{"id":3911,"date":"2018-08-15T08:52:13","date_gmt":"2018-08-15T15:52:13","guid":{"rendered":"http:\/\/www.wou.edu\/chemistry\/?page_id=3911"},"modified":"2022-09-28T13:39:57","modified_gmt":"2022-09-28T20:39:57","slug":"ch104-atoms-and-the-periodic-table","status":"publish","type":"page","link":"https:\/\/wou.edu\/chemistry\/courses\/online-chemistry-textbooks\/3890-2\/ch104-atoms-and-the-periodic-table\/","title":{"rendered":"CH104 – Chapter 2: Atoms and The Periodic Table"},"content":{"rendered":"

Chapter 2: Atoms and the Periodic Table<\/span><\/strong><\/span><\/h2>\n

This text is published under creative commons licensing, for referencing and adaptation, please click<\/span> here. <\/em><\/strong><\/a><\/span><\/p>\n

2.1 Atomic Theory with Historical Perspectives<\/span><\/strong><\/a><\/h3>\n

2.2 Introduction to Elements and the Periodic Table<\/span><\/strong><\/a><\/h3>\n

2.3 Dmitri Mendeleev and the development of the periodic table<\/strong><\/span><\/a><\/h3>\n

2.4 Families of the Periodic Table<\/span><\/strong><\/a><\/h3>\n

2.5 Defining the Atom<\/span><\/strong><\/a><\/h3>\n

Basic Atomic Structure \u2013 electrons, neutrons, and protons<\/strong><\/span><\/a><\/h4>\n

2.6 Atomic Number – <\/strong>Protons Determine the Identity of an Element<\/strong><\/span><\/a><\/h3>\n

2.7 Atomic Mass, Isotopes, and Allotropes<\/span><\/strong><\/a><\/h3>\n

2.8 Electronic Structure of Atoms<\/strong><\/span><\/a><\/h3>\n

The Four Electronic Quantum Numbers<\/span><\/strong><\/a><\/h4>\n

Electron Orbital Filling Rules<\/strong><\/span><\/a><\/h4>\n

Electron Configurations and the Periodic Table<\/span><\/strong><\/a><\/h4>\n

Electron Configuration Solitaire<\/span><\/strong><\/a><\/h4>\n

Electron Configuration Shorthand<\/span><\/strong><\/a><\/h4>\n

Electron-Dot Symbols<\/strong><\/span><\/a><\/h4>\n

2.9 Periodic Table Trends<\/strong><\/span><\/a><\/h3>\n

Atomic Size<\/strong><\/span><\/a><\/h4>\n

Electronegativity<\/strong><\/span><\/a><\/h4>\n

Ionization Energy<\/strong><\/span><\/a><\/h4>\n

Metallic and Nonmetallic Character<\/strong><\/span><\/a><\/h4>\n

2.10 Focus on the Environment – Nuclear Energy<\/strong><\/span><\/a><\/h3>\n

Introduction to Radioactivity<\/strong><\/a><\/span><\/h4>\n

Nuclear Fission Reactions and Energy<\/span><\/strong><\/span><\/a><\/h4>\n

2.11 Chapter Summary and Homework<\/strong><\/span><\/a><\/h3>\n

2.12 References<\/span><\/strong><\/a><\/h3>\n
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2.1 Atomic Theory with Historical Perspectives<\/span><\/strong><\/h3>\n

What are the smallest building blocks of everyday objects? This is a question that has interested man since the age of the Greek philosophers. Like the ancient Greeks we can perform a simple thought experiment that raises a very important question for modern chemistry: suppose you were given a piece of aluminum foil and asked to cut the foil in half over and over. How long could you continue cutting, assuming that you had no limitations based on your own abilities? Is there a limit on how small matter can be broken up into, or could you infinitely divide matter into smaller and smaller pieces? This argument dates as far back as the Greek philosophers. Most, like Aristotle, argued that matter could be divided infinitely. However, one brilliant philosopher, Democritus, argued that there is a limit. He proposed that the smallest piece that any element (like aluminum) can be divided into and still be recognized as that element is an Atom<\/strong><\/em>, a word derived from the Greek word atomos,<\/em> meaning \u201cindivisible\u201d.<\/span><\/p>\n

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Figure 2.1: Democritus<\/strong><\/span><\/p>\n

Photo taken from: <\/span>Public Domain<\/a><\/p>\n

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Philosophers, like Democritus, based most of their ideas off of thought experiments like the one above instead of actual observations and experimentation. It is for this reason that Democritus\u2019 ideas on atoms were dismissed until 1808, when John Dalton, an English scientist, proposed four fundamental assumptions based upon observations that we call Dalton\u2019s Atomic Theory.<\/strong><\/em><\/span><\/p>\n

Dalton proposed that:<\/span><\/strong><\/h4>\n
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    Matter is made up of tiny particles called atoms<\/span><\/strong><\/h4>\n<\/li>\n
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    Atoms cannot be broken into smaller pieces. During a chemical reaction, atoms are rearranged, but they do not break apart, nor are they created or destroyed<\/span><\/strong><\/h4>\n<\/li>\n
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    All atoms of the same element are identical in mass and other properties<\/span><\/strong><\/h4>\n<\/li>\n
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    Atoms of different elements differ in mass and other properties<\/span><\/strong><\/h4>\n<\/li>\n<\/ol>\n

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    2.2 Introduction to Elements and the Periodic Table<\/span><\/strong><\/h3>\n

    <\/span><\/em><\/strong>An element<\/strong> <\/em>is a substance that cannot be broken down into simpler chemical substances. There are about 90 naturally occurring elements known on Earth. Using technology, scientists have been able to create nearly 30 additional elements that are not readily found in nature. Today, chemistry recognizes a total of 118 elements which are all represented on a standard chart of the elements, called the Periodic Table of Elements. Each element is represented by a one or two letter code, where the first letter is always capitalized and, if a second letter is present, it is written in lowercase. For example, the symbol for Hydrogen is H, and the symbol for carbon is C.\u00a0 Some of the elements have seemingly strange letter codes, such as sodium which is Na. These letter codes are derived from latin terminology. For example, the symbol for sodium (Na) is derived from the latin word, natrium, which means sodium carbonate. Elements in the periodic table can be broken up into different general classes based upon similarities in their properties. Going from left to right across the periodic table, the elements can be broken up into metals, metalloids, and nonmetals.<\/span><\/p>\n

