencing GPS for all the elements. Remember, the modern table is organized in such a way to represent visually the periodicity of chemical and physical properties of the elements. The periodic table, then, can provide you with a tremendous amount of information that otherwise would have to be memorized. While you do not need to “memorize” the periodic table for the MCAT (or ever), you absolutely need to understand the trends within the periodic table that will help you predict the chemical and physical behavior of any element you encounter on the MCAT (and in your medical career).
Don’t try to memorize the periodic table. You will have access to it on Test Day. Do understand its configuration and trends so that you can use it efficiently to get a higher score!
A few basic facts to keep in mind, before we examine the trends in detail: First, as we’ve already mentioned, as you move from element to element, 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 it. This causes the electron cloud, the “outer boundary” of which is 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), which is a measure of the net positive charge experienced by the outermost electrons. For elements within the same period, then, Zeff increases from left to right.
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 you go down a group, the increase in the shielding effect of the additional insulating layer of inner shell electrons negates the increase in the positivity of the nucleus (the nuclear charge). Thus, the Zeff is more or less constant among the elements within a given group. In spite of this, the valence electrons are held less tightly to the nucleus due to the increased separation between them.
Third, and finally, we can generally say that elements behave in such a way to gain or lose electrons so as to achieve the stable octet formation possessed by the inert (noble) gases (Group VIII or Group 18). We will soon learn (Chapter 3) that this so-called “octet rule” is hardly a rule at all, as a far greater number of elements can be exceptions to this rule than elements that always follow it. Nevertheless, for now, let’s just keep in mind that elements, especially the ones that have biological roles, tend to want to have eight electrons in their valence shell.
These three facts can be your guiding principles as you work toward an understanding of the more particular trends demonstrated among the elements organized as they are in the periodic table. In fact, all you really need to remember is the trend for effective nuclear charge across a period and the impact of increasing the number of inner shells down a group in order to correctly “derive” all the other trends.
If we imagine one atom of any element to essentially be a little cloud of electrons with a dense core of protons and neutrons, then the atomic radius of an element is equal to one-half the distance between the centers of two atoms of that element that are just touching each other. (We can’t measure atomic radius by examining a single atom because the electrons are constantly moving around and it becomes impossible to mark the outer boundary of the electron cloud.) 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 inner-shell electrons (which act as insulation between the nucleus and the valence electrons) remains constant, the increasing positive charge of the nucleus holds the outer electrons more closely and more tightly. The Zeffincreases left to right across a period, and as a result, atomic radius decreases from left to right across a period.
As an electron and a proton get farther away from one another, it becomes easier to pull them apart. This will help us understand all the trends with respect to the radius.
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 Zeff remains essentially constant, atomic radius increases in a group from top to bottom. In summary, within each group, the largest atom will be at the bottom, and within each period, the largest atom will be within Group IA (Group 1).
Ionization energy (IE), also known as ionization potential, is the energy required to remove an electron completely from a gaseous atom or ion. Removing an electron from an atom always requires an input of energy (it’s always endothermic; Chapter 6). Anybody who has ever volunteered to make fundraising calls for a nonprofit organization knows that considerable energy must be invested in convincing a person to give a charitable donation, no matter how charitable the potential donor may be. Some are easier to convince than others are, but in the end, every donation is secured by a lot of hard work, effective communication, and skillful playing of the guilt and pity cards. The higher the atom’s Zeff or the closer the valence electrons are to the nucleus, the more tightly they are bound to the atom. This makes it more difficult to remove one or more electrons, so the ionization energy increases. Thus, ionization energy increases from left to right across a period and decreases in a group from top to bottom. Furthermore, the subsequent removal of a second or third electron requires increasing amounts of energy, because the removal of more than one electron necessarily means that the electrons are being removed from an increasingly cationic 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 to form the divalent cation is called the second ionization energy, and so on. For example:
Mg(g) Mg+(g) + e- First Ionization Energy = 7.646 eV
Mg+(g) Mg2+(g) + e- Second Ionization Energy = 15.035 eV
Elements in Groups I and II (Groups 1 and 2) have such relatively low ionization energies that they are called the active metals. The active metals never exist naturally in their neutral elemental (native) forms; they are always found in ionic compounds, minerals, or ores. The loss of one electron (from the alkali metals) or the loss of two electrons (from the alkaline earth metals) results in the formation of a stable, filled valence shell. As you might imagine from the trend, the Group VIIA (Group 17) elements, the halogens, are a miserly group of penny pinchers and aren’t willing to give up their electrons to anybody. In fact, in their monatomic ion form, they are found only as anions, having greedily taken one electron from another atom to complete their octets. As you might guess, the halogens have very large ionization energies and the smaller the halogen atom, the higher the ionization energy.
