Chapter 01 Atomic Structure and Periodic Table.pdf

Chapter 1 Atomic Structure and Periodic Table Figure 1.1 Minerals that your body needs Chapter Outline 1.1 Chemistry in Everyday Life 1.2 Matter 1.3 Classifying Matter 1.4 The Structure of Atom 1.5 The Periodic Table 1.6 Features of the Periodic Table 1.7 To Your Health: Transition Metals in the B...

Chapter 1 Atomic Structure and Periodic Table Figure 1.1 Minerals that your body needs Chapter Outline 1.1 Chemistry in Everyday Life 1.2 Matter 1.3 Classifying Matter 1.4 The Structure of Atom 1.5 The Periodic Table 1.6 Features of the Periodic Table 1.7 To Your Health: Transition Metals in the Body 1.8 Atomic Orbitals and Their Energies 1.9 Lewis Electron Dot Diagrams 1.10 Periodic trends in atomic properties 2 Chapter 1 | Atomic Structure and Periodic Table Introduction Your alarm goes off and, after hitting “snooze” once or twice, you pry yourself out of bed. You make a cup of coffee to help you get going, and then you shower, get dressed, eat breakfast, and check your phone for messages. On your way to school, you stop to fill your car’s gas tank, almost making you late for the first day of chemistry class. As you find a seat in the classroom, you read the question projected on the screen: “Welcome to class! Why should we study chemistry?” Do you have an answer? You may be studying chemistry because it fulfills an academic requirement, but if you consider your daily activities, you might find chemistry interesting for other reasons. Most everything you do and encounter during your day involves chemistry. Making coffee, cooking eggs, and toasting bread involve chemistry. The products you use—like soap and shampoo, the fabrics you wear, the electronics that keep you connected to your world, the gasoline that propels your car—all of these and more involve chemical substances and processes. Whether you are aware or not, chemistry is part of your everyday world. In this course, you will learn many of the essential principles underlying the chemistry of modern-day life. 1.1 Chemistry in Everyday Life Learning Objectives By the end of this section, you will be able to:  Recognize the breadth, depth, and scope of chemistry.  Provide examples of the importance of chemistry in everyday life. 3 Chapter 1 | Atomic Structure and Periodic Table Chemistry is the study of matter and the changes that material substances undergo. Of all the scientific disciplines, it is perhaps the most extensively connected to other fields of study. Geologists who want to locate new mineral or oil deposits use chemical techniques to analyze and identify rock samples. Oceanographers use chemistry to track ocean currents, determine the flux of nutrients into the sea, and measure the rate of exchange of nutrients between ocean layers. Engineers consider the relationships between the structures and the properties of substances when they specify materials for various uses. Physicists take advantage of the properties of substances to detect new subatomic particles. Astronomers use chemical signatures to determine the age and distance of stars and thus answer questions about how stars form and how old the universe is. The entire subject of environmental science depends on chemistry to explain the origin and impacts of phenomena such as air pollution, ozone layer depletion, and global warming. Examples of the practical applications of chemistry are everywhere (Figure 1.2 "Chemistry in Everyday Life"). Figure 1.2 Chemistry in Everyday Life The disciplines that focus on living organisms and their interactions with the physical world rely heavily on biochemistry, the application of chemistry to the study of biological processes. 4 Chapter 1 | Atomic Structure and Periodic Table Fields such as medicine, pharmacology, nutrition, and toxicology focus specifically on how the chemical substances that enter our bodies interact with the chemical components of the body to maintain our health and well-being. For example, in the specialized area of sports medicine, a knowledge of chemistry is needed to understand why muscles get sore after exercise as well as how prolonged exercise produces the euphoric feeling known as “runner’s high.” 1.2 Matter Learning Objectives By the end of this section, you will be able to:  Classify matter.  Describe the basic properties of each physical state of matter: solid, liquid, and gas. Chemists study the structures, physical properties, and chemical properties of material substances. These consist of matter, which is anything that occupies space and has mass. Gold and iridium are matter, as are peanuts, people, and postage stamps. Smoke, smog, and laughing gas are matter. Energy, light, and sound, however, are not matter; ideas and emotions are also not matter. The mass of an object is the quantity of matter it contains.  EXAMPLE 1.2 Which of the following are examples of matter? a) a baby b) an idea c) the Empire State Building d) an emotion 5 Chapter 1 | Atomic Structure and Periodic Table e) the air f) Alpha Centauri, the closest known star (excluding the sun) to our solar system Solution a) matter b) not matter. c) matter d) not matter. e) matter f) matter Under normal conditions, there are three distinct states of matter: solids, liquids, and gases. Solids, liquids, and gases are the three states of matter commonly found on earth (Figure 1.3 “states of matter”). A solid is rigid and possesses a definite shape. A liquid flows and takes the shape of its container, except that it forms a flat or slightly curved upper surface when acted upon by gravity. (In zero gravity, liquids assume a spherical shape.) Both liquid and solid samples have volumes that are very nearly independent of pressure. A gas takes both the shape and volume of its container. Figure 1.3 The three most common states or phases of matter are solid, liquid, and gas. 6 Chapter 1 | Atomic Structure and Periodic Table Link to Learning Watch this video (https://www.youtube.com/watch?v=wclY8FUoTE) to see the three states of matter. The following website (States of Matter (colorado.edu) ) shows the three states of matter. 1.3 Classifying Matter Learning Objectives By the end of this section, you will be able to:  Separate physical from chemical properties and changes.  Classify matter as an element, compound, homogeneous mixture, or heterogeneous mixture.  