    Metals<\/strong><\/em> are typically shiny, very dense, and have high melting points. Most metals are ductile (can be drawn out into thin wires), malleable (can be hammered into thin sheets), and good conductors of both heat as well as electricity. All metals are solids at room temperature except for mercury. In chemical reactions, metals easily lose electrons to form positive ions. Examples of metals are silver, gold, and zinc.<\/span><\/p>\n

    Nonmetals<\/strong><\/em> are generally brittle, dull, have low melting points, and they are generally poor conductors of heat as well as electricity. In chemical reactions, they tend to gain electrons to form negative ions. Examples of nonmetals are hydrogen, carbon, and nitrogen.<\/span><\/p>\n

    Metalloids<\/strong><\/em> have properties of both metals and nonmetals. Metalloids can be shiny or dull. Electricity and heat can travel through metalloids, although not as easily as they can through metals. They are also called semimetals<\/em><\/strong>. They are typically semi-conductors, which means that they are elements that conduct electricity better than insulators, but not as well as conductors. They are valuable in the computer chip industry. Examples of metalloids are silicon and boron.<\/span><\/p>\n

    A.<\/span> <\/strong><\/h4>\n

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    Periodic Table Downloadable PDF Version<\/a><\/p>\n

    Figure 2.2: Periodic Table of the Elements<\/strong>.\u00a0 All of the known chemical elements are arranged in the format of a table. The table has been set up in such a way that the characteristics of each different element can be predicted by their position on the table. (A) On this rendition of the periodic table, you can see that the pink elements on the lefthand side of the table are the metals, while the blue elements on the right are the non-metals (Hydrogen is the only exception to this rule and will be explained in the subsequent sections).\u00a0 The metalloids (also termed semi-metals) occur in a stairstep pattern between the metals and nonmetals and are represented in this diagram by the green elements. (B) Shows the positions of the metals, nonmetals and metalloids on the periodic table. During this chapter, you will learn more about these unique characteristics, called periodic trends<\/em><\/strong>.<\/span><\/p>\n

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    The elements vary widely in abundance. In the universe as a whole, the most common element is hydrogen (about 90%), followed by helium (most of the remaining 10%). All other elements are present in relatively minuscule amounts, as far as we can detect. On the planet Earth, however, the situation is rather different. Oxygen makes up 46.1% of the mass of Earth\u2019s crust (the relatively thin layer of rock forming Earth\u2019s surface), mostly in combination with other elements, while silicon makes up 28.5%. Hydrogen, the most abundant element in the universe, makes up only 0.14% of Earth\u2019s crust. Table 2.1 lists the relative abundances of elements on Earth as a whole and in Earth\u2019s crust. Table 2.2 lists the relative abundances of elements in the human body. If you compare Table 2.1 and 2.2, you will find disparities between the percentage of each element in the human body and on Earth. Oxygen has the highest percentage in both cases, but carbon, the element with the second highest percentage in the body, is relatively rare on Earth and does not even appear as a separate entry in Table 2.1; carbon is part of the 0.174% representing \u201cother\u201d elements. How does the human body concentrate so many apparently rare elements?<\/span><\/p>\n

    The relative amounts of elements in the body have less to do with their abundances on Earth than with their availability in a form we can assimilate. We obtain oxygen from the air we breathe and the water we drink. We also obtain hydrogen from water. On the other hand, although carbon is present in the atmosphere as carbon dioxide, and about 80% of the atmosphere is nitrogen, we obtain those two elements from the food we eat, not the air we breathe.<\/span><\/p>\n

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    2.3 Dmitri Mendeleev and the development of the periodic table<\/strong><\/span><\/h3>\n

    In the 19th century, many previously unknown elements were discovered, and scientists noted that certain sets of elements had similar chemical properties. For example, chlorine, bromine, and iodine react with other elements (such as sodium) to make similar compounds. Likewise, lithium, sodium, and potassium react with other elements (such as oxygen) to make similar compounds. Why is this so?<\/span><\/p>\n

    In 1864, Julius Lothar Meyer, a German chemist, organized the elements by atomic mass and grouped them according to their chemical properties. Later that decade, Dmitri Mendeleev, a Russian chemist, organized all the known elements according to similar properties. He left gaps in his table for what he thought were undiscovered elements, and he made some bold predictions regarding the properties of those undiscovered elements. Later, when elements were discovered whose properties closely matched Mendeleev\u2019s predictions, his version of the table gained favor in the scientific community. Because certain properties of the elements repeat on a regular basis throughout the table (that is, they are periodic), it became known as the periodic table<\/strong><\/em>.<\/span><\/p>\n

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    Figure 2.3: Dmitri Medeleev’s 1871 Early Version of the Periodic Table<\/strong><\/span><\/p>\n

    Photo of Dmitri Medeleev provided by:<\/span> \u043a\u0430\u0431\u0438\u043d\u0435\u0442 \u0430\u043a\u0430\u0434\u0435\u043c\u0438\u043a\u0430 \u041c\u0438\u0445\u0430\u0438\u043b\u0430 \u041c\u0438\u0445\u0430\u0439\u043b\u043e\u0432\u0438\u0447\u0430 \u0428\u0443\u043b\u044c\u0446\u0430 – \u0444\u043e\u0442\u043e \u043b\u044e\u0431\u0435\u0437\u043d\u043e \u043f\u0435\u0440\u0435\u0434\u0430\u043d\u043e \u043c\u043d\u0435 \u0432 \u0441\u043e\u0431\u0441\u0442\u0432\u0435\u043d\u043d\u043e\u0441\u0442\u044c \u0432\u0434\u043e\u0432\u043e\u0439 \u041c.\u041c.\u0428\u0443\u043b\u044c\u0446\u0430 \u041d\u0438\u043d\u043e\u0439 \u0414\u043c\u0438\u0442\u0440\u0438\u0435\u0432\u043d\u043e\u0439 \u0428\u0443\u043b\u044c\u0446<\/a><\/p>\n

    Periodic Table provided by<\/span>:<\/span> Den fj\u00e4ttrade ankan<\/a><\/p>\n

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    \u00a02.4 Families of the Periodic Table<\/span><\/strong><\/h3>\n

