Study Guide: THE PERIODIC TABLE

AND PERIODIC LAW

CHEMISTRY STUDY GUIDE 6: PERIODIC LAW P.1

TABLE OF CONTENTS

I. THE PERIODIC TABLE - ORGANIZATION .................................................. 2 A. Mendeleev ............................................................................................................... 2 B. Henry Moseley ........................................................................................................ 3 C. Other Periodic Tables ............................................................................................. 3

II. MODERN PERIODIC TABLE ........................................................... 4 A. Groups, Families & Periods, and Group Numbers ................................................. 4 B. Classification of Groups ......................................................................................... 4

1. Classification – Metals, Nonmetals, Metalloids ....................................................... 4 2. Classification – Blocks ................................................................................................ 5 3. Classification – Groups or Families .......................................................................... 5

III. THE PERIODIC TABLE - TRENDS ......................................... 8 A. Terms ...................................................................................................................... 8

Valence Electrons ....................................................................................................... 8 Octet Rule: .................................................................................................................. 8 Ion: ............................................................................................................................... 8 Anion: .......................................................................................................................... 8 Cation: ......................................................................................................................... 8 Isoelectric: ................................................................................................................... 8 Diagonal Relationship: ............................................................................................... 8 Effective Nuclear Charge: ......................................................................................... 8

B. Atomic Radius ........................................................................................................ 9 C. Ionization Energy .................................................................................................... 9 D. Electron Affinity ................................................................................................... 10 E. Electronegativity ................................................................................................... 10

IV. EXPLAINING ELEMENT PROPERTIES USING ARRANGEMENT OF ELECTRONS ...................................... 11

A. Explaining the properties of elements using electron configuration .................... 11 1. Number of Main Energy Levels: .......................................................................... 11 2. Number of Main Energy Levels: .......................................................................... 11 3. Net Effective Nuclear Charge: ............................................................................. 11

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I. THE PERIODIC TABLE - ORGANIZATION A. Mendeleev By the late 1790’s, the ‘father of chemistry’ Antoine Lavoisier had compiled a list of the 23 known elements. During this time, scientists were performing rigorous investigations on electricity – e.g., Benjamin Franklin. Electricity was used to break apart chemical compounds into their component elements. Although many elements had been known since prehistoric time (e.g., gold, silver), the number of known elements expanded greatly during this period. In addition, the advent of the industrial revolution led to the development of many chemistry-based industries (e.g., dyes, soaps, petrochemicals & fertilizers) helped fuel the discovery of new elements. By the 1860’s, more than 60 elements were known. With the relatively large number of known elements, scientists began to search for a systematic understanding of the elements. In 1860 a group of chemists met at the First International Congress of Chemists to discuss the use of atomic mass and other issues. At the meeting, the Italian chemist Stanislao Cannizzaro presented an exacting way to measure atomic mass, enabling chemists to agree on standard values for the average atomic mass of elements. In 1864, the John Newlands presented the “law of octaves” – when he arranged the elements by increasing mass, the chemical and physical properties repeated at every eighth element. However, he was met with considerable criticism – this pattern did not work for all known elements and many scientists resented the analogy to music as being unscientific. In 1869, the Russian scientist Dmitri Mendeleev presented his first periodic table. Mendeleev created a table by creating a series of cards – on each card he wrote the name of

the element and a list of its chemical and physical properties, and looked for repeating patterns or trends. When arranged according to atomic mass, Mendeleev noted a repeating pattern – periodic – in the chemical and physical properties of the elements. In fact, the reason that the periodic table is called ‘periodic’ results from the organization of elements into observable periodicity, or repeating pattern of chemical and physical properties. Although Mendeleev ordered the elements by atomic mass, he reversed iodine (atomic mass 127) with tellurium (atomic mass 128) due to their pattern of chemical and physical properties. Additionally, he left several empty spaces that predicted the existence and properties of three undiscovered elements. The discovery of these elements – scandium, gallium, and germanium – and that their predicted properties matched those in reality - led to wide acceptance of Mendeleev’s periodic table and earned him the title, “Father of the Periodic Table.” It is interesting that Mendeleev’s table had only 17 columns because the noble gases had not been discovered. Their existence was not predicted, or expected, due to

their lack of reactivity. The periodic table created by Mendeleev worked well but there were two nagging questions: (1) Why could one organize most of the elements’ properties? (2) What was the basis for the periodicity in the properties of the elements?