Matter can be classified into several categories. Two broad categories are mixtures and pure substances. A pure substance has a constant composition. They may be divided into two classes: elements and compounds. Pure substances that cannot be broken down into simpler substances by chemical changes are called elements. Iron, silver, gold, aluminum, sulfur, oxygen, and copper are familiar examples of the more than 100 known elements. A compound is a pure substance composed of two or more different atoms chemically bonded to one another. When heated in the absence of air, the compound sucrose is broken down into the element carbon and the compound water. A mixture is composed of two or more types of matter that can be 7 Chapter 1 | Atomic Structure and Periodic Table present in varying amounts and can be separated by physical changes, such as evaporation. A mixture with a composition that varies from point to point is called a heterogeneous mixture. A homogeneous mixture is a combination of two or more substances that is so intimately mixed that the mixture behaves as a single substance. Another word for a homogeneous mixture is solution. (Figure 1.4 “Types of Mixtures”) Figure 1.4 (a) Oil and vinegar salad dressing is a heterogeneous mixture. (b) A commercial sports drink is a homogeneous mixture. (credit a “left”: modification of work by John Mayer; credit a “right”: modification of work by Umberto Salvagnini; credit b “left: modification of work by Jeff Bedford). A summary of how to distinguish between the various major classifications of matter is shown in (Figure 1.5). Figure 1.5 Depending on its properties, a given substance can be classified as a homogeneous mixture, a heterogeneous mixture, a compound, or an element. 8 Chapter 1 | Atomic Structure and Periodic Table Link to Learning Watch this video (https://www.youtube.com/watch?v=anuDPYoG4HM) to see the Classification of Matter. Watch this video (https://www.youtube.com/watch?v=dggHWvFJ8Xs) to see the Types of Matter: Elements, Compounds, and Mixtures. 1.4 The Structure of Atom Learning Objectives By the end of this section, you will be able to: Identify the subatomic particles that make up of an atom and R LEARNING their properties.  Identify atoms of an element based on the number of protons in the nucleus.  Describe the different isotopes of an element. The basic building block of all matter is the atom (from the Greek Atomos, meaning “indivisible”). An atom is the smallest particle of an element that has the properties of that element and can enter into a chemical combination. It was learned that an atom contains a very small nucleus composed of positively charged protons and uncharged neutrons, surrounded by a much larger volume of space containing negatively charged electrons. The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume. (Figure 1.6“The Structure of the Atom”) 9 Chapter 1 | Atomic Structure and Periodic Table Figure 1.6 The Structure of the Atom When describing the properties of tiny objects such as atoms, we use appropriately small units of measure, such as the atomic mass unit (amu). Thus, 1 amu is exactly 1/12 of the mass of one carbon12 atom: 1 amu = 1.6605 ×× 10−24 g. Some properties of these subatomic particles are summarized in Table 1.1 "Properties of Subatomic Particles", which illustrates three important points. 1. Electrons and protons have electrical charges that are identical in magnitude but opposite in sign. We usually assign relative charges of −1 and +1 to the electron and proton, respectively. 2. Neutrons have approximately the same mass as protons but no charge. They are electrically neutral. 3.The mass of a proton or a neutron is about 1836 times greater than the mass of an electron. 4. Electrons, protons and neutrons constitute by far the bulk of the mass of atoms. 10 Chapter 1 | Atomic Structure and Periodic Table Table 1.1 Properties of Subatomic Particles Symbol Mass (approx.; Name Charge kg) Proton p+ 1.6 × 10−27 1+ Neutron n, n0 1.6 × 10−27 none Electron e− 9.1 × 10−31 1- The identity of an element is defined by its atomic number (Z), the number of protons in the nucleus of an atom of the element. The atomic number is therefore different for each element. The chemistry of each element is determined by its number of protons and electrons. In a neutral atom, the number of electrons equals the number of protons. Neutral atoms have the same number of electrons as they have protons, so their overall charge is zero. However, as we shall see later, this will not always be the case. (Table 1.2 “Selected Atomic Masses of Some Elements”) lists the atomic masses of some elements. Table 1.2 Selected Atomic Masses of Some Elements Element Name Atomic Mass (u) Aluminum 26.981 Argon 39.948 Arsenic 74.922 Barium 137.327 The mass number of an atom is the sum of the numbers of protons and neutrons in the nucleus. Given the mass number for a nucleus (and knowing the atomic number of that particular atom), you can determine the number of neutrons by subtracting the atomic number from the mass number. The element carbon (C) has an 11 Chapter 1 | Atomic Structure and Periodic Table atomic number of 6, which means that all neutral carbon atoms contain 6 protons and 6 electrons. (Figure 1.7) Figure 1.7 The Structure of the Atom Link to Learning Watch this video (https://www.youtube.com/watch?v=03iWCjxjCdA) to see the Structure of an Atom. Watch this video (https://www.youtube.com/watch?v=YKZv9bsFD3w) to understand Atomic Number and Atomic Mass. The (Build an Atom) simulator (https://phet.colorado.edu/sims/html/build-an-atom/latest/build-anatom_en.html) lets you use the number of protons, neutrons. and electrons to draw a model of an atom. 12 Chapter 1 | Atomic Structure and Periodic Table  EXAMPLE 1.2 What is the number of protons in the nucleus of each element? a) aluminum b) iron c) carbon Solution According to the periodic table, a) aluminum has an atomic number of 13. Therefore, every aluminum atom has 13 protons in its nucleus. b) Iron has an atomic number of 26. Therefore, every iron atom has 26 protons in its nucleus. c) Carbon has an atomic number of 6. Therefore, every carbon atom has 6 protons in its nucleus.  