    Remember that Mendeleev arranged the periodic table so that elements with the most similar properties were placed in the same group. A group, or family of elements<\/strong><\/em>, is a vertical column of the periodic table. Elements are placed into families due to their similar properties, characteristics, and reactivities.\u00a0\u00a0 For example, all of the elements in group 1 (except for hydrogen, which has unique properties) are very reactive and form compounds in the same ratios and with similar properties as other 1 elements. Due to the similarities in their chemical properties, Mendeleev put these elements into the same group and they came to be known as the alkali metals<\/strong><\/em>. The alkali metals<\/em><\/strong> include: lithium, sodium, potassium, rubidium, cesium, and francium. Alkali metals are among the most reactive metals. This is due in part to their larger atomic radii and low ionization energies, that will be discussed in more details in section 2.8 below. They get their name from ancient Arabic (al qali) because “scientists” of the time found that the ashes of the vegetation they were burning contained a large amount of sodium and potassium. In Arabic, al qali<\/i> means ashes<\/i>. Although most metals tend to be very hard, alkali metals have a soft texture, are silvery in color and can be easily cut. They also have low boiling and melting points and are less dense than most elements. Figure 2.4 shows some of the most common families on the periodic table.<\/span>
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    Figure 2.4 Common Families and Groups of the Periodic Table.<\/strong><\/span><\/p>\n

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    The same pattern is true of other vertical groups on the periodic table. Group 2 is called the alkaline earth metals<\/strong><\/em>. Once again these elements have similar properties to each other. alkaline earth metals include Beryllium, Magnesium, Calcium, Barium, Strontium and Radium and are soft, silver metals that are less metallic in character than the Group 1 alkali metals. Although many characteristics are common throughout the group, the heavier metals such as Ca, Sr, Ba, and Ra are almost as reactive as the Group 1 alkali metals. They get their name because early “scientists” found that all of the alkaline earth metals were found in the earth’s crust.
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    The transition metals<\/em><\/strong> are the larger block of elements shown in purple on Figure 2.4 extending from Groups 3-12 (also known as the group B elements). Transition elements differ from the main group elements (group A elements) in that they tend to be hard and have high densities.\u00a0 They have high melting points and boiling points and can show various oxidation states when forming chemical bonds (this will be discussed further in chapter 3). They often form colored compounds that are highly stable and they can serve as good catalysts<\/em><\/strong>.\u00a0 A catalyst<\/em><\/strong> is an agent that helps to speed up a chemical reaction without itself being changed in the process.\u00a0<\/span><\/p>\n

    Group 17 elements are also called halogens<\/strong><\/em>. This group contains very reactive nonmetals. The halogens<\/b><\/em> are an interesting group. Halogens are members of Group 17, which is also referred to as 7A. It is the only group in the Periodic Table that contains all of the states of matter at room temperature. Fluorine, F2 <\/sub>and <\/sub>chlorine, Cl2<\/sub> are gases, while Bromine, Br2<\/sub>, is a liquid and iodine, I2<\/sub>, and astatine, At, are both solids. Another interesting feature about Group 17 is that it houses four (4) of the seven (7) diatomic elements<\/em><\/strong>. Diatomic elements<\/em><\/strong> only exist in nature as a pair of atoms of the same element that are bonded together. The seven diatomic elements are H2<\/sub>, N2<\/sub>, O2<\/sub>, F2<\/sub>, Cl2<\/sub>, Br2<\/sub>, and I2<\/sub>. Notice that the latter four are Group 17 elements. The word halogen comes from the Greek meaning salt forming<\/i>. French chemists discovered that the majority of halogen ions will form salts when combined with metals.
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    The noble gases<\/strong> are in group 18. The two most significant properties of noble gases is that they are extremely unreactive, rarely forming compounds, and that they all exist as gases at room temperature. We will learn the reason for their unreactivity when we discuss how compounds form in chapters 3 and 4. The first person to isolate a noble gas was Henry Cavendish, who isolated argon in the late 1700s. The noble gases were actually considered inert gases until the 1960s when a compound was formed between xenon and fluorine which changed the way chemists viewed the “inert” gases. In the English language, inert means to be lifeless or motionless; in the chemical world, inert means does not react<\/i>. Later, the name “noble gas<\/b>” replaced “inert gas” for the name of Group 18. The elements in this group are also gases at room temperature.<\/span><\/p>\n

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    2.5 Defining the Atom<\/span><\/strong><\/h3>\n

    Basic Atomic Structure \u2013 electrons, neutrons, and protons<\/strong><\/span><\/h4>\n

    The modern atomic theory, proposed about 1803 by the English chemist John Dalton, is a fundamental concept that states that all elements are composed of atoms. An atom is the smallest part of an element that maintains the identity of that element. Individual atoms are extremely small; even the largest atom has an approximate diameter of only 5.4 \u00d7 10\u221210<\/sup> m. With that size, it takes over 18 million of these atoms, lined up side by side, to equal the width of your little finger (about 1 cm).<\/span><\/p>\n

    Most elements in their pure form exist as individual atoms. For example, a macroscopic chunk of iron metal is composed, microscopically, of individual iron atoms. Some elements, however, exist as groups of atoms called molecules. Several important elements exist as two-atom combinations and are called diatomic molecules. In representing a diatomic molecule, we use the symbol of the element and include the subscript 2 to indicate that two atoms of that element are joined together. The elements that exist as diatomic molecules are hydrogen (H2<\/sub>), oxygen (O2<\/sub>), nitrogen (N2<\/sub>), fluorine (F2<\/sub>), chlorine (Cl2<\/sub>), bromine (Br2<\/sub>), and iodine (I2<\/sub>).<\/span><\/p>\n

    Atoms are made up of extremely small subatomic particles called protons, neutrons, and electrons. Protons<\/strong><\/em> are positively charged particles with a relative mass of 1.672622×10-24<\/sup>g, which form part of the core nucleus<\/strong> <\/em>of an atom. The other part of the atomic nucleus is made up of neutrons<\/em>, <\/strong>electrically neutral particles with a relative mass almost identical to a proton (1.674927×10-24<\/sup>g). Electrons<\/strong><\/em> are extremely small (9.109328×10-28<\/sup>g) negatively charged particles that form an electron cloud, which orbits the nucleus.Table 2.3 summarizes some of the general properties of subatomic particles.<\/span>
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    Experiment have shown that protons and neutrons are concentrated in a central region of each atom called the nucleus (plural, nuclei). Electrons are outside the nucleus and orbit about it because they are attracted to the positive charge in the nucleus. Figures 2.5 and 2.6 depict the structure of an atom.<\/span>
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    Figure 2.5 The Anatomy of an Atom.<\/strong> The protons and neutrons of an atom are found clustered at the center of the atom in a structure called the nucleus.\u00a0 The electrons orbit the nucleus of the atom within an electron cloud, or the empty space that surrounds the atom\u2019s nucleus.\u00a0 Note that most of the area of an atom is taken up by the empty space of the electron cloud.<\/span><\/p>\n

