1 324 Chapter 6 Electronic Structure and Periodic Properties of Elements 6.5 Periodic Variations in Element Properties By the end of this section, you will be able to: Describe and explain the observed trends in atomic size, ionization energy, and electron affinity of the elements The elements in groups (vertical columns) of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number and distribution of electrons in their valence shells. However, there are also other patterns in chemical properties on the periodic table. For example, as we move down a group, the metallic character of the atoms increases. Oxygen, at the top of group 16 (6A), is a colorless gas; in the middle of the group, selenium is a semiconducting solid; and, toward the bottom, polonium is a silver-grey solid that conducts electricity. As we go across a period from left to right, we add a proton to the nucleus and an electron to the valence shell with each successive element. As we go down the elements in a group, the number of electrons in the valence shell remains constant, but the principal quantum number increases by one each time. An understanding of the electronic structure of the elements allows us to examine some of the properties that govern their chemical behavior. These properties vary periodically as the electronic structure of the elements changes. They are (1) size (radius) of atoms and ions, (2) ionization energies, and (3) electron affinities. This content is available for free at
2 Chapter 6 Electronic Structure and Periodic Properties of Elements 325 Link to Learning Explore visualizations (http://openstaxcollege.org/l/16pertrends) of the periodic trends discussed in this section (and many more trends). With just a few clicks, you can create three-dimensional versions of the periodic table showing atomic size or graphs of ionization energies from all measured elements. Variation in Covalent Radius The quantum mechanical picture makes it difficult to establish a definite size of an atom. However, there are several practical ways to define the radius of atoms and, thus, to determine their relative sizes that give roughly similar values. We will use the covalent radius (Figure 6.31), which is defined as one-half the distance between the nuclei of two identical atoms when they are joined by a covalent bond (this measurement is possible because atoms within molecules still retain much of their atomic identity). We know that as we scan down a group, the principal quantum number, n, increases by one for each element. Thus, the electrons are being added to a region of space that is increasingly distant from the nucleus. Consequently, the size of the atom (and its covalent radius) must increase as we increase the distance of the outermost electrons from the nucleus. This trend is illustrated for the covalent radii of the halogens in Table 6.2 and Figure The trends for the entire periodic table can be seen in Figure Covalent Radii of the Halogen Group Elements Atom Covalent radius (pm) Nuclear charge F Cl Br I At Table 6.2
3 326 Chapter 6 Electronic Structure and Periodic Properties of Elements Figure 6.31 (a) The radius of an atom is defined as one-half the distance between the nuclei in a molecule consisting of two identical atoms joined by a covalent bond. The atomic radius for the halogens increases down the group as n increases. (b) Covalent radii of the elements are shown to scale. The general trend is that radii increase down a group and decrease across a period. This content is available for free at
4 Chapter 6 Electronic Structure and Periodic Properties of Elements 327 Figure 6.32 Within each period, the trend in atomic radius decreases as Z increases; for example, from K to Kr. Within each group (e.g., the alkali metals shown in purple), the trend is that atomic radius increases as Z increases. As shown in Figure 6.32, as we move across a period from left to right, we generally find that each element has a smaller covalent radius than the element preceding it. This might seem counterintuitive because it implies that atoms with more electrons have a smaller atomic radius. This can be explained with the concept of effective nuclear charge, Z eff. This is the pull exerted on a specific electron by the nucleus, taking into account any electron electron repulsions. For hydrogen, there is only one electron and so the nuclear charge (Z) and the effective nuclear charge (Z eff ) are equal. For all other atoms, the inner electrons partially shield the outer electrons from the pull of the nucleus, and thus: Z eff = Z shielding Shielding is determined by the probability of another electron being between the electron of interest and the nucleus, as well as by the electron electron repulsions the electron of interest encounters. Core electrons are adept at shielding, while electrons in the same valence shell do not block the nuclear attraction experienced by each other as efficiently. Thus, each time we move from one element to the next across a period, Z increases by one, but the shielding increases only slightly. Thus, Z eff increases as we move from left to right across a period. The stronger pull (higher effective nuclear charge) experienced by electrons on the right side of the periodic table draws them closer to the nucleus, making the covalent radii smaller. Thus, as we would expect, the outermost or valence electrons are easiest to remove because they have the highest energies, are shielded more, and are farthest from the nucleus. As a general rule, when the representative elements form cations, they do so by the loss of the ns or np electrons that were added last in the Aufbau process. The transition