EXAMPLE 1.3 How many electrons are present in the atoms of each element? a) sulfur b) tungsten c) argon Solution a) The atomic number of sulfur is 16. Therefore, in a neutral atom of sulfur, there are 16 electrons. b) The atomic number of tungsten is 74. Therefore, in a neutral atom of tungsten, there are 74 electrons. c) The atomic number of argon is 18. Therefore, in a neutral atom of argon, there are 18 electrons. C 13 Chapter 1 | Atomic Structure and Periodic Table Medicinal Chemistry Iodine Iodine is an essential trace element in our diet; it is needed to produce thyroid hormone. Insufficient iodine in the diet can lead to the development of a goiter, an enlargement of the thyroid gland (Figure a,b). a) Insufficient iodine in the diet can cause an enlargement of the thyroid gland called a goiter. (b) The addition of small amounts of iodine to salt, which prevents the formation of goiters, has helped to eliminate this concern in the US where salt consumption is high. (Credit a: modification of work by “Almazi”/Wikimedia Commons; credit b: modification of work by Mike Mozart). The addition of small amounts of iodine to table salt (iodized salt) has essentially eliminated this health concern in the United States, but as much as 40% of the world’s population is still at risk of iodine deficiency. The iodine atoms are added as anions, and each has a 1− charge and a mass number of 127. 14 Chapter 1 | Atomic Structure and Periodic Table Isotopes Atoms of the same element can have different numbers of neutrons, however. Atoms of the same element (i.e., atoms with the same number of protons) with different numbers of neutrons are called isotopes. Most naturally occurring elements exist as isotopes. For example, most hydrogen atoms have a single proton in their nucleus. However, a small number (about one in a million) of hydrogen atoms have a proton and a neutron in their nuclei. This isotope particular isotope of hydrogen is called deuterium. A very rare form of hydrogen has one proton and two neutrons in the nucleus; this isotope of hydrogen is called tritium. Figure 1.8 "Isotopes of Hydrogen". Figure 1.8 Isotopes of Hydrogen. Most hydrogen atoms have only a proton in the nucleus (a). A small amount of hydrogen exists as the isotope deuterium, which has one proton and one neutron in its nucleus (b). A tiny amount of the hydrogen isotope tritium, with one proton and two neutrons in its nucleus, also exists on Earth (c). The nuclei and electrons are proportionately much smaller than depicted here. Link to Learning Watch this video (https://www.youtube.com/watch?v=EboWeWmh5Pg) to see; What are Isotopes. The (Isotopes and the atomic mass) simulator (https://phet.colorado.edu/sims/html/isotopes-and-atomicmass/latest/isotopes-and-atomic-mass_en.html) lets you know how to identify one isotope from another. 15 Chapter 1 | Atomic Structure and Periodic Table  LEARNING CHECK a) Why is the atomic number so important to the identity of an atom? b) What is the relationship between the number of protons and the number of electrons in an atom? c) How do isotopes of an element differ from each other? d) What is the mass number of an element?  LEARNING CHECK State the number of protons, neutrons, and electrons in neutral atoms of each isotope. a) 3 H b) 135 Cs c) 54 Fe d) 206 Pb 1.5 The Periodic Table Learning Objectives By the end of this section, you will be able to:  Explain how elements are organized into the periodic table.  Describe how some characteristics of elements relate to their positions on the periodic table. 16 Chapter 1 | Atomic Structure and Periodic Table 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? 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. When elements were later discovered whose properties closely matched Mendeleev’s 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. The periodic table is one of the cornerstones of chemistry because it organizes all the known elements based on their chemical properties. A modern version is shown in (Figure 1.9 "A Modern Periodic Table"). Most periodic tables provide additional data (such as atomic mass) in a box that contains each element’s symbol. The elements are listed in order of atomic number. 17 Chapter 1 | Atomic Structure and Periodic Table Figure 1.9 A Modern Periodic Table Each element name is abbreviated as a one- or two-letter chemical symbol (Table 1.3 “Element Names and Symbols”). By convention, the first letter of a chemical symbol is a capital letter, while the second letter (if there is one) is a lowercase letter. The first letter of the symbol is usually the first letter of the element’s name, while the second letter is some other letter from the name. Some elements have symbols that derive from earlier, mostly Latin names, so the symbols may not contain any letters from the English name. 18 Chapter 1 | Atomic Structure and Periodic Table Table 1.3 Element Names and Symbols (*The symbol comes from the Latin name of element.) Element Name Symbol Element Name Symbol Aluminum Al Magnesium Mg Argon Ar Manganese Mn Arsenic As Iron (Ferrum)* Fe Barium Ba Neon Ne Bismuth Bi Nickel Ni Boron B Nitrogen N Bromine Br Oxygen O Calcium Ca Phosphorus P Carbon C Tin (Stannum)* Sn Link to Learning Watch this video (https://www.youtube.com/watch?v=uPkEGAHo78o) to see the Periodic Table Introduction. Watch this video (https://www.youtube.com/watch?v=NPfOPOa5L30&t=1s to see the Periodic Table of Elements Watch this video (https://www.youtube.com/watch?v=KkUKVrpAlb8) to see the 8 Lesser-Known, Useful Elements 19 Chapter 1 | Atomic Structure and Periodic Table 1.6 Features of the Periodic Table Learning Objectives By the end of this section, you will be able to:  Predict the general properties of elements based on their location within the periodic table.  Classify them as metals, nonmetals, or metalloids. Elements that have similar chemical properties are grouped in columns called groups (or families). As well as being numbered, some of these groups have names—for example, alkali metals (the first column of elements), alkaline earth metals (the second column of elements), halogens (the next-to-last column of elements), and noble gases (the last column of elements). The alkali metals are lithium, sodium, potassium, rubidium, cesium, and francium. Hydrogen is unique in that it is generally placed in group 1, but it is not a metal. The compounds of the alkali metals are common in nature and daily life. One example is table salt (sodium chloride); lithium compounds are used in greases, in batteries, and as drugs to treat patients who exhibit manic-depressive, or bipolar, behavior. Although lithium, rubidium, and cesium are relatively rare in nature, and francium is so unstable and highly radioactive that it exists in only trace amounts, sodium and potassium are the seventh and eighth most abundant elements in Earth’s crust, respectively. 