    Diagram provided by: <\/span>CNX OpenStax<\/a><\/p>\n

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    Fig 2.6 The path of the electron in a hydrogen atom.\u00a0<\/strong> Electrons are not in discrete orbits like planets around the sun. Instead there is a probability that an electron may occupy a certain space within the electron cloud (a) The darker the color, the higher the probability that the hydrogen\u2019s one electron will be at that point at any given time. (b) Similarly, the more crowded the dots, the higher the probability that hydrogen\u2019s one electron will be at that point.\u00a0 In both diagrams, the nucleus is in the center of the diagram.<\/span><\/p>\n

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    Most of the mass of an atom is in the nucleus, while the orbiting electrons account for an atom\u2019s size. As a result, an atom consists largely of empty space. Atomic particles are so small that it is impractical to measure them in grams, instead we use a relative mass scale which makes the numbers much more manageable. We use Atomic Mass Units (AMU or u)<\/strong><\/em> to measure the mass of atomic particles, one AMU is equal to 1\/12th the mass of an atom of carbon-12. You may also see Atomic Mass Units referred to as Daltons (Da)<\/strong><\/em> after John Dalton, the English Chemist that first proposed the atomic theory.\u00a0 Carbon-12 has 6 protons and 6 neutrons in its nucleus, meaning that one amu is equal to the average of the masses of a proton and a neutron. The mass of an atom in AMUs is equal to the number of protons and neutrons making up the atom. For example the atomic mass of bromine is roughly 80 amu and its proton number is 35, meaning that bromine has 35 protons and 45 neutrons in its nucleus.<\/span><\/p>\n

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    2.6 Atomic Number – <\/strong>Protons Determine the Identity of an Element<\/strong><\/span><\/h3>\n

    When looking at the periodic table you might notice that for each element there are two sets of numbers around the symbol. These symbols correspond to important values that give you important information about each element (Figure 2.7). The most important value corresponding to characteristics of an element is the proton number, which is also called a<\/strong>tomic number<\/strong><\/em> (represented by the mathematical term, Z<\/em>). As it turns out, the number of protons that an atom holds in its nucleus is the key determining feature for its chemical properties.\u00a0 In short, an element is defined by the number of protons found in its nucleus. If you refer back to the Periodic Table of Elements shown in figure 2.2, you will see that the periodic table is organized by the number of protons that an element contains. Thus, as you read across each row of the Periodic Table (left to right), each element increases by one proton (or one Atomic Number, Z<\/em>). When atoms are in their elemental states, their overall charge is zero and the atoms are neutral.\u00a0 Since protons are positively charged and electrons are negatively charged, this means that when atoms are in their elemental form, the number of protons equals the number of electrons. Therefore, if you know the atomic number of an atom, you also know how many electrons are present in that atom when it is in its elemental form.\u00a0 When atoms combine with one another to form compounds, like water (H2<\/sub>O), they will either share or donate\/accept electrons from their bonding partners.\u00a0 It is this movement of electrons that facilitates chemical bond formation. Thus, during bond formation the number of electrons around an atom may change, but the atomic number (or number of protons) remains constant and does not change.<\/span>
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    Fig 2.7 Structure of the Periodic Table.<\/strong> Each element on the periodic table is represented by the atomic symbol (Cu for Copper, and Te for Tellurium).\u00a0 Sometimes the Atomic Number is written in the upper lefthand corner, and the Atomic Mass in the righthand corner, as shown in this figure.\u00a0 Sometimes, periodic tables will show the atomic number above the element symbol and the atomic mass below the element symbol, as shown in the periodic table in Figure 2.2.<\/span>
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    2.7 Atomic Mass, Isotopes, and Allotropes<\/span><\/strong><\/h3>\n

    Atomic mass (A)<\/em><\/strong> is the total mass of an atom of a specific element and can be calculated by adding up the number of protons and neutrons present within an atom. The electrons are ignored in the mass calculation because they are so small that they barely add any mass to the atom.\u00a0<\/span> <\/span><\/p>\n

    # Protons + # Neutrons = Atomic Mass<\/span><\/strong><\/h4>\n

    Thus, if you know any two of the the three values (atomic mass, atomic number, or number of neutrons), you can calculate the third value.\u00a0 For example, nitrogen has an atomic mass of 14.007 and an atomic number of 7.\u00a0 Thus, it contains 7 protons, and 7 neutrons (14.007 – 7 = 7.007, which is then rounded to 7).\u00a0 Note that the number of neutrons in an atom does not have to equal the number of protons in the atom. For example, lead (Pb), contains 82 protons and has an atomic mass of 207.2.\u00a0 Thus it contains 125 neutrons (207.2 – 82 = 125).\u00a0 Thus, if you know the atomic mass and the atomic number of an element, you can calculate the number of neutrons present, or if you know the atomic mass and the number of neutrons, you can calculate the atomic number.
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    So…how many neutrons are in atoms of a particular element? At first it was thought that the number of neutrons in a nucleus was also defining characteristic of an element. However, it was found that atoms of the same element can have different numbers of neutrons. Atoms of the same element that have different numbers of neutrons are called isotopes. <\/strong><\/em><\/strong>An example of the three common isotopes of hydrogen are shown in Figure 2.8. Note that each of the hydrogen isotopes is known by a unique name, hydrogen, deuterium, and tritium.\u00a0 Not all elemental isotopes have such unique names.\u00a0 Many of the isotopes are distinguished from one another by including the atomic mass in the definition. For example, 99% of the carbon atoms on Earth have 6 neutrons and 6 protons in their nuclei, this is known as carbon-12; just under 1% of the carbon atoms have 7 neutrons and 6 protons in their nuclei, which is known as carbon-13, and an even smaller percent is carbon with 8 neutrons and 6 protons, or carbon-14. Carbon-14 is unstable and will decay over time making it a radioactive form of carbon.\u00a0 The half-life of carbon, or the time it takes for half of the isotope to breakdown is 5,700 years.\u00a0 Overall, there are 15 known isotopes of carbon!\u00a0 Thus, naturally occurring carbon on Earth, therefore, is actually a mixture of isotopes, albeit a mixture that is 99% carbon-12. Isotope composition has proven to be a useful method for dating many rock layers and fossils.<\/span><\/p>\n