5 328 Chapter 6 Electronic Structure and Periodic Properties of Elements elements, on the other hand, lose the ns electrons before they begin to lose the (n 1)d electrons, even though the ns electrons are added first, according to the Aufbau principle. Example 6.12 Sorting Atomic Radii Predict the order of increasing covalent radius for Ge, Fl, Br, Kr. Solution Radius increases as we move down a group, so Ge < Fl (Note: Fl is the symbol for flerovium, element 114, NOT fluorine). Radius decreases as we move across a period, so Kr < Br < Ge. Putting the trends together, we obtain Kr < Br < Ge < Fl. Check Your Learning Give an example of an atom whose size is smaller than fluorine. Answer: Ne or He Variation in Ionic Radii Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it is smaller than the atom from which it is derived (Figure 6.33). For example, the covalent radius of an aluminum atom (1s 2 2s 2 2p 6 3s 2 3p 1 ) is 118 pm, whereas the ionic radius of an Al 3+ (1s 2 2s 2 2p 6 ) is 68 pm. As electrons are removed from the outer valence shell, the remaining core electrons occupying smaller shells experience a greater effective nuclear charge Z eff (as discussed) and are drawn even closer to the nucleus. Figure 6.33 The radius for a cation is smaller than the parent atom (Al), due to the lost electrons; the radius for an anion is larger than the parent (S), due to the gained electrons. Cations with larger charges are smaller than cations with smaller charges (e.g., V 2+ has an ionic radius of 79 pm, while that of V 3+ is 64 pm). Proceeding down the groups of the periodic table, we find that cations of successive elements with the same charge generally have larger radii, corresponding to an increase in the principal quantum number, n. An anion (negative ion) is formed by the addition of one or more electrons to the valence shell of an atom. This results in a greater repulsion among the electrons and a decrease in Z eff per electron. Both effects (the increased number of electrons and the decreased Z eff ) cause the radius of an anion to be larger than that of the parent atom (Figure 6.33). For example, a sulfur atom ([Ne]3s 2 3p 4 ) has a covalent radius of 104 pm, whereas the ionic radius of the sulfide anion ([Ne]3s 2 3p 6 ) is 170 pm. For consecutive elements proceeding down any group, anions have larger principal quantum numbers and, thus, larger radii. Atoms and ions that have the same electron configuration are said to be isoelectronic. Examples of isoelectronic species are N 3, O 2, F, Ne, Na +, Mg 2+, and Al 3+ (1s 2 2s 2 2p 6 ). Another isoelectronic series is P 3, S 2, Cl, Ar, K +, Ca 2+, and Sc 3+ ([Ne]3s 2 3p 6 ). For atoms or ions that are isoelectronic, the number of protons determines the size. The greater the nuclear charge, the smaller the radius in a series of isoelectronic ions and atoms. This content is available for free at
6 Chapter 6 Electronic Structure and Periodic Properties of Elements 329 Variation in Ionization Energies The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy (IE 1 ). The first ionization energy for an element, X, is the energy required to form a cation with +1 charge: X(g) X + (g) + e IE 1 The energy required to remove the second most loosely bound electron is called the second ionization energy (IE 2 ). X + (g) X 2+ (g) + e IE 2 The energy required to remove the third electron is the third ionization energy, and so on. Energy is always required to remove electrons from atoms or ions, so ionization processes are endothermic and IE values are always positive. For larger atoms, the most loosely bound electron is located farther from the nucleus and so is easier to remove. Thus, as size (atomic radius) increases, the ionization energy should decrease. Relating this logic to what we have just learned about radii, we would expect first ionization energies to decrease down a group and to increase across a period. Figure 6.34 graphs the relationship between the first ionization energy and the atomic number of several elements. The values of first ionization energy for the elements are given in Figure Within a period, the IE 1 generally increases with increasing Z. Down a group, the IE 1 value generally decreases with increasing Z. There are some systematic deviations from this trend, however. Note that the ionization energy of boron (atomic number 5) is less than that of beryllium (atomic number 4) even though the nuclear charge of boron is greater by one proton. This can be explained because the energy of the subshells increases as l increases, due to penetration and shielding (as discussed previously in this chapter). Within any one shell, the s electrons are lower in energy than the p electrons. This means that an s electron is harder to remove from an atom than a p electron in the same shell. The electron removed during the ionization of beryllium ([He]2s 2 ) is an s electron, whereas the electron removed during the ionization of boron ([He]2s 2 2p 1 ) is a p electron; this results in a lower first ionization energy for boron, even though its nuclear charge is greater by one proton. Thus, we see a small deviation from the predicted trend occurring each time a new subshell begins.