20 Chapter 1 | Atomic Structure and Periodic Table The Alkaline Earth Metals The alkaline earth metals are beryllium, magnesium, calcium, strontium, barium, and radium. Beryllium, strontium, and barium are rather rare, and radium is unstable and highly radioactive. In contrast, calcium and magnesium are the fifth and sixth most abundant elements on Earth, respectively; they are found in huge deposits of limestone and other minerals. The Halogens The halogens are fluorine, chlorine, bromine, iodine, and astatine. The name halogen is derived from the Greek for “salt forming,” which reflects that all the halogens react readily with metals to form compounds, such as sodium chloride and calcium chloride (used in some areas as road salt). Compounds that contain the fluoride ion are added to toothpaste and the water supply to prevent dental cavities. Fluorine is also found in Teflon coatings on kitchen utensils. Although chlorofluorocarbon propellants and refrigerants are believed to lead to the depletion of Earth’s ozone layer and contain both fluorine and chlorine, the latter is responsible for the adverse effect on the ozone layer. Bromine and iodine are less abundant than chlorine, and astatine is so radioactive that it exists in only negligible amounts in nature. The Noble Gases The noble gases are helium, neon, argon, krypton, xenon, and radon. Because the noble gases are composed of only single atoms, they are monatomic. At room temperature and pressure, they are unreactive gases. Because of their lack of reactivity, for many years they were called inert gases or rare gases. However, the first chemical compounds containing the noble gases were prepared in 1962. Although the noble gases are relatively minor constituents of the atmosphere, natural gas contains substantial amounts of helium. Because of its low reactivity, argon is often used as an unreactive (inert) atmosphere for welding and in light bulbs. The red light emitted by neon in a gas discharge tube is used in neon lights. 21 Chapter 1 | Atomic Structure and Periodic Table Each row of elements on the periodic table is called a period. Periods have different lengths; the first period has only 2 elements (hydrogen and helium), while the second and third periods have 8 elements each. The fourth and fifth periods have 18 elements each, and later periods are so long that a segment from each is removed and placed beneath the main body of the table. Certain elemental properties become apparent in a survey of the periodic table. Every element can be classified as either a metal, a nonmetal, or a semimetal, as shown in Figure 1.10 "Types of Elements". A metal is a substance that is shiny, typically (but not always) silvery in color, and an excellent conductor of electricity and heat. Metals are also malleable (they can be beaten into thin sheets) and ductile (they can be drawn into thin wires). A nonmetal is typically dull and a poor conductor of electricity and heat. Figure 1.10 Types of Elements Solid nonmetals are also very brittle. As shown in Figure 1.10 "Types of Elements", metals occupy the left three-fourths of the 22 Chapter 1 | Atomic Structure and Periodic Table periodic table, while nonmetals (except for hydrogen) are clustered in the upper right-hand corner of the periodic table. The elements with properties intermediate between those of metals and nonmetals are called semimetals (or metalloids). Elements adjacent to the bold line in the right-hand portion of the periodic table have semimetal properties. Another way to categorize the elements of the periodic table is shown in Figure 1.11 "Special Names for Sections of the Periodic Table". The first two columns on the left and the last six columns on the right are called the main group elements. The ten-column block between these columns contains the transition metals. The two rows beneath the main body of the periodic table contain the inner transition metals. The elements in these two rows are also referred to as, respectively, the lanthanide metals and the actinide metals. Figure 1.11 Special Names for Sections of the Periodic Table 23 Chapter 1 | Atomic Structure and Periodic Table Link to Learning Watch this video (https://www.youtube.com/watch?v=uPkEGAHo78o) to see the Periodic Table Introduction. Watch this video (https://www.youtube.com/watch?v=NPfOPOa5L30&t=1s to see the Periodic Table of Elements. Watch this video (https://www.youtube.com/watch?v=KkUKVrpAlb8) to see the 8 Lesser-Known, Useful Elements. Watch this video (https://www.youtube.com/watch?v=rrDJ8Ri3pXk) to see the Strontium: It Knows Where You've Been  EXAMPLE 1.4 Classify each element as a metal, a semimetal, or a nonmetal. If a metal, state whether it is an alkali metal, an alkaline earth metal, or a transition metal. a) iron b) tantalum c) sulfur d) silicon e) chlorine f) nickel g) potassium h) radon i) zirconium 24 Chapter 1 | Atomic Structure and Periodic Table Solution Fe metal: transition metal Ta metal: transition metal. S nonmetal Si semimetal Cl nonmetal (halogen) Ni metal: transition metal K metal: alkali metal Rn nonmetal (noble gas) Zr metal: transition metal  LEARNING CHECK Which of these sets of elements are all in the same group? a) iron, ruthenium, and osmium b) nickel, palladium, and lead c) iodine, fluorine, and oxygen d) boron, aluminum, and gallium 1.7 To Your Health: Transition Metals in the Body Learning Objectives By the end of this section, you will be able to:  Become familiar with some of the roles of transition-metal complexes in biological systems. Most of the elemental composition of the human body consists of main group elements (Table 1.4 “Elemental Composition of a Human Body”). The first element appearing on the list that is not a main group element is iron, at 0.006 percentage by mass. Iron is a transition metal. The chemistry of iron makes it a key component in the proper functioning of red blood cells. Red blood cells are 25 Chapter 1 | Atomic Structure and Periodic Table cells that transport oxygen from the lungs to cells of the body and then transport carbon dioxide from the cells to the lungs. Without red blood cells, animal respiration as we know it would not exist. The critical part of the red blood cell is a protein called hemoglobin. Hemoglobin combines with oxygen and carbon dioxide, transporting these gases from one location to another in the body. The crucial atom in the hemoglobin protein is iron. Each hemoglobin molecule has four iron atoms, which act as binding sites for oxygen. It is the presence of this particular transition metal in your red blood cells that allows you to use the oxygen you inhale. Other transition metals have important functions in the body, despite being present in low amounts. Zinc is needed for the body’s