The only group less willing to give up their valence electrons is the inert elements (noble gases). They already have a very stable electron configuration and are unwilling to disrupt that stability by losing an electron. Inert gases are among the elements with the highest ionization energies.
The greedy halogens are among the worst of the bunch of elements that tend to hoard their electrons toward themselves. These elements also tend to be very anxious to gain the number of electrons necessary to complete their octets. Like nervous little squirrels frantically running around in search of nuts to pack into their accommodating cheek pouches, these elements go in search of other atoms that are willing to give up their electrons. When a gaseous atom of a particular elemental identity gains one or more electrons to complete its octet, it relaxes and breathes a sigh of relief. This “sigh of relief ” is a release of a quantity of energy called the electron affinity. Because energy is released when an atom gains an electron, we can describe this process as exothermic. By convention, the electron affinity is reported as a positive energy value, even though by the conventions of thermodynamics, exothermic processes have negative energy changes. Regardless of the sign, just remember that electron affinity is released energy. The stronger the electrostatic pull (that is, 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 period from top to bottom. Groups IA and IIA (Group 1 and 2) have very low electron affinities, preferring rather to give up one or two electrons, respectively, to achieve the octet configuration of the prior noble gas. 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 immediately following noble gases in Group VIIIA (Group 18). Although the noble gases are the group of elements farthest to the right and would be predicted to have the highest electron affinities according to the trend, they actually have electron affinities on the order of zero, since they already possess a stable octet and cannot readily accept an electron. Elements of other groups generally have low electron affinity values.
To recall the various trends, remember this: Cesium, Cs, is the largest, most metallic, and least electronegative of all naturally occurring elements. It also has the smallest ionization energy and the least exothermic electron affinity.
In contrast to cesium, fluorine (F ) is the smallest, most electronegative element. It also has the largest ionization energy and most exothermic electron affinity.
No, we are not referring to pessimistic electrons. 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 greater is its attraction for bonding electrons. 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 electronegativity value for any element is not measured directly 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 for fluorine, the most electronegative element. Electronegativity increases across a period from left to right and decreases in a period from top to bottom. Figure 2.1Summarizes the behavior of atoms in terms of the periodic table.
Electronegativity might better be called “nuclear positivity.” it is a result of the nucleus’ attraction for electrons; that is, the Zeff perceived by the electrons in a bond.
Types of Elements
It’s often been said that birds of a feather flock together, and this is no less true for the elements. When we consider the trends of chemical reactivity and physical properties taken together, we begin to identify whole clans, if you will, of elements that share sets of similarities. These larger collections of elements, which span groups and periods, are divided into three categories: metals, nonmetals, and metalloids (semimetals).
Note: Atomic radius is always opposite the other trends.
Metals, found both on the left side and in the middle of the periodic table, include the active metals, the transition metals, and the lanthanide and actinide series of elements. Metals are shiny solids, except for mercury, which is a liquid under standard state conditions. They generally have high melting points and densities, but there are exceptions, such as lithium, which has a density that is 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 drawn into wires is calledductility. At the atomic level, low Zeff, low electronegativity (high electropositivity), large atomic radius, and low ionization energy define metals. These characteristics make it fairly easy for metals to give up one or more electrons. Many of the transition metals, for example, are known to have twooxidation states, and some have more than that. Because the valence electrons of all metals are only loosely held to their atoms, they are essentially free to move, which makes metals generally good conductors of heat and electricity (some are better than others). 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 found in the f subshell. Some transition metals—such as copper, nickel, silver, gold, palladium, and platinum—are relatively nonreactive, a property that makes them ideal substances for the production of coins and jewelry.