7 330 Chapter 6 Electronic Structure and Periodic Properties of Elements Figure 6.34 number. The first ionization energy of the elements in the first five periods are plotted against their atomic Figure 6.35 elements. This version of the periodic table shows the first ionization energy of (IE 1), in kj/mol, of selected This content is available for free at
8 Chapter 6 Electronic Structure and Periodic Properties of Elements 331 Another deviation occurs as orbitals become more than one-half filled. The first ionization energy for oxygen is slightly less than that for nitrogen, despite the trend in increasing IE 1 values across a period. Looking at the orbital diagram of oxygen, we can see that removing one electron will eliminate the electron electron repulsion caused by pairing the electrons in the 2p orbital and will result in a half-filled orbital (which is energetically favorable). Analogous changes occur in succeeding periods (note the dip for sulfur after phosphorus in Figure 6.35). Removing an electron from a cation is more difficult than removing an electron from a neutral atom because of the greater electrostatic attraction to the cation. Likewise, removing an electron from a cation with a higher positive charge is more difficult than removing an electron from an ion with a lower charge. Thus, successive ionization energies for one element always increase. As seen in Table 6.3, there is a large increase in the ionization energies (color change) for each element. This jump corresponds to removal of the core electrons, which are harder to remove than the valence electrons. For example, Sc and Ga both have three valence electrons, so the rapid increase in ionization energy occurs after the third ionization. Successive Ionization Energies for Selected Elements (kj/mol) Element IE 1 IE 2 IE 3 IE 4 IE 5 IE 6 IE 7 K Ca Sc Ga Ge Not available Not available As Not available Table 6.3 Example 6.13 Ranking Ionization Energies Predict the order of increasing energy for the following processes: IE 1 for Al, IE 1 for Tl, IE 2 for Na, IE 3 for Al. Solution Removing the 6p 1 electron from Tl is easier than removing the 3p 1 electron from Al because the higher n orbital is farther from the nucleus, so IE 1 (Tl) < IE 1 (Al). Ionizing the third electron from Al (Al 2+ Al 3+ + e ) requires more energy because the cation Al 2+ exerts a stronger pull on the electron than the neutral Al atom, so IE 1 (Al) < IE 3 (Al). The second ionization energy for sodium removes a core electron, which is a much higher energy process than removing valence electrons. Putting this all together, we obtain: IE 1 (Tl) < IE 1 (Al) < IE 3 (Al) < IE 2 (Na). Check Your Learning
9 332 Chapter 6 Electronic Structure and Periodic Properties of Elements Which has the lowest value for IE 1 : O, Po, Pb, or Ba? Answer: Ba Variation in Electron Affinities The electron affinity [EA] is the energy change for the process of adding an electron to a gaseous atom to form an anion (negative ion). X(g) + e X (g) EA 1 This process can be either endothermic or exothermic, depending on the element. The EA of some of the elements is given in Figure You can see that many of these elements have negative values of EA, which means that energy is released when the gaseous atom accepts an electron. However, for some elements, energy is required for the atom to become negatively charged and the value of their EA is positive. Just as with ionization energy, subsequent EA values are associated with forming ions with more charge. The second EA is the energy associated with adding an electron to an anion to form a 2 ion, and so on. As we might predict, it becomes easier to add an electron across a series of atoms as the effective nuclear charge of the atoms increases. We find, as we go from left to right across a period, EAs tend to become more negative. The exceptions found among the elements of group 2 (2A), group 15 (5A), and group 18 (8A) can be understood based on the electronic structure of these groups. The noble gases, group 18 (8A), have a completely filled shell and the incoming electron must be added to a higher n level, which is more difficult to do. Group 2 (2A) has a filled ns subshell, and so the next electron added goes into the higher energy np, so, again, the observed EA value is not as the trend would predict. Finally, group 15 (5A) has a half-filled np subshell and the next electron must be paired with an existing np electron. In all of these cases, the initial relative stability of the electron configuration disrupts the trend in EA. We also might expect the atom at the top of each group to have the largest EA; their first ionization potentials suggest that these atoms have the largest effective nuclear charges. However, as we move down a group, we