immune system to function properly, as well as for protein synthesis and tissue and cell growth. Copper is also needed for several proteins to function properly in the body. Manganese is needed for the body to metabolize oxygen properly. Cobalt is a necessary component of vitamin B-12, a vital nutrient. Table 1.4 Elemental Composition of a Human Body Element Percentage by Mass Oxygen 61 Carbon 23 Hydrogen 10 Nitrogen 2.6 Calcium 1.4 Phosphorus 1.1 Sulfur 0.2 Potassium 0.2 Sodium 0.14 Chlorine 0.12 26 Chapter 1 | Atomic Structure and Periodic Table Magnesium 0.027 Silicon 0.026 Iron 0.006 Fluorine 0.0037 Zinc 0.0033 All others 0.174 Source: D. R. Lide, ed. CRC Handbook of Chemistry and Physics, 89th ed. (Boca Raton, FL: CRC Press, 2008–9), 7–24. Link to Learning Watch this video (https://www.youtube.com/watch?v=g2BPJ7nO1jk&t=1s) to see the Elements of the Human Body Watch this video (https://www.youtube.com/watch?v=7QXKBVpQ7GA ) to see the elements that make up the human body 1.8 Atomic Orbitals and Their Energies Learning Objectives By the end of this section, you will be able to:  Describe the distribution of electrons into shells, subshells, and orbitals around the nucleus of an atom.  Use the rules of an electronic configuration to describe how electrons are distributed into specific orbitals of atoms and ions.  Identify the valence shell electrons for an atom. 27 Chapter 1 | Atomic Structure and Periodic Table 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? The modern theory of electron behavior is called quantum mechanics. It makes the following statements about electrons in atoms: • Electrons in atoms can have only certain specific energies. We say that the energies of the electrons are quantized. • Electrons are organized according to their energies into sets called shells (energy levels or principal quantum numbers, n= 1,2, 3…). Generally, the higher the energy of a shell, the farther it is (on average) from the nucleus. Shells do not have specific, fixed distances from the nucleus, but an electron in a higher-energy shell will spend more time farther from the nucleus than does an electron in a lower-energy shell. • Shells are further divided into subsets called subshells. The first shell has only one subshell, the second shell has two subshells, the third shell has three subshells, and so on. The subshells of each shell are labeled, in order, with the letters s, p, d, and f. Thus, the first shell has only an s subshell, the second shell has an s and p subshells, the third shell has s, p, and d subshells, and so forth. • Different subshells hold a different maximum number of electrons. Any s subshell can hold up to 2 electrons; p, 6; d, 10; and f, 14. • We use numbers to indicate which shell an electron is in. The first shell, closest to the nucleus and with the lowest-energy electrons, is shell (1). This first shell has only one subshell, which is labeled s and can hold a maximum of (2) electrons. We combine the shell and subshell labels when referring to the organization of electrons about a nucleus and use a superscript to indicate how many electrons are in a subshell. 28 Chapter 1 | Atomic Structure and Periodic Table • Electrons are typically organized around an atom by starting at the lowest possible quantum numbers first, which are the shellssubshells with lower energies. • For larger atoms, the order of filling the shells and subshells seems to become even more complicated. There are some useful ways to remember the order, like that shown in Figure 1.12 “Electron Shell Filling Order”. If you follow the arrows in order, they pass through the subshells in the order that they are filled with electrons in larger atoms. Figure 1.12 Electron Shell Filling Order • Thus, because a hydrogen atom has its single electron in the s subshell of the first shell, we use 1s1 to describe the electronic structure of hydrogen. This structure is called an electron configuration. Electron configurations are shorthand descriptions of the arrangements of electrons in atoms. The electron configuration of a hydrogen atom is spoken out loud as “one-essone.” 29 Chapter 1 | Atomic Structure and Periodic Table • Helium atoms have 2 electrons. Both electrons fit into the 1s subshell because s subshells can hold up to 2 electrons; therefore, the electron configuration for helium atoms is 1s2 (spoken as “one-ess-two”). • The 1s subshell cannot hold 3 electrons (because an s subshell can hold a maximum of 2 electrons), so the electron configuration for a lithium atom cannot be 1s3. Two of the lithium electrons can fit into the 1s subshell, but the third electron must go into the second shell. The second shell has two subshells, s and p, which fill with electrons in that order. The 2s subshell holds a maximum of 2 electrons, and the 2p subshell holds a maximum of 6 electrons. Because lithium’s final electron goes into the 2s subshell, we write the electron configuration of a lithium atom as 1s22s1. • The next largest atom, beryllium, has 4 electrons, so its electron configuration is 1s22s2. Now that the 2s subshell is filled, electrons in larger atoms start filling the 2p subshell. Thus, the electron configurations for the next six atoms are as follows: • B: 1s22s22p1 • C: 1s22s22p2 • N: 1s22s22p3 • O: 1s22s22p4 • F: 1s22s22p5 • Ne: 1s22s22p6 • With neon, the 2p subshell is completely filled. Because the second shell has only two subshells, atoms with more electrons now must begin the third shell. The third shell has three subshells, labeled s, p, and d. The d subshell can hold a maximum of 10 electrons. The first two subshells of the third shell are filled in order—for example, the electron configuration of aluminum, with 13 electrons, is 1s22s22p63s23p1. However, a curious thing happens after the 3p subshell is filled: the 4s subshell begins to fill before 30 Chapter 1 | Atomic Structure and Periodic Table the 3d subshell does. In fact, the exact ordering of subshells becomes more complicated at this point (after argon, with its 18 electrons), so we will not consider the electron configurations of larger atoms. • A fourth subshell, the f subshell, is needed to complete the electron configurations for all elements. An f subshell can hold up to 14 electrons.  