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    Fig 2.8 Isotopes of Hydrogen.<\/strong> All hydrogen atoms have one proton and one electron. However, they can differ in the number of neutrons. (a) Most hydrogen atoms only contain one proton and one electron and no neutrons (b) A small amount of hydrogen exists as the isotope deuterium, which has one proton and one neutron in its nucleus, and (c) an even smaller amount contains one proton and two neutrons in its nucleus and is termed Tritium. Note that Tritium is unstable isotope and will breakdown over time.\u00a0Thus, Tritium is a radioactive element.<\/span><\/p>\n

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    Most elements exist as mixtures of isotopes. In fact, there are currently over 3,500 isotopes known for all the elements. When scientists discuss individual isotopes, they need an efficient way to specify the number of neutrons in any particular nucleus. A simple way of indicating the mass number of a particular isotope is to list it as a superscript on the left side of an element\u2019s symbol. Atomic numbers are often listed as a subscript on the left side of an element\u2019s symbol. Thus, we might see<\/span><\/p>\n

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    which indicates a particular isotope of copper. The 29 is the atomic number, Z<\/em>, (which is the same for all copper atoms), while the 63 is the atomic mass (A) of the isotope. To determine the number of neutrons in this isotope, we subtract 29 from 63: 63 \u2212 29 = 34, so there are 34 neutrons in this atom.<\/span><\/p>\n

    The atomic masses indicated on the periodic table represents an average mass for each element based on the proportion of each isotope present on the Earth. This is why most of the atomic masses on the periodic table are not exact numbers.\u00a0 For example, the atomic mass of copper is 63.546 amu. This mass is an average of an element\u2019s atomic masses, weighted by the natural abundance of each isotope. So how could we calculate atomic mass based on the natural abundance of different isotopes of an element?<\/span>
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    Example: Calculating Atomic Mass Using Isotope Abundance<\/span><\/strong><\/em><\/h3>\n

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    Extra Practice:<\/strong><\/em><\/span><\/h3>\n

    Try to work out the atomic mass for boron.\u00a0 Boron exists as a mixture that is 19.9% 10<\/sup>B and 80.1% 11<\/sup>B. Calculate the atomic mass.\u00a0 Check the periodic table for the correct answer!<\/strong><\/span><\/h3>\n
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    Allotropes<\/strong><\/em> of an element are different and separate from the term isotope and should not be confused. Some chemical elements can form more than one type of structural lattice, these different structural lattices are known as allotropes<\/strong><\/em>.\u00a0 This is the case for phosphorus as shown in Figure 2.9.\u00a0White or yellow phosphorus forms when four phosphorus atoms align in a tetrahedral conformation (Fig 2.9).\u00a0The other crystal lattices of phosphorus are more complex and can be formed by exposing phosphorus to different temperatures and pressures. \u00a0For example, the cage-like lattice of red phosphorus can be formed by heating white phosphorus over 280o<\/sup>C (Fig 2.9).\u00a0Note that allotropic changes\u00a0affect how the atoms of the element interact with one another to form a 3-dimensional structure.\u00a0\u00a0They do not alter the\u00a0sample with regard to the atomic\u00a0isotope forms that are present, and DO NOT<\/em> alter or affect\u00a0the atomic mass (A<\/em>) of the element.<\/span><\/p>\n

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    Figure 2.9 Allotropes of Phosphorus.<\/strong>\u00a0 (A) White phosphorus exists as a (B) tetrahedral form of phosphorus, whereas (C) red phosphorus has a more (D) cage-like crystal lattice.\u00a0 (E) The different elemental forms of phosphorus can be created by treating samples of white phosphorus with increasing temperature and pressure.<\/span><\/p>\n

    Source: <\/span>https:\/\/en.wikipedia.org\/wiki\/Allotropes_of_phosphorus<\/a><\/span><\/p>\n

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    Different allotropes of different elements can have different physical and chemical\u00a0properties and are thus, still important to consider. For example, oxygen has two different allotropes with the dominant allotrope\u00a0being the diatomic form of oxygen, O2<\/sub>.\u00a0 However, oxygen can also exist as O3<\/sub>, ozone. In the lower atmosphere, ozone is produced as a by-product in automobile exhaust, and other industrial processes\u00a0where it contributes to pollution. It has a very pungent smell and is a very powerful oxidant.\u00a0 It can cause damage to mucous membranes and respiratory tissues in animals.\u00a0 Exposure to ozone has been linked to premature death, asthma, bronchitis, heart attacks and other cardiopulmonary diseases.\u00a0 In the upper atmosphere, it is created\u00a0by natural electrical discharges\u00a0and exists at very low concentrations. The presence of ozone in the upper atmosphere\u00a0is critically\u00a0important as it intercepts very damaging ultraviolet radiation from the sun, preventing it from reaching the Earth\u2019s surface. Thus, ozone can be a health hazard or a health protector, depending upon where it is found!<\/span>
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    (Back to the Top)<\/em><\/strong><\/span><\/a><\/p>\n

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    2.8 Electronic Structure of Atoms<\/strong><\/span><\/h3>\n

    Although we have discussed the general arrangement of subatomic particles in atoms, we have said little about how electrons occupy the space about the nucleus. Do they move around the nucleus at random, or do they exist in some ordered arrangement?<\/span><\/p>\n

    The modern theory of electron behavior is predicted by the field of quantum mechanics<\/strong><\/em>. The term quantum<\/em><\/strong> is defined as a discrete quantity or ‘packet’ of energy. For example, light has qualities of both a wave and a particle.\u00a0 Since it has properties of a particle, the energy of light is said to be quantized into discrete ‘packets’ of light particles called photons<\/strong><\/em>. Each photon of light, then carries a quantized amount of energy that is dependent on its frequency.
    \n<\/span><\/p>\n