see that the second element in the group most often has the greatest EA. The reduction of the EA of the first member can be attributed to the small size of the n = 2 shell and the resulting large electron electron repulsions. For example, chlorine, with an EA value of 348 kj/mol, has the highest value of any element in the periodic table. The EA of fluorine is 322 kj/mol. When we add an electron to a fluorine atom to form a fluoride anion (F ), we add an electron to the n = 2 shell. The electron is attracted to the nucleus, but there is also significant repulsion from the other electrons already present in this small valence shell. The chlorine atom has the same electron configuration in the valence shell, but because the entering electron is going into the n = 3 shell, it occupies a considerably larger region of space and the electron electron repulsions are reduced. The entering electron does not experience as much repulsion and the chlorine atom accepts an additional electron more readily. This content is available for free at
10 Chapter 6 Electronic Structure and Periodic Properties of Elements 333 Figure 6.36 This version of the periodic table displays the electron affinity values (in kj/mol) for selected elements. The properties discussed in this section (size of atoms and ions, effective nuclear charge, ionization energies, and electron affinities) are central to understanding chemical reactivity. For example, because fluorine has an energetically favorable EA and a large energy barrier to ionization (IE), it is much easier to form fluorine anions than cations. Metallic properties including conductivity and malleability (the ability to be formed into sheets) depend on having electrons that can be removed easily. Thus, metallic character increases as we move down a group and decreases across a period in the same trend observed for atomic size because it is easier to remove an electron that is farther away from the nucleus.
11 Chapter 6 Electronic Structure and Periodic Properties of Elements 343 Summary 6.5 Periodic Variations in Element Properties Electron configurations allow us to understand many periodic trends. Covalent radius increases as we move down a group because the n level (orbital size) increases. Covalent radius mostly decreases as we move left to right across a period because the effective nuclear charge experienced by the electrons increases, and the electrons are pulled in tighter to the nucleus. Anionic radii are larger than the parent atom, while cationic radii are smaller, because the number of valence electrons has changed while the nuclear charge has remained constant. Ionization energy (the energy associated with forming a cation) decreases down a group and mostly increases across a period because it is easier to remove an electron from a larger, higher energy orbital. Electron affinity (the energy associated with forming an anion) is more favorable (exothermic) when electrons are placed into lower energy orbitals, closer to the nucleus. Therefore, electron affinity becomes increasingly negative as we move left to right across the periodic table and decreases as we move down a group. For both IE and electron affinity data, there are exceptions to the trends when dealing with completely filled or half-filled subshells. Exercises 6.5 Periodic Variations in Element Properties 67. Based on their positions in the periodic table, predict which has the smallest atomic radius: Mg, Sr, Si, Cl, I. 68. Based on their positions in the periodic table, predict which has the largest atomic radius: Li, Rb, N, F, I. 69. Based on their positions in the periodic table, predict which has the largest first ionization energy: Mg, Ba, B, O, Te. 70. Based on their positions in the periodic table, predict which has the smallest first ionization energy: Li, Cs, N, F, I. 71. Based on their positions in the periodic table, rank the following atoms in order of increasing first ionization energy: F, Li, N, Rb 72. Based on their positions in the periodic table, rank the following atoms or compounds in order of increasing first ionization energy: Mg, O, S, Si 73. Atoms of which group in the periodic table have a valence shell electron configuration of ns 2 np 3? 74. Atoms of which group in the periodic table have a valence shell electron configuration of ns 2? 75. Based on their positions in the periodic table, list the following atoms in order of increasing radius: Mg, Ca, Rb, Cs. 76. Based on their positions in the periodic table, list the following atoms in order of increasing radius: Sr, Ca, Si, Cl. 77. Based on their positions in the periodic table, list the following ions in order of increasing radius: K +, Ca 2+, Al 3+, Si List the following ions in order of increasing radius: Li +, Mg 2+, Br, Te 2.