EXAMPLE 1.5 What is the electron configuration of a neutral phosphorus atom? Solution A neutral phosphorus atom has 15 electrons. Two electrons can go into the 1s subshell, 2 can go into the 2s subshell, and 6 can go into the 2p subshell. That leaves 5 electrons. Of those 5 electrons, 2 can go into the 3s subshell, and the remaining 3 electrons can go into the 3p subshell. Thus, the electron configuration of neutral phosphorus atoms is 1s22s22p63s23p3 The atomic number of phosphorus is 15. Thus, a phosphorus atom contains 15 electrons. The order of filling of the energy levels is 1s, 2s, 2p, 3s, 3p, 4s, . . . The 15 electrons of the phosphorus atom are shown in the following diagram. Notice that the "3p" orbitals contain 3 electrons, one electron in each orbital.  LEARNING CHECK What is the electron configuration of a neutral chlorine atom? Chemistry results from interactions between the outermost shells of electrons on different atoms. Thus, it is convenient to separate electrons into two groups. Valence shell electrons (or, more simply, the valence electrons) are the electrons in the highest- 31 Chapter 1 | Atomic Structure and Periodic Table numbered shell, or valence shell, while core electrons are the electrons in lower-numbered shells. We can see from the electron configuration of a carbon atom—1s22s22p2—that it has 4 valence electrons (2s22p2) and 2 core electrons (1s2).  EXAMPLE 1.6 From the electron configuration of neutral phosphorus atoms in Example 7, how many valence electrons and how many core electrons does a neutral phosphorus atom have? Solution The highest-numbered shell is the third shell, which has 2 electrons in the 3s subshell and 3 electrons in the 3p subshell. That gives a total of 5 electrons, so neutral phosphorus atoms have 5 valence electrons. The 10 remaining electrons, from the first and second shells, are core electrons. In an external magnetic field, the electron has two possible orientations (Figure 1.13 "Electron Spin"). Any electron can have only two possible values, designated +½ (up) and −½ (down) to indicate that the two orientations are opposites. An electron behaves like a magnet that has one of two possible orientations, aligned either with the magnetic field or against it. 32 Chapter 1 | Atomic Structure and Periodic Table Figure 1.13 Electron Spin We have just seen that electrons fill orbitals in shells and subshells in a regular pattern, but why does it follow this pattern? There are three principles which should be followed to properly fill electron orbital energy diagrams: 1. The Aufbau principle 2. The Pauli exclusion principle 3. Hund’s rule The overall pattern of the electron shell filling order emerges from the Aufbau principle (German for “building up”): electrons fill the lowest energy orbitals first. Increasing the principal quantum number, n, increases orbital energy levels, as the electron density becomes more spread out away from the nucleus. Figure 1.14, demonstrates that each line represents an orbital, and each set of lines at the same energy represents a subshell of orbitals. 33 Chapter 1 | Atomic Structure and Periodic Table Figure 1.14 Generic energy diagram of orbitals in a multi-electron atom. the Pauli exclusion principle states that we can only fill each orbital with a maximum of two electrons of opposite spin. But how should we fill multiple orbitals of the same energy level within a subshell (e.g., The three orbitals in the 2p subshell)? Orbitals of the same energy level are known as degenerate orbitals, and we fill them using Hund’s rule: place one electron into each degenerate orbital first, before pairing them in the same orbital. Let’s examine a few examples to demonstrate the use of the three principles. Boron is atomic number 5, and therefore has 5 electrons. First fill the lowest energy 1s orbital with two electrons of opposite spin, then the 2s orbital with 2 electrons of opposite spin and finally place the last electron into any of the three degenerate 2p orbitals (Figure 1.15). 34 Chapter 1 | Atomic Structure and Periodic Table Figure 1.15 Boron electron configuration energy diagram Moving across the periodic table, we follow Hund’s rule and add an additional electron to each degenerate 2p orbital for each subsequent element (Figure 1.16). At oxygen we can finally pair up and fill one of the degenerate 2p orbitals. Figure 1.16 Electron configuration energy diagrams for carbon, nitrogen and oxygen. 35 Chapter 1 | Atomic Structure and Periodic Table  LEARNING CHECK 1. How many subshells are completely filled with electrons for Na? How many subshells are unfilled? 2. How many electrons in the valence shell for Mg? How many electrons in the core shells? 3. What is the maximum number of electrons in the entire n = 2 shell? 4. What is the maximum number of electrons in the entire n = 4 shell? Based on electron configurations, the periodic table can be divided into blocks denoting which sublevel is in the process of being filled. The s, p, d, and f blocks are illustrated below, Figure 1.17 “Blocks on the Periodic Table”. Figure 1.17 Blocks on the Periodic Table 36 Chapter 1 | Atomic Structure and Periodic Table Link to Learning Watch this video (https://www.youtube.com/watch?v=wAagHeE2SI0) to see Electronic Configuration of Elements The following website (e-1 Configuration Menu Page (fscj.me)) provides practice with the electron configuration of an atom Watch this video (https://www.youtube.com/watch?v=K2JwCPBgn2g) to see spdf Block Elements 1.9 Lewis Electron Dot Diagrams Learning Objectives By the end of this section, you will be able to:  Draw the Lewis dot symbols to represent the valence electrons for a given atom. A Lewis electron dot diagram (or electron dot diagram or a Lewis diagram or a Lewis structure) is a representation of the valence electrons of an atom that uses dots around the symbol of the element. The number of dots equals the number of valence electrons in the atom. These dots are arranged to the right and left and above and below the symbol, with no more than two dots on a side. (It does not matter what order the positions are used.) For example, the Lewis electron dot diagram for hydrogen is simply 37 Chapter 1 | Atomic Structure and Periodic Table Because the side is not important, the Lewis electron dot diagram could also be drawn as follows: The electron dot diagram for helium, with two valence electrons, is as follows: Fluorine and neon have seven and eight dots, respectively:  EXAMPLE 1.7 What is the Lewis electron dot diagram for each element? a) Aluminum b) Sulfur Solution a) b)   LEARNING CHECK 1. Represent the valence electrons of an atom that uses dots around the symbol of the element for group 3. 2. How you can relate the group number and Lewis dots formula. 