    Electrons in atoms are also quantized and can exist at different energy levels depending on how far away from the nucleus of that atom that they are positioned. This is of interest, because the electrons are the mobile part of the atom and they are involved in forming chemical bonds. Understanding their positioning around the nucleus of the atom helps to predict how they will combine with other atoms to form chemical compounds. Electrons typically have higher energy, the farther away they are (on average) from the nucleus. However, the increasing energy of an electron does not<\/em><\/strong> proceed continuously like a ramp the farther away from the nucleus that it travels. Instead, the energy of an electron is quantized, meaning that the energy of an electron does not increase continuously, but instead increases by levels adding discrete packets of energy. This would be similar to walking up a staircase rather than a ramp.\u00a0 On a staircase, you can step up one step at a time, or maybe two steps at a time, but you can never step up a half of a step.\u00a0 Similarly, electrons travel between energy levels instantly, never having an energy amount in between two levels.<\/span><\/p>\n

    A total of four quantum numbers are used to describe completely the movement and trajectories of each electron\u00a0within an atom. According to the Pauli Exclusion Principle<\/em><\/strong>, no two electrons can share the same combination of four quantum numbers. Thus, each electron in an atom has a unique set of quantum numbers that helps define the probable location of an electron within an atom. The quantum numbers that help to define electron location exist in a hierarchical order. This would be very similar to finding an address of a friend. If you don’t know your friend very well, you may only know the city where your friend lives. This doesn’t provide very much information. As you get to know them better, you may find out what neighborhood in the city that they live in. At a deeper level, you may find our what street they live on, and then at the deepest level, what their house number is on that street.\u00a0 However, knowing which house your friend lives in does not tell you where in the house you friend is at, maybe they are in the kitchen or the backyard.\u00a0 Determining the exact location of an electron is not possible.\u00a0 We can only find the approximate\u00a0 location or “address” of the electron.<\/span> <\/span><\/p>\n

    The Four Electronic Quantum Numbers<\/span><\/strong><\/h4>\n

    There are a total of four quantum numbers: the principal quantum number (n<\/em>), the orbital angular momentum quantum number (l<\/em>), the magnetic quantum number (ml<\/sub><\/em>), and the electron spin quantum number (ms<\/sub><\/em>) that are used to describe the major characteristics and spatial distribution of electrons within an atom. The quantum numbers correspond to the following hierarchical layers: the principle quantum number (n<\/em>) is the broadest classification and corresponds to the energy shell (this would be equivalent to the city in our address analogy), the orbital angular momentum quantum number (l<\/em>) is the next layer and corresponds to the subshells (this would be the neighborhood in our address analogy), the magnetic quantum number (ml<\/sub><\/em>) corresponds to the electron orbitals (or the street where your friend lives), and electron spin quantum number (ms<\/sub><\/em>) defines the spins of electrons (or the house number in our address analogy).<\/span> <\/span><\/span><\/span><\/span><\/span><\/span><\/p>\n

    n<\/em> = shell<\/span><\/strong><\/h4>\n

    l<\/em> = subshell<\/span><\/strong><\/h4>\n

    ml<\/sub><\/em> = orbital<\/span><\/strong><\/h4>\n

    ms<\/sub><\/em> = electron spin<\/span><\/strong><\/h4>\n
    <\/div>\n

    The principal quantum number, n<\/span><\/span><\/em><\/span><\/span>,\u00a0designates the electron shell. Because\u00a0n<\/em>\u00a0describes the most probable distance of the electrons from the nucleus, the larger the number\u00a0n<\/em>\u00a0is, the farther the electron is from the nucleus and the larger the atom. The principal quantum number (n)<\/em> can be any positive integer between 1 and 7. An n<\/span>=<\/span>1<\/span><\/span><\/span><\/span> designates the first principal shell (the innermost shell). The first principal shell is also called the ground state, or lowest energy state. When an electron is in an excited state or gains energy, it may jump to the second principle shell, where n<\/span>=<\/span>2<\/span><\/span><\/span><\/span>. This process is called absorption<\/em><\/strong> because the electron is “absorbing” photons, or gaining energy. Thus, as the energy of the electron increases during absorption, so does the principal quantum number, e.g., if an electron in the n<\/em> = 3\u00a0 shell absorbs energy, it will jump to the fourth principal shell, n<\/em> = 4.\u00a0The opposite process is emission<\/em><\/strong>, where electrons “emit” or release energy as they fall from higher to lower principle shells. In this case, n<\/em>\u00a0decreases by whole numbers (packets or quanta of energy).\u00a0 Within atoms there is a maximum number of seven electron shells (n<\/em> = 7), after which the atom becomes too unstable to hold electrons in place.
    \n<\/span><\/p>\n

    At a basic level, the energy shells can be thought of like the layering of an onion, with the first layer being the smallest and closest to the core of the nucleus, and the subsequent layers being larger and farther away from the core. Larger shells have more space and can hold more electrons. The general formula for determining how many electrons can be housed in each shell is 2n<\/em>2 <\/sup>and is shown in Table 2.4.\u00a0 While this table predicts that that outer shells, which are the largest shells can contain 50, 72, and 98 electrons, elements that have this many electrons in those shells have never been discovered.\u00a0 It is likely that they are too big and unstable to exist.<\/span>
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    Table 2.4 Maximum Number of Electrons Per Shell<\/span><\/strong><\/h4>\n

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    The periodic table allows us to easily determine how many protons (and therefore electrons) an element will have by showing the atomic number (Z<\/em>) in each element box.\u00a0 Equally useful, the periodic table is set up to represent the different quantum numbers to facilitate the determination of electron configuration of an element. The principle quantum number (n<\/em>) or number of shells present in an element are shown in the periods or rows of the periodic table (Figure 2.10). For example, if we look at the\u00a0 sodium atom, we will see that it is in row 3 of the periodic table.\u00a0 This means that it has three electron shells that can house electrons. As electrons fill their available orbital spaces, they always fill the shells starting at the lowest energy levels and going up to higher levels as needed. This rule can be formally noted as the Aufbau principle<\/strong> which states that electrons orbiting one or more atoms will fill the lowest available energy levels before filling higher energy levels.<\/span><\/p>\n

    (Back to the Top)<\/em><\/strong><\/span><\/a><\/p>\n

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    Figure 2.10 The Periods of the Periodic Table Represent Electron Shells<\/strong> (A) Each electron shell is represented by a row or period on the the periodic table.\u00a0 (B) A representation of the electron shells of the sodium atom.\u00a0 Sodium has an atomic number of 11, thus it has 11 protons and 11 electrons in the elemental form. The first and innermost electron shell can hold a total of 2 electrons, as indicated by the gray circles in the diagram.\u00a0 The second electron shell can hold 8 electrons and is completely full. The third electron shell has the capacity to hold 18 electrons, however, in sodium, there is only one more electron that needs to be placed after filling the first and second electron shells.\u00a0 Thus, the third electron shell of sodium only has 1 electron.\u00a0<\/span><\/p>\n