12 344 Chapter 6 Electronic Structure and Periodic Properties of Elements 79. Which atom and/or ion is (are) isoelectronic with Br + : Se 2+, Se, As, Kr, Ga 3+, Cl? 80. Which of the following atoms and ions is (are) isoelectronic with S 2+ : Si 4+, Cl 3+, Ar, As 3+, Si, Al 3+? 81. Compare both the numbers of protons and electrons present in each to rank the following ions in order of increasing radius: As 3, Br, K +, Mg Of the five elements Al, Cl, I, Na, Rb, which has the most exothermic reaction? (E represents an atom.) What name is given to the energy for the reaction? Hint: note the process depicted does not correspond to electron affinity E + (g) + e E(g) 83. Of the five elements Sn, Si, Sb, O, Te, which has the most endothermic reaction? (E represents an atom.) What name is given to the energy for the reaction? E(g) E + (g) + e 84. The ionic radii of the ions S 2, Cl, and K + are 184, 181, 138 pm respectively. Explain why these ions have different sizes even though they contain the same number of electrons. 85. Which main group atom would be expected to have the lowest second ionization energy? 86. Explain why Al is a member of group 13 rather than group 3? This content is available for free at
Worksheet 11 - Periodic Trends A number of physical and chemical properties of elements can be predicted from their position in the Periodic Table. Among these properties are Ionization Energy, Electron
SCPS Chemistry Worksheet Periodicity A. Periodic table 1. Which are metals? Circle your answers: C, Na, F, Cs, Ba, Ni Which metal in the list above has the most metallic character? Explain. Cesium as the
Answer Key ALE 7. Ionization Energies & Electron Affinities (Reference: Sections 8.4 - Silberberg 5 th edition) How can one use the Periodic Table to make predictions of atomic properties? The Model: First
The Periodic Table: Chapter Problems Periodic Table 1. As you move from left to right across the periodic table, how does atomic number change? 2. What element is located in period 3, group 13? 3. What
Name Period Date Honors Chemistry - Periodic Trends Check Your Understanding Answer the following, formulating responses in your own words. (This helps you better understand the concepts) 1. Define shielding
The following are topics and sample questions for the first exam. Topics 1. Mendeleev and the first periodic Table 2. Information in the Periodic Table a. Groups (families) i. Alkali (group 1) ii. Alkaline
Chemistry 101 ANSWER KEY REVIEW QUESTIONS Chapter 8 Use only a periodic table to answer the following questions. 1. Write complete electron configuration for each of the following elements: a) Aluminum
Chapter 8: The Periodic Table 8.1: Development of the Periodic Table Johann Dobereiner: - first to discover a pattern of a group of elements like Cl, Br, and I (called triads). John Newland: - suggested
Shielding effect Effective nuclear charge, Z eff, experienced by an electron is less than the actual nuclear charge, Z Electrons in the outermost shell are repelled (shielded) by electrons in the inner
TRENDS OF CHEMICAL AND PHYSICAL PROPERTIES IN PERIODIC TABLE Sixth Course (General Chemistry) by Dr. Istadi 1 Trends in Atomic Size All physical and chemical behavior of the elements is based ultimately
Chapter 7 Periodic Properties of the Elements 1. Elements in the modern version of the periodic table are arranged in order of increasing. (a). oxidation number (b). atomic mass (c). average atomic mass
PSI AP Chemistry Periodic Trends MC Review Name Periodic Law and the Quantum Model Use the PES spectrum of Phosphorus below to answer questions 1-3. 1. Which peak corresponds to the 1s orbital? (A) 1.06
PSI AP Chemistry Unit 2: Free Response CW/HW Name The Periodic Law and Ionic Charge Classwork: 1. The PES spectrum for an element can be found below: Intensity 0.63 0.77 3.24 5.44 39.2 48.5 433 Binding
Name: Class: Date: ID: A Unit 7 Review Multiple Choice Identify the choice that best completes the statement or answers the question. 1) In which set of elements would all members be expected to have very
UNIT-3 Classification of elements and periodicity in properties One mark questions:. For the triad of elements A, B and C if the atomic weights of A and C are 7 and 39. Predict the atomic weight of B..