The effective nuclear charge, Zeff, can explain all periodic trends as well as chemical properties.
Nonmetals are found on the upper right side of the periodic table. The metals claim that the nonmetals are jealous of them for their shiny hair and sparkly personalities. The nonmetals scoff back, yet quietly steal electrons from them. Nonmetals are generally brittle in the solid state and show little or no metallic luster. They have high ionization energies, electron affinities, and electronegativities; have small atomic radii; and are usually poor conductors of heat and electricity. Nonmetals are less unified in their chemical and physical properties than are the metals. They are separated from the metals by a diagonal band of elements called the metalloids.
The metalloids are also called the semimetals because they possess characteristics that are between those of metals and nonmetals. The electronegativities and ionization energies of the metalloids lie between those of metals and nonmetals. Their physical properties, such as densities, melting points, and boiling points, vary widely and can be combinations of metallic and nonmetallic characteristics. For example, silicon has a metallic luster but is brittle and a poor conductor. The particular reactivity of the metalloids is dependent upon the elements with which they are reacting. Boron (B), for example, behaves as a nonmetal when reacting with sodium (Na) and as a metal when reacting with fluorine (F). The elements classified as metalloids form a “staircase” on the periodic table and include boron, silicon, germanium, arsenic, antimony, tellurium, and polonium.
Metalloids share some properties with metals, and others with nonmetals. For instance, metalloids make good semiconductors due to their electrical conductivity.
The Chemistry of Groups
The alkali metals, Group IA (Group 1), possess most of the classic physical properties of metals, except that their densities are lower than those of other metals (as is true of lithium). The alkali metals have only one loosely bound electron in their outermost shells, and their Zeff values are very low, giving them the largest atomic radii of all the elements in their respective periods. The very low Zeff values also result in low ionization energies, low electron affinities, and low electronegativities, and these atoms easily lose one electron to form univalent cations. They react very readily with nonmetals, especially halogens.
Alkali and alkaline earth metals are both metallic in nature because they both lose electrons easily from the s-orbital of their valence shells.
ALKALINE EARTH METALS
The alkaline earth metals, Group IIA (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 so have 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.
The halogens, Group VIIA (Group 17), are highly reactive nonmetals with seven valence electrons. They are rather “desperate” to complete their octets by each gaining an additional electron. The halogens are highly variable in their physical properties. For instance, the halogens range from gaseous (F2 and Cl2) to liquid (Br2) to solid (I2) at room temperature. 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 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).
Halogens are seen often on the MCAT. Remember that they only need one more electron to become “noble” (have that full valence shell).
The noble gases, Group VIIIA (Group 18), are also known as the inert gases because they have very low chemical reactivities as a result of their filled valence shells. They have high ionization energies, little or no tendency to gain or lose electrons, and no real electronegativities. They are essentially snobby elements, as they refuse to mingle with the hoi polloi. After all, they already have everything they need. The noble gases have low boiling points, and all exist as gases at room temperature.
The transition elements, Groups IB to VIIIB (Groups 3 to 12), are all considered 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 are progressively filling the d subshell orbitals in the valence shell. 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 various numbers of electrons from the s- and d-orbitals of the valence shell. For instance, copper (Cu), in Group 1B (Group 11), can exist in either the +1 or the +2 oxidation state, and manganese (Mn), in Group VIIB (Group 7), can have 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 and partially ionic compounds. The dissolved ions can form complex ions either with molecules of water (hydration complexes) or with nonmetals, forming highly colored solutions and compounds (e.g., CuSO4·5H2O), and this complexation may enhance the relatively low solubility of certain compounds. For example, AgCl is insoluble in water but quite soluble in aqueous ammonia due to the formation of the complex ion [Ag(NH3)2]+. 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 higher d sublevel. The frequencies not absorbed (known as the subtraction frequencies) give the complexes their characteristic colors.
Transition metals are present in biological systems and are therefore often seen on the MCAT (think iron in hemoglobin). You don’t need to memorize them, but be able to use your knowledge from these first two chapters to understand how the transition metals ionize and act.