38 Chapter 1 | Atomic Structure and Periodic Table Link to Learning Watch this video (https://www.youtube.com/watch?v=vOIX4UQ0jJU) to see how to draw Lewis Dot Symbols Watch this video (https://www.youtube.com/watch?v=79l5ucPy2RY) to see Lewis Structures of Atoms 1.10 Periodic trends in atomic properties Learning Objectives By the end of this section, you will be able to:  Use the electron arrangement of elements to explain the trends in periodic properties. Atomic Radius The periodic table is useful for understanding atomic properties that show periodic trends. One such property is the atomic radius (Figure 1.18 "Trends on the Periodic Table"). As mentioned earlier, the higher the shell number, the farther from the nucleus the electrons in that shell are likely to be. In other words, the size of an atom is generally determined by the number of the valence electron shell. Therefore, as we go down a column on the 39 Chapter 1 | Atomic Structure and Periodic Table periodic table, the atomic radius increases. As we go across a period on the periodic table, however, electrons are being added to the same valence shell; meanwhile, more protons are being added to the nucleus, so the positive charge of the nucleus is increasing. The increasing positive charge attracts the electrons more strongly, pulling them closer to the nucleus. Consequently, as we go across a period, the atomic radius decreases. These trends are seen clearly in Figure 1.18 "Trends on the Periodic Table". Figure 1.18 Trends on the Periodic Table (The relative sizes of the atoms show several trends with regard to the structure of the periodic table. Atoms become larger going down a column and smaller going across a period.) 40 Chapter 1 | Atomic Structure and Periodic Table  EXAMPLE 1.8 Using the periodic table (Figure 1.18 "Trends on the Periodic Table"), which atom is larger? N or Bi Mg or Cl Solution Because Bi is below N on the periodic table and has electrons in higher-numbered shells, we expect that Bi atoms are larger than N atoms. Both Mg and Cl are in period 3 of the periodic table, but Cl lies farther to the right. Therefore, we expect Mg atoms to be larger than Cl atoms.  LEARNING CHECK 1.Which elements have chemical properties similar to those of magnesium? a) sodium b) calcium c) barium d) selenium 2.In each pair of atoms, which atom has the greater atomic radius? a) H and Li b) N and P c) Cl and Ar d) Al and Cl Ionization energy (IE) is the amount of energy required to remove an electron from an atom in the gas phase? IE is usually expressed in kJ/mol of atoms. It is always positive because the removal of an electron always requires that energy be put in (i.e., it is endothermic). IE also shows periodic trends. As you go down the periodic table, it becomes easier to remove an electron from an atom (i.e., IE 41 Chapter 1 | Atomic Structure and Periodic Table decreases) because the valence electron is farther away from the nucleus. However, as you go across the periodic table and the electrons get drawn closer in, it takes more energy to remove an electron; as a result, IE increases. Figure 1.19 " The first ionization energy of the elements ". X(g)⟶X+(g) + e− ( IE ) Figure 1.19 The first ionization energy of the elements. Electronegativity The ability of an atom in a molecule to attract a shared electron pair to itself, forming a polar covalent bond, is called its electronegativity. Imagine a game of tug-of-war. If the two teams are of equal strength, the rope stays centered. If one team is stronger, the rope is pulled in that team's direction. If one team is overwhelmingly stronger, the weaker team is no longer able to hold onto the rope and the entire rope ends up on the side of the stronger team. This is analogous to chemical bonds. If the two atoms of the bond are of equal electronegativity, the electrons are equally shared. If one atom is more electronegative, the electrons of the bond are 42 Chapter 1 | Atomic Structure and Periodic Table more attracted to that atom. If one atom is overwhelmingly more electronegative than the other atom, the electrons will not be shared and an ionic bond will result. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to cesium and francium which are the least electronegative at 0.7. Electronegativities increase from left to right across the periodic table Figure 1.20 “Electronegativity trend across the periodic table ". Elements on the left of the periodic table have low electronegativities and are often called electropositive elements. The order of electronegativities F > O > N > C is an important property that we will use to explain the chemical properties of organic compounds. Electronegativities decrease from top to bottom within a group of elements. The order of decreasing electronegativities F > Cl > Br > I is another sequence that we will use to interpret the chemical and physical properties of organic compounds. Figure 1.20 Electronegativity trend across the periodic table. 43 Chapter 1 | Atomic Structure and Periodic Table Metallic Character (MC) The metallic character of an element can be defined as how readily an atom can lose an electron. Metallic characteristics increase from right to left across a period. This is caused by the increase in radius of the atom that allows the outer electrons to ionize more readily. Metallic characteristics increase down a group. Electron shielding causes the atomic radius to increase thus the outer electrons ionizes more readily than electrons in smaller atoms. Figure 1.21 " Metallic Character Trend "). Figure 1.21 Metallic Character Trend.  EXAMPLE 1.9 a) Which has more metallic character, Lead (Pb) or Tin (Sn)? b) Which element is more electronegative, sulfur (S) or selenium (Se)? c) Why is the electronegativity value of most noble gases zero? 44 Chapter 1 | Atomic Structure and Periodic Table Solution a) Lead (Pb) b) Sulfur (S) c) Most noble gases have full valence shells. Electron Affinity (EA) The amount of energy released when an electron is added to a neutral atom to form an anion. X(g) + e− → X−(g) + EA Electron affinity increases going left to right across a period. Going down the group the electron affinity should decrease. Figure 1.22 Electron Affinity Trend.  LEARNING CHECK 1. Which element in group 4 has higher metallic characteristic? 2. Which element in group 2 has higher electronegativity? 