Use the PES spectrum of Phosphorus below to answer questions 1-3. 1. Which peak corresponds to the 1s orbital? (A) 1.06 (B) 1.95 (C) 13.5 (D) 18.7 (E) 208 2. Which peak corresponds to the valence 3p orbital?
Chemistry Chapter 14 Review Name answer key General Concept Questions 1) is credited with developing the concept of atomic numbers. A) Dmitri Mendeleev B) Lothar Meyer C) Henry Moseley D) Ernest Rutherford
Electron Configuration and Chemical Periodicity Orbital approximation: the e density of an isolated many-electron atom is approximated by the sum of the e densities of each of the individual e taken separately.
DEVELOPMENT OF THE PERIODIC TABLE Prior to the 1700s, relatively few element were known, and consisted mostly of metals used for coinage, jewelry and weapons. From early 1700s to mid-1800s, chemists discovered
Chapter 3 Atomic Structure and Properties Introduction The nuclear atom and quantum theory are the accepted theories for the atom. In this chapter, we demonstrate their utility by using them to explain
Chapter 6. Periodic Relationships Among the Elements Student: 1. The nineteenth century chemists arranged elements in the periodic table according to increasing A. atomic number. B. number of electrons.
Key Concepts: Periodic Properties of the Elements 1. Understand and be able to predict and explain trends in effective nuclear charge, Z eff. 2. Understand and be able to predict and explain the periodic
The Periodic Table Organizing the Elements A few elements, such as gold and copper, have been known for thousands of years - since ancient times Yet, only about 13 had been identified by the year 1700.
5.4 Trends in the Periodic Table Think about all the things that change over time or in a predictable way. For example, the size of the computer has continually decreased over time. You may become more
Periodic Table Questions 1. The elements characterized as nonmetals are located in the periodic table at the (1) far left; (2) bottom; (3) center; (4) top right. 2. An element that is a liquid at STP is
Coincidence? I Think Not! 0 As you have realized, the Periodic Table provides a great deal more information than just atomic number and atomic mass! 0 Each period (row) corresponds to an energy level 0
1. Which statement best describes Group 2 elements as they are considered in order from top to bottom of the Periodic Table? (A) The number of principal energy levels increases, and the number of valence
Chapter 6 The Periodic Table Organizing the Periodic Table In a grocery store, the products are grouped according to similar characteristics. With a logical classification system, finding and comparing
Slide 1 of 39 Chapter One (continued) Many Electron Atoms and The Periodic Table Slide 2 of 39 Multielectron Atoms In the hydrogen atom, all subshells of a principal shell are at the same energy level.
CHEMISTRY TEST: THE PERIODIC TABLE Directions: Multiple Choice For each of the following questions, choose the answer that best answers the question and place it on your answer sheet. 1. Which of the following
Concept/Skills Development [Wolfgram, D. (1987, December). ChemMatters, p. 16.] FAMILY RESEMBLANCE Can you place these elements in their proper positions on the Periodic Table without knowing their identity?