Now that we have completed our review of the periodic table of the elements, commit to understanding (not just to memorizing) the trends of physical and chemical properties that will allow you to answer quickly the questions on the MCAT. You will find, as you progress through the chapters of this book, that your foundational understanding of the elements will help you develop a richer, more nuanced understanding of their general and particular behaviors. Topics in general chemistry that may have given you trouble in the past will be understandable from the perspective of the behaviors and characteristics that you have reviewed here.
CONCEPTS TO REMEMBER
The periodic table of the elements organizes the elements according to their atomic numbers and reveals a repeating pattern of similar chemical and physical properties. Elements in the same row are in a period, while those elements in a column are in a group. Elements in the same period have the same principal energy level, n. Elements in the same group have the same valence shell electron configuration.
The valence electrons are those located in the outer shell and/or are available for interaction (bonding) with other atoms. The representative elements have their valence electrons in either s- or s- and p-orbitals. The nonrepresentative elements (the transition elements) have their valence electrons either in s- and d- or in s-, d-, and f-orbitals.
Effective nuclear charge (Zeff) is the net positive charge experienced by electrons in the valence shell. Zeff increases from left to right across a period, with little change in value from top to bottom in a group. Valence electrons become increasingly separated from the nucleus as the principal energy level, n, increases from top to bottom in a group. These two trends are the basis for all the other trends exhibited by the elements in the periodic table.
Atomic radius decreases from left to right across a period and increases from top to bottom in a group.
Ionization energy (IE) is the amount of energy necessary to remove an electron from the valence shell. It increases from left to right across a period and decreases from top to bottom in a group.
Electron affinity is the amount of energy released when an atom gains an electron in its valence shell. It increases from left to right across a period and decreases from top to bottom in a group.
Electronegativity is a measure of the attractive force that an atom in a chemical bond will exert on the electron pair of the bond. It increases from left to right across a period and decreases from top to bottom in a group.
There are three general classes of elements:
—The metals, located on the left and middle of the periodic table, including the active metals and the transition metals
—The nonmetals, located in the upper right side of the periodic table, including hydrogen, carbon, oxygen, nitrogen, and phosphorus, among others
—The metalloids or semimetals, located in a staircase formation between the metals and nonmetals, with qualities and behaviors that are combinations of those of the metals and nonmetals
The alkali and alkaline earth metals are the most reactive of all metals; they exist only in their ionic forms, having given up one or two electrons, respectively, in order to achieve the electronic configuration of the prior noble gas. The transition metals are less reactive, and many can have two or more oxidation states.
The halogens are very reactive nonmetals and are highly electronegative. They need only one electron to complete their octets and are naturally found only in the anionic state. The noble gases are the least reactive of all the elements because they have the stable octet in their valence shell. They have very high ionization energies and virtually nonexistent electronegativities.
1. Lithium and sodium have similar chemical properties. For example, both can form ionic bonds with chloride. Which of the following best explains this similarity?
A. Both lithium and sodium ions are positively charged.
B. Lithium and sodium are in the same group of the periodic table.
C. Lithium and sodium are in the same period of the periodic table.
D. Both lithium and sodium have low atomic weights.
2. Carbon and silicon, elements used as the basis of biological life and synthetic computing, respectively, have some similar chemical properties. Which of the following describes a difference between the two elements?
A. Carbon has a smaller atomic radius than silicon.
B. Silicon has a smaller atomic radius than carbon.
C. Carbon has fewer valence electrons than silicon.
D. Silicon has fewer valence electrons than carbon.
3. What determines the length of an element’s atomic radius?
I. The number of valence electrons
II. The number of electron shells
III. The number of neutrons in the nucleus
A. I only
B. II only
C. I and II only
D. I, II, and III
4. Ionization energy contributes to an atom’s chemical reactivity. An accurate ordering of ionization energies, from lowest ionization energy to highest ionization energy, would be
A. Be, first ionization energy < Be, second ionization energy < Li, first ionization energy.
B. Be, second ionization energy < Be, first ionization energy < Li, first ionization energy.
C. Li, first ionization energy < Be, first ionization energy < Be, second ionization energy.
D. Li, first ionization energy < Be, second ionization energy < Be, first ionization energy.
5. Selenium is often an active component of treatments for scalp dermatitis. What type of element is selenium?
6. The properties of atoms can be predicted, to some extent, by their location within the periodic table. Which property or properties increase in the direction of the arrows shown?