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B. Henry Moseley In 1911, the English scientist Henry Moseley, a student of Ernst Rutherford, recognized that the elements periodic trends resulted not from increasing mass but from increasing atomic number – the number of protons.. This led to the periodic law: The physical and chemical properties of the elements are functions of their atomic numbers. C. Other Periodic Tables The periodic table with which you are most familiar is not the only way to organize the elements. Many scientists have developed alternative schemes but the most useful is still the conventional system. Notice, also, that although the two lower tables are in languages other than English (left is Chinese; right is Korean), the same element symbols are used throughout the world.

II. MODERN PERIODIC TABLE

The periodic table is the most important tool for organizing, remembering, and predicting chemical facts.

A. Groups, Families & Periods, and Group Numbers Recall that the vertical columns of the periodic table are called groups or families; the horizontal rows, periods. Elements in the same group have similar chemical and physical properties. There is always some confusion about the numbering system used to label the groups. The currently accepted system (IUPAC) labels the columns at the top of the periodic table from 1 to 18a.

B. Classification of Groups

There are many ways to classify the elements. Three common and useful ways are: (a) dividing the periodic table into metals, nonmetals, and metalloids, (b) by the orbital shape of the valence electrons, i.e., s-block, p-block, d-block, or f-block), and (c) the number of valence electrons, which in turn determines much of the chemical and physical properties – by group or family names.

1. Classification – Metals, Nonmetals, Metalloids Recall that the staircase on the right side of the periodic table separates the metals from the nonmetals, and many of the elements on the staircase, itself, are metalloids. Characteristic properties of the metals, nonmetals, and metalloids are given in Table 1. Color in the elements with the appropriate color designations. However, it is important to keep in mind that these are only general properties that may, or may not, be present in the material. For example, diamonds are made of carbon atoms. Carbon is a nonmetal yet diamonds have luster. Identify the metals, nonmetals and metalloids by coloring in the boxes on the above periodic table. Use the key at the lower left to identify the group.

Table 1. Properties of Metals, Nonmetals, and Metalloids Property Metals Metalloids Nonmetals o conduct heat & electricity: good moderate poor o luster (shiny) high moderate low o ductile (ability to be drawn out in a wire) high moderate low o malleable (ability to be pounded into shape) high moderate low o high tensile strength (ability to support mass) high moderate low

• nonmetals: § most are gases at room temperature § typically gain electron(s) to form negative ions (anions)

. ⇒ oxidation states are typically negative • metalloids: § used as semiconductors in electronics (e.g., computers, memory chips)

§ 7 elements: boron (5), silicon (14), germanium (32), arsenic (33), antimony (51), tellurium (52), polonium (84)

• metals: § all are solids at room temperature except mercury (Hg) § the more reactive metals react with acids to form hydrogen gas and a salt § typically lose electron(s) to become positive ions (cations)

⇒ oxidation states are typically positive

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2. Classification – Blocks

Many properties of elements can be explained by the electron configurations of the last electron. For example, many of the visually attractive metals – e.g., iron, copper, cobalt, are associated with the available d-block electrons of the transition metals. The s- and p-blocks are grouped together into the main group elements; the d- block elements, the transition metals; the f-block elements, the inner transition metals. s-Block elements have metallic characteristics and are softest elements (e.g., can be cut with a knife). Metals of the p-block (e.g., lead and tin) are harder than s-block elements but typically softer than d-block metals. Typically, p-block metals are too reactive to be found in free state in nature (exception: bismuth).