Chemistry Unit 5- Periodic Table and Periodic Law Name: History of the Periodic Table I. Mendeleev and Chemical Periodicity A. Wanted to organize elements according to their B. When elements were arranged
Ionic bonds and main group chemistry Learning objectives Write Lewis dot structures of atoms and ions Describe physical basis underlying octet rule Predict ionic charges using periodic table Define lattice
Chapter 8 Basic Concepts of the Chemical Bonding 1. There are paired and unpaired electrons in the Lewis symbol for a phosphorus atom. (a). 4, 2 (b). 2, 4 (c). 4, 3 (d). 2, 3 Explanation: Read the question
CHAPTER 1 1 Electrons and Chemical Bonding SECTION Chemical Bonding BEFORE YOU READ After you read this section, you should be able to answer these questions: What is chemical bonding? What are valence
1. The elements on the Periodic Table are arranged in order of increasing A) atomic mass B) atomic number C) molar mass D) oxidation number 2. Which list of elements consists of a metal, a metalloid, and
Chapter 5: The Periodic Law Section 5.1: The History of the Periodic Table Dmitri Mendeleev (1869) first person to organize the elements in a chart Organized about 70 elements by increasing atomic mass
The Periodic Table of The Elements Elements are like a collection As more and more elements were discovered it became more important to organize and classify them Between the late 1700 s and mid 1800 s
Chapter 7 Electron Configurations and the Properties of Atoms In this Chapter In the last chapter we introduced and explored the concept of orbitals, which define the shapes electrons take around the nucleus
Periodic Table Extra Practice 1. Which of the following elements in Period 3 has the greatest metallic character? 1) Ar 3) Mg 2) Si 4) S 2. Which sequence of atomic numbers represents elements which have
IONISATION ENERGY IONISATION ENERGY CONTENTS What is Ionisation Energy? Definition of t Ionisation Energy What affects Ionisation Energy? General variation across periods Variation down groups Variation
HEINS10-118-128v4.qxd 12/30/06 2:05 PM Page 118 CHAPTER 10 MODERN ATOMIC THEORY AND THE PERIODIC TABLE SOLUTIONS TO REVIEW QUESTIONS 1. An electron orbital is a region in space around the nucleus of an
Some Basic Concepts & Classification of Elements Test No. 15 Time: 2 Hrs. Total Marks: 311 Date: 07/07/2013 Name Batch: CLASSIFICATION OF ELEMENTS Section A (Objective Question only one correct answer)
Question (1) Chapter Two Periodic Table Choose the correct answer for each statement of the following:- 1- The elements of same vertical group are identical in the number of----------- a valence electrons
Unit 2 Periodic Behavior and Ionic Bonding 6.1 Organizing the Elements I. The Periodic Law A. The physical and chemical properties of the elements are periodic functions of their atomic numbers B. Elements
Name: Wednesday, January 16, 2008 Test 7: Periodic Table Review Questions 1. Which halogen is a solid at STP? 1. fluorine 3. bromine 2. chlorine 4. iodine 2. Element M is a metal and its chloride has the
Click: http://www.pbs.org/wgbh/nova/physics/hunting-elements.html Periodic Table: Organizes and classifying the elements Dmitri Mendeleev: Russian chemist who arranged according to their increasing atomic.
Building up the atoms in the periodic table 1) The Aufbau ( building up ) principle: lowest energy orbitals are filled first 1s, then 2s, then 2p, then 3s, then 3p, etc. 2) Remember the Pauli exclusion
Composition and Structure of the Atom Atom: basic unit of an element; smallest unit that retains chemical properties of an element Subatomic particles: Small particles that are the building blocks from
Mendeleev s Periodic Table of the Elements Dmitri Mendeleev born 1834 in the Soviet Union. In 1869 he organised the 63 known elements into a periodic table based on atomic masses. He predicted the existence
Electron Configuration and Chemical Periodicity Chapter 8 8.1 Development of the Periodic Table 8.2 Characteristics of Many-Electron Atoms Electron Configuration and Chemical Periodicity 8.3 The Quantum-Mechanical
Matter and the Periodic Table Purpose The purpose of this station is to reinforce students understanding of the organization and predictive power of the Periodic Table of the Elements and students ability
The Octet Rule Our discussion of valence electron configurations leads us to one of the cardinal tenets of chemical bonding, the octet rule. The octet rule states that atoms become especially stable when
Ch 8 Atomic Electron Configuration and Chemical Periodicity Pauli Exclusion principle No two electrons in an atom have the same set of quantum numbers. If two electrons occupy the same orbital, they must
179 ELECTRON CONFIGURATION (SHORT FORM) - We can represent the electron configuration without drawing a diagram or writing down pages of quantum numbers every time. We write the "electron configuration".