II. Atomic radius
III. First ionization energy
A. I only
B. II only
C. I and III only
D. II and III only
7. Metals are often used for making wires that conduct electricity. Which of the following properties of metals explains why?
A. Metals are malleable.
B. Metals have high electronegativities.
C. Metals have valence electrons that can move freely.
D. Metals have high melting points.
8. Which of the following is an important property of the set of elements shaded in the periodic table shown?
A. These elements are the best electrical conductors in the periodic table.
B. These elements form divalent cations.
C. The second ionization energy for these elements is lower than the first ionization energy.
D. The atomic radii of these elements decrease as one moves down the column.
9. When dissolved in water, which of the following ions is most likely to form a complex ion with H2O?
10. How many valence electrons are present in elements in the third period?
C. The number decreases as the atomic number increases.
D. The number increases as the atomic number increases.
11. Which of the following elements has the highest electronegativity?
12. Of the four atoms depicted here, which has the highest electron affinity?
13. Which of the following is true of an atom with a large atomic radius?
A. The atom is most likely to be on the right side of the periodic table.
B. The atom is likely to have a high second ionization energy.
C. The atom is likely to have low electronegativity.
D. The atom is likely to form ionic bonds.
14. Which of the following atoms/ions has the largest effective nuclear charge?
15. Why do halogens often form ionic bonds with alkaline earth metals?
A. The alkaline earth metals have much higher electron affinities than the halogens.
B. By sharing electrons equally, the alkaline earth metals and halogens both form full octets.
C. Within the same row, the halogens have smaller atomic radii than the alkaline earth metals.
D. The halogens have much higher electron affinities than the alkaline earth metals.
16. What is the outermost orbital of elements in the third period?
Small Group Questions
1. Mercury (Hg) exists as a liquid at room temperature. Why, then, is it classified as a metal? What metallic properties might it possess?
2. The transition metals in group VIIIB could theoretically have eight different oxidation states. In reality, that does not hold true. Why not?
Explanations to Practice Questions
To answer this question, one must first recall that the periodic table is organized with periods (rows) and groups (columns). This method of organization allows elements to be organized such that some chemical properties can be predicted based on an element’s position in the table. Groups (columns) are particularly significant because they represent sets of elements with the same outer electron configuration. In other words, all elements within the same group will have the same configuration of valence electrons, which in turn will dictate many of the chemical properties of those similar elements. Although (A) is true, the fact that both ions are positively charged does not explain the similarity in chemical properties as effectively as does answer (B); most other metals, whether or not they are similar to lithium and sodium, produce positively charged ions. (C) is not true, because periods are rows and lithium and sodium are in the same column. Finally, although lithium and sodium do have relatively low atomic weights, so do several other elements that do not share those same properties.
This question assesses understanding of a key periodic trend: atomic radii. As one moves from left to right across a period (row), atomic radii decrease. This occurs because as more protons are added to the nucleus and more electrons are added within the same shell, there is no increased shielding between the protons and electrons, but there is increased attractive electrostatic force. This effect decreases the atomic radius. In contrast, as one moves from top to bottom down a group (column), extra electron shells accumulate, despite the fact that the valence configurations remain identical. These extra electron shells provide shielding between the positive nucleus and the outermost electrons, decreasing the electrostatic forces and increasing the atomic radius. Because carbon and silicon are in the same group, and silicon is farther down the periodic table, it will have a larger atomic radius because of its extra electron shell. (C) and (D) are incorrect because all elements in the same group have the same number of valence electrons.
Atomic radius is determined by multiple factors. Of the choices given, the number of valence electrons does have an impact on the atomic radius. As one moves across a period (row), protons and valence electrons are added, and the electrons are more strongly attracted to the central protons. This attraction tightens the atom, shrinking the atomic radius. The number of electron shells is also significant, as demonstrated by the trend when moving down a group (column). As more electron shells are added that separate the positively charged nucleus from the outermost electrons, the electrostatic forces are weakened, and the atomic radius increases. The number of neutrons, III, is irrelevant, because it does not impact these attractive forces. Statements I and II are correct, and III is incorrect.