3. Classification – Groups or Families

This is the way that elements are frequently classified. Recall that the vertical columns of the periodic table are called groups or families; the horizontal rows, periods. Elements in the same group have similar chemical and physical properties. Elements in the same group have similar properties because they have the same number of valance electrons.

a. Alkali Metals:

§ ns1 (denoting n as the number of the highest occupied energy level); lose the one valence electron to form 1+ ions

§ silvery, lustrous appearance § soft enough to cut with a knife (even a butter knife) § most reactive of the metals; so reactive they are not found in the free (metallic) state in

nature § react vigorously with nonmetals § react vigorously with water to produce hydrogen gas and basic solutions (alkaline) § typically less dense than water b. Alkaline Earth Metals:

§ ns2 (denoting n as the number of the highest occupied energy level); lose the two valence electrons to form 2+ ions

§ properties similar to alkali metals but: § harder, denser, and stronger than alkali metals § higher melting points (alkali metals have typically low melting points)

c. Hydrogen: § Although hydrogen, like the other elements in group 1, has an ns1 electron

configuration, it is not an alkali metal – it is a unique element classified by itself.

The image

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d. Transition Elements (or transition metals) § The electron configuration within this d-block sometimes deviates from the orderly

progression of filling orbitals. However, the sum of the s- and d- electrons equals the Group number.

§ Typical properties of metals (e.g., conduct electricity and heat)

§ Some are so unreactive as to be found in the free (metallic) state in nature – e.g., the ‘coinage metals’: copper, silver, and gold. Platinum, palladium, and gold are among the least reactive of all the elements.

e. Halogens

§ Most reactive nonmetals. Reactivity results from 7 valence electrons – one short of noble-gas configuration.

§ React vigorously with most metals to form salts.

f. Noble Gases:

§ John William Strutt (Lord Rayleigh) and Sir William Ramsay discovered argon in 1894. This was remarkable because the noble gases, up to that time, had escaped detection because they are unreactive and not easily observable (e.g., colorless gases). In 1868 helium had been detected on the sun, based on its line emission spectrum, but was not demonstrated to exist on Earth until 1895 by Ramsay. Interesting note: although argon makes up ~1% of air, it was not discovered until 1894 (Rayleigh and Ramsay). The name for argon comes from the Greek work argos for ‘inactive.’

§ Noble gases are unreactive because they have a full shell of valence electrons. (However classically unreactive, some noble gas compounds have been made. The first was in 1962 Bartlett and Lohmann created O2

+[PtF6]–. )

g. Actinides & Lanthanides: § The f-block elements are wedged between Groups 3 and 4 on the periodic table. These

elements are typically grouped together because their properties are so similar. Elements above neptunium (93) are all man-made.

§ The lanthanides (in the row beginning with lanthanide (57) are shiny metals with reactivity similar to the alkaline-earth metals (Group 2).

§ The actinides (in the row beginning with actinide (89) are all radioactive. Only the first four (Th, Pa, U, and Np) are naturally occurring.

h. Classification – States and Structures of Matter The three common states of matter1 are solid, liquid and gas. All of the metals, except mercury (Hg), are solids at room temperature. Mercury is a liquid. Most of the nonmetals are gases at room temperature.

1 There are actually five known states of matter. The other two, plasma and Bose-Einstein condensate, are not

within the scope of this course.

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In addition, many elements can have more than one structural form or allotrope (Table 2). For example, pure carbon can exist as several different allotropes including amorphous, diamond (like the jewelry), graphite, or fullerene (Figure 1). The atomic arrangement accounts for the properties we associate with each – as graphite, or pencil lead, sheets of carbon atoms are easily shaved off and deposited on the paper as one writes; as diamond, each carbon atom is bonded to five other atoms in fixed and immovable placement accounting for diamonds being one of the hardest substances. Amorphous structures are without a definite or fixed shape or form’. In this form, the atoms are not arranged in a regular pattern. For example, soot, which is the amorphous form of carbon, the black residue from a wood fire) is easily spread or separated apart and is very different from the carbon allotropes of graphite, diamond or fullerene. Table 2. Some Classic Examples of Allotropes Element Forms Carbon amorphous, graphite, diamond, fullerene Oxygen diatomic oxygen (O2), ozone (O3) Phosphorus white, red, and black. (White phosphorus is poisonous and can ignite

spontaneously when it comes in contact with air. Red phosphorus, used in safety matches, is not as dangerous or poisonous as white phosphorus. Black phosphorus is the least reactive form of phosphorus and has no significant commercial uses.)