    Figure 2.10B adapted from <\/span>Pumbaa<\/a><\/p>\n

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    The electron shells are further subdivided into subshells that are represented by the second quantum number. There are four major types of subshells that are used to house electrons in known elements.\u00a0 These are termed the s, p, d<\/em>, and f<\/em> subshells. The first electron shell (n<\/em> = 1) is so small that it only contains an s<\/em>-subshell. The second electron shell\u00a0(n<\/em> = 2) is large enough to house an s-<\/em> and a p-<\/em>subshell. The third electron shell (n<\/em> = 3) contains s-, p-<\/em>, and d-<\/em>subshells. The fourth shell and above (n<\/em> = 4, 5, 6, and 7) contain all four electron subshells (s, p, d, and f<\/em>). Shells 5-7 are predicted to have additional subshells, but elements using these subshells have never been discovered, so we will limit our discussion with the s, p, d,<\/em> and f<\/em> subshells.<\/span>
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    Each subshell, in turn,\u00a0 contains a specified number of electron orbitals arranged in unique shapes (Figure 2.11). Each orbital can then contain a total of 2 electrons rotating in opposite spin states. The s<\/em>-subshell only contains 1 circular orbital that can house a total of two electrons. The p<\/em>-subshell contains a total of 3 dumbbell-shaped orbitals that can house a total of 6 electrons. The d<\/em>-subshell contains 5 uniquely shaped orbitals that can house a total of 10 electrons and the f<\/em>-subshell contains 7 orbitals with a potential of housing 14 electrons.\u00a0 While the 5th – 7th shells are predicted to be able to house more electrons in even stranger g, h,<\/em> and i<\/em> subshells, there have never been any atoms discovered that contain enough electrons to utilize these locations.\u00a0 Atoms appear to become too large and unstable after the filling of the f<\/em>-subshell orbitals in the 4th and 5th shell.<\/span><\/p>\n

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    Figure 2.11 Orbital Shape Diagrams for the s, p, d,<\/em> and f<\/em> Subshells. <\/strong>A pair of electrons with opposite spin can be housed within each of the electron orbitals shown above. <\/span><\/p>\n

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    We can summarize the availability of electron orbitals below:<\/span><\/p>\n

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    (Back to the Top)<\/em><\/strong><\/span><\/a><\/p>\n

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    Electron Orbital Filling Rules<\/strong><\/span><\/h4>\n

    Now that we have defined the physical spaces that electrons can occupy, we need to determine the order of electron orbital filling. There are three major rules that we need to follow when filling electron orbitals.\u00a0 Recall that the Pauli Exclusion Principle<\/strong><\/em> excludes electrons from having the same quantum numbers, and thus, they cannot exist in the same place at the same time. Second according to the Aufbau principle<\/em>, <\/strong>electrons orbiting one or more atoms will fill the lowest available energy levels before filling higher energy levels. Thus, the lowest energy levels must be filled with electrons before orbitals with higher energy can house electrons. The final rule that we need to follow is Hund\u2019s rule<\/strong> which states that when electron orbitals have equal energy level, electrons must fill each of those orbitals as single electrons before they can begin to pair with electrons of opposite spin states. This is because electrons are negatively charged and naturally repel each other.\u00a0 Thus, they exist as far apart as possible. <\/strong>This is the most energetically stable way that electrons can fill orbitals.
    \n<\/span><\/p>\n

    Due to the strange shapes of the orbitals within the larger subshells, the energy levels of the orbitals don’t always correspond directly with their shell level.\u00a0 For example, the d<\/em>-subshell orbitals have a higher energy than the s<\/em>-subshell orbital from next shell level\u00a0 (i.e. the d-<\/em>subshell orbitals from the 3rd shell have a higher energy and will fill with electrons later that the s-<\/em>subshell from the 4th shell). Figure 2.12 depicts the energy levels from orbitals in the different subshells.<\/span> <\/span><\/p>\n

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    Figure 2.12 Electron Energy Filling Diagram.\u00a0 <\/strong>Orbitals with the lowest energy are filled with electrons before orbitals at higher energy levels.<\/span><\/p>\n

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    Note that in the energy diagram above that there are many orbitals that share the same energy level.\u00a0 For example the three 2p orbitals have the same energy level.\u00a0 Thus, when these orbitals are filled with electrons, they are filled according to Hund’s Rule which states that single electrons will occupy orbitals at the same energy level before they will pair up.\u00a0 (If it helps, you can think to yourself…’electrons find each other repulsive and like to live alone’) The following example shows the correct and the incorrect way to fill the 2p orbitals.<\/span> <\/span><\/p>\n

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    Image provided by<\/span> Wikimedia<\/a><\/p>\n

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    We can use the electron energy filling diagram to complete an example of how electrons in an element are arranged within the electron orbitals (Figure 2.13).\u00a0 Let’s take the example of Iron, (Fe).\u00a0 Iron has an atomic number of 26 or 26 protons.\u00a0 Thus in the elemental state, iron will also have 26 electrons. We will use arrows to represent our electrons within our diagram.\u00a0 Arrows pointing in opposite directions denote electrons that have opposite spin states. Note that the first set of electrons fills the 1s orbital.\u00a0 The next 8 electrons then fill the 2s and 2p orbitals. Note that when filling orbitals that have the same energy level, that the electrons have to follow Hund’s Rule, and fill these orbitals as single electrons first, and then pair up when there are no other options to exist in an orbital as a single electron.\u00a0 This filling pattern is noted in the lefthand diagram of the iron electron configuration.\u00a0 The lefthand diagram shows the active filling of the 2p orbitals with one electron at a time.\u00a0 After all three orbitals receive an electron, they will then pair up before filling the higher energy 3s orbital.\u00a0 Electron filling continues until all 26 electrons are placed.\u00a0 Note that in the case of iron, that the final electrons are placed in the 3d subshell orbitals following Hund’s Rule.\u00a0<\/span><\/p>\n