45 Chapter 1 | Atomic Structure and Periodic Table Link to Learning Watch this video (https://www.youtube.com/watch?v=hePb00CqvP0) to see The Periodic Table: Atomic Radius, Ionization Energy, and Electronegativity trends Watch this video (https://www.youtube.com/watch?v=dss59DrvIk) to see Trends in the Periodic Table 46 Chapter 1 | Atomic Structure and Periodic Table PROBLEMS Choose the correct statement for each of the following questions: 1. Dot diagrams are used to represent. A. Atomic numbers B. Atomic mass C. Isotopes D. Outer level electrons 2. How many valence electrons does carbon have? A. 12 B. 6 C. 4 D. 2 3. Which element is found in period 4 and group 2? A. Helium B. Beryllium C. Calcium D. Strontium 4. In order to determine the number of neutrons A. It is always equal to number of protons B. It is always equal to number of electrons C. Atomic mass minus atomic number D. Subtract electrons from atomic number 5. Why is group 1 highly reactive? A. It has the most energy levels B. It has 1 valence electron C. It has the same number of protons D. It has the same number of electrons 6. The elements potassium, vanadium, manganese and calcium are examples of A. Members of a period B. Members of a family C. Molecules D. Gases 47 Chapter 1 | Atomic Structure and Periodic Table 7. Where are valence electrons found? A. In the outermost energy level B. In the innermost energy level C. In the nucleus D. In the electron cloud 8. The modern periodic table is organized with the elements in order of increasing _____. A. Atomic number B. Atomic mass C. Nuclear number D. Quantum number 9. Each row on the periodic table represents: A. An energy level B. A sublevel C. An electron D. An orbital 10. Energy levels are denoted by: A. Letters B. Numbers C. A combination of letters and numbers D. Subscripts 11. Which periodic table family could have electrons filling orbitals in the s-block? A. Alkaline Earth B. Halogens C. Noble Gases D. Transition Metals 12. What is the electron configuration of iodine? A. [Kr]5s24d105p6 B. [Kr]5s24d105p5 C. [Kr]5s25d106p5 D. [Kr]5s25d106p6 13. What atom matches this electron configuration? [Xe] 6s24f145d9 A. Mercury B. Gold C. Platinum D. Thallium 48 Chapter 1 | Atomic Structure and Periodic Table 14. How many atomic orbitals are there in the p-sublevel? A. 2 B. 3 C. 4 D. 5 15. What is the noble gas configuration for phosphorus? A. [Ar] 3p5 B. [He] 3s2 3p5 C. [Ne] 3s2 3p3 D. [Na] 3s2 3p5 16. Sodium (Na) and potassium (K) are in the same group on the periodic table. Based on their locations, which statement about sodium and potassium is true? A. Sodium is less electronegative than potassium. B. Sodium has fewer energy levels than potassium. C. Sodium has a larger ionic radius than potassium. D. Sodium has lower ionization energy than potassium. 17. According to the periodic table, which of the following indicates the correct decreasing order of ionization energy? A. Li > Na > K > Cs B. Na > K > Li > Cs C. Li > K > Na > Cs D. Cs > K > Na > Li 18. Which atom has the largest atomic radius? A. Potassium B. Rubidium C. Francium D. Cesium 19. Iron-60 (atomic number 26) is an isotope that is often used to study meteorites. How many neutrons does iron-60 have? A. 26 B. 60 C. 34 D. 86 49 Chapter 1 | Atomic Structure and Periodic Table 20. Which natural element is the most metallic? A. Francium B. Cesium C. Helium D. Cobalt ADDITIONAL PROBLEMS 1. Which atom has the electron configuration 2 2 6 2 6 2 10 6 2 2 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d ? 2. Which atom has the electron configuration 2 2 6 2 6 7 2 1s 2s 2p 3s 3p 3d 4s ? 3. Which of the following atoms contains only three valence electrons: Li, B, N, F, Ne? 4. Which of the following has only two unpaired electrons? Br, Si, S and P? 5. Cobalt–60 and iodine–131 are radioactive isotopes commonly used in nuclear medicine. How many protons, neutrons, and electrons are in atoms of these isotopes? Write the complete electron configuration for each isotope. 6. Based on their positions in the periodic table, predict which has the smallest atomic radius: Mg, Sr, Si, Cl, I. 7. Based on their positions in the periodic table, predict which has the largest atomic radius: Li, Rb, N, F, I. 8. Based on their positions in the periodic table, predict which has the largest first ionization energy: Mg, Ba, B, O, Te. 9. Atoms of which group in the periodic table have a valence shell electron configuration of ns2np3? 10. Atoms of which group in the periodic table have a valence shell electron configuration of ns2? 50 Chapter 1 | Atomic Structure and Periodic Table 11. In what way are isotopes of a given element always different? In what way(s) are they always the same? 12. Give the number of protons, electrons, and neutrons in neutral atoms of each of the following isotopes: (a) 105B (b) 19980Hg 13. Atoms of each of the following elements are essential for life. Give the group name for the following elements: (a) chlorine (b) calcium (c) sodium (d) sulfur 14. Use the periodic table to give the name and symbol for each of the following elements: (a) The noble gas in the same period as germanium (b) The alkaline earth metal in the same period as selenium (c) The halogen in the same period as lithium 15. Write a symbol for each of the following neutral isotopes. Include the atomic number and mass number for each. (a) The alkali metal with 11 protons and a mass number of 23 (b) The noble gas element with 75 neutrons in its nucleus and 54 electrons in the neutral atom (c) The isotope with 33 protons and 40 neutrons in its nucleus (d) The alkaline earth metal with 88 electrons and 138 neutrons 16. The electron configuration of the Ti atom is 1s22s22p63s23p64s23d2. How many valence electrons of Ti? 17. What might be the electron configuration of Cl−? 18. Describe the trends in atomic radii as related to an element’s position on the periodic table. 51 Chapter 1 | Atomic Structure and Periodic Table 19. What orbital is filled when iodine gains an electron to become a negative ion? 20. Positive ions are smaller than the atoms from which they are formed, but negative ions are larger than the atoms from which they are formed. Explain why this is so. 52 Chapter 1 | Atomic Structure and Periodic Table RESOURCES 1234567- https://openstax.org/details/books/chemistry-2e https://open.umn.edu/opentextbooks/subjects/chemistry https://open.umn.edu/opentextbooks/textbooks/generalchemistry-principles-patterns-and- applications https://open.umn.edu/opentextbooks/textbooks/the-basics-ofgeneral-organic-and-biological- chemistry https://open.umn.edu/opentextbooks/textbooks/introductorychemistry https://open.umn.edu/opentextbooks/textbooks/organicchemistry-with-a-biological-emphasis- volume-i https://www2.chemistry.msu.edu/faculty/reusch/VirtTxtJml/int ro1.htm https://opentextbc.ca/introductorychemistry/