Sulfur rhombic (S8, the most stable sulfur allotrope), polymeric

Graphite

Diamond

Fullerene

Figure 1. Three Allotropes of Carbon. (source: http://www.nyu.edu/pages/mathmol/library/carbon/)

Some elements exist as diatomic molecules: nonmetallic elements that, when in their elemental state, exist in pairs of atoms (Table 3). A helpful mnemonic for remembering the diatomic molecules is HONClBrIF (pronounced hon-k’l-brif).

Table 3. Diatomic Molecules Element Diatomic Molecule Hydrogen (H) H2 Oxygen (O) O2 Nitrogen (N) N2 Chlorine (Cl) Cl2 Bromine (Br) Br2 Iodine (I) I2 Fluorine (F) F2

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III. THE PERIODIC TABLE - TRENDS A. Terms Valence Electrons: outermost electrons. As we will see, these are the electrons that are available to be

lost, gained, or shared to form chemical compounds. Valence electrons hold chemical compounds together.

Octet Rule: atoms tend to gain, lose, or share electrons in order to have a full set of eight valence

electrons. Elements on the right side of the periodic table tend to gain electrons, becoming anions, in order to acquire a noble gas electron configuration. Elements on the left side of the periodic table tend to lose electrons, becoming cations, in order to acquire a noble gas electron configuration.

Ion: a negatively or positively charged atom or group of atoms. Ionic charge results from the

difference between the number of electrons and the number of protons. For example, an oxygen atom (all atoms are neutral) has eight protons and eight electrons (8+ + 8– = 0). However, an oxygen ion has eight protons (otherwise it wouldn’t be oxygen) and ten electrons, producing a net charge of 2– (8+ + 10– = 2–).

Anion: negatively-charged ion. (Pronounced “an-eye-on”; help: the ‘t’ in cation looks like a ‘+’).

Cation: positively-charged ion. (Pronounced “cat-eye-on”). Isoelectric: Two or more atoms or groups of atoms having the same electron configuration. An important consideration for determining, understanding, and predicting the reactivity of an element is based on its electronic configuration – e.g., whether the atom readily forms a cation, an anion, or stays neutral. The arrangement and movement (i.e., loss and/or gain) of electrons results from the atom achieving a noble gas electronic configuration. For example, atomic fluorine has nine protons and nine electrons. However, it does not have a complete octet of valence electrons. To achieve this, fluorine will gain an electron, forming the ion fluoride (F–). The added electron produces a complete octet of valence electrons and the entire ion has a noble gas configuration isoelectric with neon. An ion, however, doesn’t exist by itself – positive cations and negative anions will combine with each other to obtain an overall net charge of zero. Hence, Na+ ions will combine with Cl– ions to form a NaCl structure with a zero net charge. Diagonal Relationship: The close relationship between elements in adjacent groups of the periodic table (Figure at right). Often the lightest element in a group has more in common with the element located diagonal to it – e.g., Li and Mg, Be and Al, and B and Si – than the element directly below it. Effective Nuclear Charge: The net positive charge experienced by an electron in a multi-electron atom. It is involved in determining the atomic radius, ionization energy, etc. (see below), which determine the atom’s (or ion’s) reactivity. Effective nuclear charge is essentially the charge of a nucleus seen by a given electron in the atom. The term ‘effective’ is added because other (inner shell) electrons may act to shield (= shielding effect) some of the effects of the nucleus’s positive charge on outer shell electrons. This is an important consideration in understanding atomic radius, electronegativity, etc. given below.

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B. Atomic Radius The atomic radius, like the radius of a circle, is one-half the distance across it. In formal definition, the atomic radius is the one-half the distance between the nuclei of identical atoms bounded together. The atomic radius is an important consideration for the reactivity of an atom.