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    Figure 2.13\u00a0 Electron Filling Diagram for the Element Iron (Fe).\u00a0<\/strong> The lefthand diagram shows the partial filling of the electron orbitals during the assembly of an iron atom.\u00a0 Filling begins with the lowest energy orbital, which is the 1s obital.\u00a0 This is followed by the 2s orbital and the three 2p orbitals.\u00a0 Note that all of the 2p orbitals are at the same energy level. Thus, they are filled according to Hund’s Rule:\u00a0 each orbital at the same energy level will fill with one electron alone, until there are no other options.\u00a0 Then the electrons will pair to complete the filling of the 2p orbitals. The figure to the right denotes the completed electron filling diagram for the element iron.\u00a0 Note that the final electrons are placed in the 3d orbitals, and that they are placed to obey Hund’s Rule, leaving 4 of the 6 electrons at this energy level in the unpaired state.<\/span><\/p>\n

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    We can write the final electron configuration of iron (Fe) using a notation that denotes each electron shell and subshell and the superscript notation that refers the number of electrons present in the orbitals at that energy level.\u00a0 Remember that s<\/em>-subshells can house 2 electrons, p<\/em>-subshells can house 6 electrons, d-<\/em>subshells can house 10 electrons, and f-<\/em>subshells can house a maximum of 14 electrons.<\/span> <\/span><\/p>\n

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    From the diagram above, we can notice a few things.\u00a0 First, if we add up the superscript numbers, this should equal the total number of electrons in that atom.\u00a0 For example if we take the sum of the superscript numbers from the diagram above, we get:<\/span><\/p>\n

    2 + 2 + 6 +2 + 6 + 2 + 6 = 26<\/strong><\/span><\/h4>\n

    We can use this as a way to double check and make sure that we have constructed the electron configuration correctly.<\/span><\/p>\n

    The second piece of information that we can find very quickly is the number of electrons present in the outermost shell of the atom.\u00a0 The outermost shell of the atom is referred to as the valence shell<\/strong><\/em>, and it contains the most available and reactive electrons within the atom. The valence shell will always correspond to the highest shell number present in the atom that contains electrons. The other electron shells are known as inner electron shells. In the case of iron, the valence shell would be the 4th shell (or period on the periodic table). Shells 1-3 would then be the inner shells. Electrons housed in the valence shell are known as the valence electrons<\/em><\/strong>. Valence electrons are important because they are the ones that are most involved in chemical bonding, although sometimes electrons that are close to the valence electrons in energy, may also participate in bond formation. Most often these exceptions involve electrons in the d<\/em>-subshell. \u00a0<\/span><\/p>\n

    Note that due to the energy levels of the different subshells, that the only subshells that can exist in the valence shell are the s-<\/em> and p<\/em>-subshells.\u00a0 Both the d-<\/em> and the f-<\/em>subshells only fill after the s-<\/em>orbital of the next higher shell level.\u00a0 Thus, they can never contribute to the number of valence electrons.\u00a0 Since the s-<\/em>orbital can only house 2 electrons, and the p-<\/em>orbitals can house 6 electrons, the maximum number of valence electrons possible for any atom is 8.<\/span><\/p>\n

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    Concept Review<\/em><\/strong><\/span><\/h3>\n

    How many electrons are in the valence shell of Phosphorus?\u00a0 Phosphorus has an electron configuration of:<\/span><\/p>\n

    P = 1s2<\/sup>2s2<\/sup>2p6<\/sup>3s2<\/sup>3p3<\/sup><\/strong><\/span><\/h4>\n

    Solution<\/span><\/strong><\/em><\/h3>\n

    In the analysis of the electron configuration of phosphorus, we can see that the 3rd shell is the valence shell.\u00a0 For phosphorus there are 2 electrons in the 3s orbital, and there are 3 electrons in the 3p orbitals.\u00a0 This is a total of 5 electrons in the outer shell or 5 valence electrons. Note that due to the energy levels of the different subshells, that the only subshells that can exist in the valence shell are the s-<\/em> and p<\/em>-subshells<\/span><\/p>\n

    (Back to the Top)<\/em><\/strong><\/span><\/a><\/p>\n

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    Practice using the electron energy diagram to predict the electron configurations of the following atoms:<\/em><\/strong><\/span><\/h3>\n

    Boron, B<\/span><\/strong><\/h4>\n

    Carbon, C<\/span><\/strong><\/h4>\n

    Nitrogen, N<\/span><\/strong><\/h4>\n

    Oxygen, O<\/span><\/strong><\/h4>\n

    Fluorine, F<\/span><\/strong><\/h4>\n

    Neon, Ne<\/span><\/strong><\/h4>\n
    \n

    Solution: <\/strong><\/em><\/span><\/h3>\n

    B: 1s<\/em>2<\/sup>2s<\/em>2<\/sup>2p<\/em>1<\/sup><\/span><\/p>\n

    C: 1s<\/em>2<\/sup>2s<\/em>2<\/sup>2p<\/em>2<\/sup><\/span><\/p>\n

    N: 1s<\/em>2<\/sup>2s<\/em>2<\/sup>2p<\/em>3<\/sup><\/span><\/p>\n

    O: 1s<\/em>2<\/sup>2s<\/em>2<\/sup>2p<\/em>4<\/sup><\/span><\/p>\n

    F: 1s<\/em>2<\/sup>2s<\/em>2<\/sup>2p<\/em>5<\/sup><\/span><\/p>\n

    Ne: 1s<\/em>2<\/sup>2s<\/em>2<\/sup>2p<\/em>6<\/sup><\/span><\/p>\n

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    Electron Configurations and the Periodic Table<\/span><\/strong><\/h4>\n

    It is not always convenient to use the energy filling diagram to predict the electron configuration of the elements.\u00a0 Therefore, to be able to predict electron configurations more easily, the periodic table has been arranged to represent the electron orbital filling patterns.\u00a0 This is what gives the periodic table its stair-step or jagged appearance. Once you know how to read the periodic table for electron configurations, it becomes very easy to write out the electron configuration of any element. Figure 2.14 shows a version of the periodic table that depicts the layout of the electron configurations.<\/span><\/p>\n

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    Figure 2.14 The Periodic Table Corresponds to the Filling Patterns of the Electron Shells and Subshells.<\/strong><\/span><\/p>\n

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    This organization on the periodic table makes it easy to predict where an atom’s electron configuration ends and explains why elements within a family share so many similar characteristics!\u00a0 This is because they all share the same outer shell or valence shell electron configuration. <\/span><\/p>\n