Ionization energy is related to the same set of forces that explains atomic radius, as well as the rules governing maintenance of a full valence shell octet. The first set of rules dictates that the stronger the attractive forces between the outer electron (the electron to be ionized) and the positively charged nucleus, the more energy will be required to ionize. As a result, strong attractive forces, which make the atomic radius smaller toward the right of a period or the top of a group, will also increase the first ionization energy. With this information alone, one could guess that ionization energies for beryllium (Be) should be higher than those for Li (lithium), eliminating (A) and (B). How should we choose between the two ionization energies for beryllium? The first ionization energy is usually dramatically lower than the second. This property holds true for the same reasons previously discussed. For example, upon removing one electron from beryllium, the ion is Be1+, which has one more proton in its nucleus than it has electrons surrounding it. Thus, there is a heightened electrostatic force between the positive nucleus and the now less-negative electron cloud, meaning that all remaining electrons will be held more tightly than before. Removing a second electron will be even more difficult than the initial electron removal, making the second ionization energy higher than the first. Furthermore, the first ionization energy for Li is 520.2 kJ/ mol, the first ionization energy for Be is 899.5 kJ/mol, and the second ionization energy for Be is 1,757.1 kJ/mol.
Selenium is on the right side of the periodic table (atomic number 34), too far right to be a metal or metalloid. However, it does not lie far enough to the right to fall under column 7A, which would classify it as a halogen. It is a nonmetal.
The trend in the periodic table demonstrated by the figure is correct for increasing electronegativity and first ionization energy. Electronegativity describes how strong an attraction an element will have for electrons in a bond. A nucleus with a stronger electrostatic pull due to its positive charge will have a higher electronegativity. An arrow pointing toward the right represents this because effective nuclear charge increases toward the right side of a period. This mirrors the trend for ionization energies, because a stronger nuclear pull will also lead to increased first ionization energy, as the forces make it more difficult to remove an electron. The vertical arrow can be explained by the size of the atoms. As size decreases, the proximity of the outermost electrons to the positive inner nucleus increases, making the positive charge more effective at attracting new electrons in a chemical bond (which leads to higher electronegativity). Similarly, the more effective the positive nuclear charge, the higher the first ionization energy. Because I and III are both correct, we can choose answer choice (C).
All four descriptions of metals are true, but the most significant property that contributes to their ability to conduct electricity is the fact that they have valence electrons that can move freely (C). Large atomic radii, low ionization energies, and low electronegativities all contribute to the ability of metals’ outermost electrons to be easily removed, but it’s the free movement of electrons that actually conducts the electricity.
This block represents the alkaline earth metals, which form divalent cations, or ions with a +2 charge. All of the elements in Group IIA have two electrons in their outermost s-orbital. Because loss of these two electrons would then leave a full octet as the outermost shell, becoming a divalent cation is a stable configuration for all of the alkaline earth metals. Although some of these elements might be great conductors, it’s an exaggeration to say that they are the best ones on the periodic table, so (A) is incorrect. (C) is also incorrect because although forming a divalent cation is a stable configuration for the alkaline earths, the second ionization energy is still always higher than the first due to the increased positive nuclear charge when compared with the outer negative charge from the electrons. Finally, (D) is incorrect because atomic radii increase when moving down a group of elements because the number of electron shells increases.
Iron, Fe2+, is a transition metal. Transition metals can often form more than one ion. Iron, for example, can be Fe2+ or Fe3+. The transition metals, in these various oxidation states, can often form hydration complexes (complexes with water). Part of the significance of these complexes is that when a transition metal can form a complex, its solubility within the complex solvent will increase. The other ions given might dissolve readily in water, but because none of them are transition metals, they won’t form complexes.
This question is simple if you recall that “periods” name the horizontal rows of the periodic table, while “families” refer to its columns. Within the same period, an additional valence electron is added with each step towards the right side of the table.