The radius of each atom gets larger as one goes down a group (see left). This is expected: each row adds another shell of electrons and electron repel electrons. However, what causes the radius to become smaller as one goes from left to right across a given row (see below)?

Atomic Radius Across Period 3

Element # Electrons Radius (pm) Sodium 11 186

Magnesium 12 160 Aluminum 13 143

Silicon 14 118 Phosphorus 15 108

Sulfur 16 106 Chlorine 17 99 Argon 18 97

However, the data shows that the atomic radius decreases as one goes across a period (see above table and graph for sodium through argon). The basis for this results from the balance between (1) the attraction between electrons and the positively charged nucleus and (2) the repulsion between electrons. As one moves left to right across a period, one adds not only electrons but also protons. The increased pull by the more highly charged nucleus on the electrons is stronger than the repulsion between the electrons. Thus, the atomic radii across a period decrease. Another point is the relative size of the atom and its corresponding ion. When an atom becomes a cation, it loses electron(s), so the radius of the atom is larger than the corresponding cation (Nao > Na+). Similarly, an atom becomes an anion, it gains electron(s), so the radius of the atom is smaller than the corresponding anion (Fo < F–). C. Ionization Energy An ion is an atom or group of bonded atoms that has a positive or negative charge. The ionization energy is the energy required to remove one electron from a neutral atom of an element, or in other words, the amount of energy required to form an ion.

A0 + energy → A+

+ e–

Atomic Radius Down Group 1

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This is important because the movement of electrons, much of which results from the formation of ions, can explain a lot of chemistry. This results from the balance between having the number of electrons surrounding an atom equal to the number of protons and the formation of a noble-gas configuration. The trend in the ionization energy is opposite the trend in atomic radius – as the atom gets larger, it takes less energy to remove an outer electron. Ionization energy, unless otherwise stated, is assumed to be the first ionization energy. The second ionization energy is the amount of energy to remove a second electron; the third ionization energy, the third electron, and so on. Notice in Table 4 that the 1st ionization energy for sodium is the lowest on the table. This is the easiest electron to remove: Nao → Na+ + e–. It is also possible to remove a second electron from a positive ion. For example, Group 2 atoms typically lose both s-orbital, valence electrons. Magnesium’s first two electrons are relatively easy to remove – see, in Table 4, the jump in energy between the 2nd and 3rd electrons when the two valence electrons have been removed and the 3rd electron represents removal of an electron from a noble-gas configuration of electrons. Fluorine has a propensity to gain an electron to form a noble-gas electron configuration, so the ionization energy for any of its electrons is high.

Table 4. Ionization Energies for Selected Elements. Ionization Energy (kJ/mol) Na Mg F

1st 496 738 1,681 2nd 4,562 1,451 3,374 3rd 6,912 7,733 6,050 4th 9,544 10,540 8,408 5th 13,353 13,628 11,023

D. Electron Affinity Sometimes thought of as the opposite of ionization energy. Electron affinity is the change in energy associated with the gaining of an electron by a neutral atom.

A0 + e– → A

– + energy

By convention, the quantity of energy released is represented by a negative number; energy absorbed is represented by a positive number (q.v., thermodynamic chapter). Generally, more energy is released by the halogens (Group 17). The ease with which halogens gain electrons accounts for their high reactivity. Generally, electron affinities become more negative (easier to gain electrons; higher electron affinity) as one goes across a period. Group trends, going down a given group, are not as regular. Although the halogens never add more than one electron, it is common for some atoms to gain more than one electron. For example, oxygen gains two electrons to form the O2– ion ([He]2s22p6), which is isoelectric with neon; nitrogen gains three electrons (N3–) to become isoelectric with neon. E. Electronegativity

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Electronegativity is an especially important concept in understanding chemical bonding (formation of chemical compounds). It is the measure of the ability of an atom in a chemical compound to attract electrons. Electrons are not distributed evenly throughout most chemical compounds. This uneven distribution of charge has a significant effect on the chemical properties of compounds. The most electronegative element is fluorine and is assigned an electronegativity value of 4. Other significantly electronegative elements are nitrogen, oxygen, and the halogens. The alkali and alkaline earth elements are the least electronegative. Generally, the trend in electronegativity is to decrease, or stay the same, down a group. Across a period, electronegativity tends to increase.

IV. EXPLAINING ELEMENT PROPERTIES

USING ARRANGEMENT OF ELECTRONS A. Explaining the properties of elements using electron configuration Much of how elements behave – chemically and physically, can be explained and predicted by the patterns of how electrons are arranged in an atom. For example, the overwhelming similarities between lithium ([He]2s1) and sodium ([Ne]3s1) can be explained by electron configuration, but so can their differences. The three factors influencing element behavior are explained below.

1. Number of Main Energy Levels: As one goes down a group on the periodic table, the number of shells (main energy levels) increases. Concomitantly, so does the size of the atoms. This means the valence electrons are further from the nucleus which, in turn, reduces the attractive force holding the valence electrons to the nucleus. These inner shell (non-valence) electrons produce a ‘shielding effect’ that blocks, or interferes with the attraction between the positive nucleus and the outermost electrons. Thus, the valence electrons are easier to remove.

2. Number of Main Energy Levels: The noble gases (Group 18) have a full set of eight valence electrons2. This makes them very stable and inert3. Elements react to gain, lose or share electrons in order to acquire eight valence electrons – the Octet Rule. Atoms that have completely filled energy sublevels (e.g., d-orbitals), as well as those with half-filled p-, d-, or f-sublevels, also have been found to have more stable electron arrangements.

3. Net Effective Nuclear Charge: The Effective Nuclear Charge, or kernel charge, is a relative measure of the attraction between the protons in the nucleus and the valence electrons. To calculate the kernel charge subtract the number of lower energy level electrons from the number of protons. Going from left to right across a period on the periodic table, the kernel charge increases from +1 to +8 pulling the outermost electrons closer to the nucleus. Most transition elements have a +2 kernel charge. Although successive transition metals in a period add another proton to the nucleus (e.g., lanthanide to mercury), they also add another inner-shell electron and also increase the shielding effect experienced by the valence electrons.

2 Helium has only two valence electrons but it, too, is considered to have a full set of valence electrons because

only two electrons are allowed in the s-shell available to helium. 3 The noble gases do react but, for our purposes, they are considered inert.

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ENDNOTES: a There are three ways of numbering the groups of the periodic table, one using Arabic numerals and the other two

using Roman numerals. The Roman numeral names are the original traditional names of the groups; the Arabic numeral names are those recommended by the International Union of Pure and Applied Chemistry (IUPAC) to replace the old names in an attempt to reduce the confusion generated by the two older, but mutually confusing, schemes.

There is considerable confusion surrounding the two old systems in use (old IUPAC and CAS) that combined the use of Roman numerals with letters. In the old IUPAC system the letters A and B were designated to the left (A) and right (B) part of the table, while in the CAS system the letters A and B were designated to main group elements (A) and transition elements (B). The former system was frequently used in Europe while the latter was most common in America. The new IUPAC scheme was developed to replace both systems as they confusingly used the same names to mean different things. (http://en.wikipedia.org/wiki/Periodic_table_group; Nov. 05, 2005)

“The vertical columns of the periodic table are called groups. The way in which the groups are labeled is somewhat arbitrary, and three labeling schemes are common in use... The [bottom] set of labels, which have A and B designations, is widely used in North America. Roman numerals, rather than Arabic ones, are often employed in this scheme. Group 7A, for example, is often labeled VIIA. Europeans use a similar convention that numbers the columns from 1A through 8A and then from 1B through 8B, thereby giving the label 7B (or VIIB) instead of 7A to the group headed by fluorine (F). In an effort to eliminate this confusion, the International Union of Pure and Applied Chemistry (IUPAC) has proposed a convention that numbers the groups from 1 thorugh 18 with no A or B designations, as shown... in the table... (Brown, LeMay and Bursten. Chemistrry: The Central Science AP Edition, 10th edition. p. 50) In the European system, the A represents the main group or representative elements; B, the transition metals. (Glencoe Chemistry: Matter and Change. ©2002. p.155)