This question requires knowledge of the trends of electronegativity within the periodic table. Electronegativity increases as one moves from left to right across periods for the same reasons that effective nuclear force increases. Electronegativity decreases as one moves down the periodic table, because there are more electron shells separating the nucleus from the outermost electrons. The noble gases, however, also have extremely low electronegativities because they already have full valence shells and do not desire additional electrons. The most electronegative atom in the periodic table is fluorine. The answer choice closest to fluorine is (B), chlorine. Although iodine (D) will be fairly electronegative, its higher atomic radius and position farther down on the periodic table make it less electronegative than Cl. The remaining answer choices, Mg and Li, are elements with very low electronegativities. Because they have only two and one valence electrons, respectively, they are more likely to lose these electrons in a bond than to gain electrons; the loss of electrons would leave them with a full octet. Metals like these are often called electropositive.
Electron affinity is related to several factors, including atomic size (radius) and filling of the valence shell. As atomic radius increases, the distance between the inner protons in the nucleus and the outermost electrons increases, thereby decreasing the attractive forces between protons and electrons. Additionally, as more electron shells are added from period to period, these shells shield the outermost electrons increasingly from inner protons. As a result, increased atomic radius will lead to lower electron affinity. Because atoms are in a low-energy state when their outermost valence electron shell is filled, atoms needing only one or two electrons to complete this shell will have high electron affinities. In contrast, atoms with already full valence shells (a full octet of electrons) will have very low electron affinities, because adding an extra electron would require a new shell. With this information, it is clear that (C) and (D) will likely have lower electron affinities than (A) and (B) because there is an extra electron shell “shielding” the nucleus from the outer electrons. Answer (A) is incorrect because its valence electron shell is already full with a complete octet, granting it extremely low electron affinity. Finally, (B) has one electron missing from its outermost shell, as does (D). This valence electron configuration is conducive to wanting to accept electrons readily—or to having a high electron affinity. (B) is the configuration of chlorine, while (D) is bromine. (B) is a better answer than (D), however, because the additional shell of electrons shielding the nucleus in (D) will decrease its electron affinity when compared with (B).
Electronegativity is the only property listed that has a consistent inverse correlation with atomic radius. Highly electronegative atoms hold bonding electrons tightly, while atoms with low electronegativity hold bonding electrons loosely. In atoms with large atomic radii, the distance between the outermost electrons used for bonding and the central positively charged nucleus is large. This increased distance means that the positively charged nucleus has little ability to attract new, bonding electrons toward it. In comparison, if the atomic radius is small, the force from the positively charged protons will have a stronger effect, because the distance through which they have to act is decreased. (A) is incorrect because atomic radius decreases when moving from left to right across periods in the periodic table. (B) is incorrect because atomic radius alone does not give enough information for one to ascertain the second ionization energy; it is also significant to consider the valence electron configuration. Additionally, all atoms have high second ionization energies. Finally, there is insufficient information supporting (D).
The effective nuclear charge refers to the strength with which the protons in the nucleus can “pull” additional electrons. This phenomenon helps to explain electron affinity, electronegativity, and ionization energy. In Cl, the nonionized chlorine atom, the nuclear charge is balanced by the surrounding electrons: 17+/17-. The chloride ion, in contrast, has a lower effective nuclear charge, because there are more electrons than protons: 17+/18-. Next, elemental potassium also has a “balanced” effective nuclear charge: 19+/19-. K+, ionic potassium, has a higher effective nuclear charge than any of the other options do, because it has more protons than electrons: 19+/18-. Thus, the potassium ion, (D), is the correct answer.
Ionic bonds are bonds formed through unequal sharing of electrons. These bonds typically occur because the electron affinities of the two bonded atoms differ greatly. For example, the halogens have high electron affinities because adding a single electron to their valence shells would create full outer octets. In contrast, the alkaline earth metals have very low electron affinities and are more likely to be electron donors because the loss of two electrons would leave them with full outer octets. This marked difference in electron affinity is the best explanation for the formation of ionic bonds between these two groups. (A) states the opposite and is incorrect, because the halogens have high electron affinity and the alkaline earth metals have low affinity. (B) is incorrect because equal sharing of electrons is a classic description of covalent bonding, not ionic. (C) is a true statement but is not relevant to why ionic bonds form.
In the first period, all elements have only an s-orbital. In the second period onwards, a 2p-orbital is present. In the third period, we find 3s-, 3p-, and 3d-orbitals. Though 3d appears to be part of the fourth period, it still shares the same principal quantum number as 3s and 3p (n = 3) and is therefore still applicable.