CHAPTER EIGHT BNDING: GENERAL CNCEPT or Review 1. Electronegativity is the ability of an atom in a molecule to attract electrons to itself. Electronegativity is a bonding term. Electron affinity is the
Chemistry: The Periodic Table and Periodicity Name: per: Date:. 1. By what property did Mendeleev arrange the elements? 2. By what property did Moseley suggest that the periodic table be arranged? 3. What
IONISATION ENERGY IONISATION ENERGY CONTENTS What is Ionisation Energy? Definition of t Ionisation Energy What affects Ionisation Energy? General variation across periods Variation down groups Variation
TRENDS IN ATOMIC PROPERTIES: THE PERIODIC TABLE Electron configurations determine organization of the periodic table Next properties of elements and their periodic behavior Elemental properties determined
Chemistry, The Central Science, 11th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 7 John D. Bookstaver St. Charles Community College Cottleville, MO Development of Table
Bonds hapter 8 Bonding: General oncepts Forces that hold groups of atoms together and make them function as a unit. Bond Energy Bond Length It is the energy required to break a bond. The distance where
The Periodic Table; Chapter 5: Section 1 - History of the Periodic Table Objectives: Explain the roles of Mendeleev and Moseley in the development of the periodic table. Describe the modern periodic table.
Periodicity and the Periodic Table the result Dmitri Mendeleev arranged elements in order of their atomic numbers, such that elements with similar properties fell into the same column or group. The Periodic
6 THE PERIODIC TABLE SECTION 6.1 ORGANIZING THE ELEMENTS (pages 155 160) This section describes the development of the periodic table and explains the periodic law. It also describes the classification
M. Prakash Academy Weekly workout 6 Periodic properties Q1. According to modern periodic law the properties of elements repeat at regular intervals when the elements are arranged in order of: (a) decreasing
Chemistry: The Periodic Table and Periodicity Name: Hour: Date: Directions: Answer each of the following questions. You need not use complete sentences. 1. Who first published the classification of the
Chapter 3. Elements, Atoms, Ions, and the Periodic Table The Periodic Law and the Periodic Table In the early 1800's many elements had been discovered and found to have different properties. In 1817 Döbreiner's
2.13 Ionisation Energies Definition :First ionisation energy The first ionisation energy is the energy required when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge
Chapter 6: The Periodic Table Study Guide I. General organization of table A. Modern periodic table 1. Increasing atomic number B. 3 major blocks 1. Metals a. Mostly solids at room temperature b. Conduct
3.06 Periodic Table and Periodic Trends Dr. Fred Omega Garces Chemistry 100, Miramar College 1 3.06 Periodic Table and Periodic Trend The Periodic Table and the Elements What is the periodic table? What
Chapter 5 TEST: The Periodic Table name HPS # date: Multiple Choice Identify the choice that best completes the statement or answers the question. 1. The order of elements in the periodic table is based
Chapter 11 Chemical Bonds: The Formation of Compounds from Atoms 1 11.1 Periodic Trends in atomic properties 11.1 Periodic Trends in atomic properties design of periodic table is based on observing properties
11 Chemical Bonds The Formation of Compounds from Atoms Chapter Outline 11.1 11.2 Lewis Structures of Atoms 11.3 The Ionic Bond Transfer of Electrons from One Atom to Another 11.4 Predicting Formulas of
Unit 5 Elements and their Properties 1. In 1871, Russian chemist Mendeleev created the forerunner of the modern periodic table. 2. The elements in Mendeleev's table were arranged in order of increasing
Chemistry Periodic Table Name: Period: 1 2 3 4 5 6 7 8 Periodic Table Study Guide Directions: Please use this packet as practice and review. DO NOT try to answer these questions during presentations, take
Chapter 2 1 Chapter 2 Atoms and the Periodic Table Solutions to In-Chapter Problems 2.1 Each element is identified by a one- or two-letter symbol. Use the periodic table to find the symbol for each element.
Unit 3.2: The Periodic Table and Periodic Trends Notes The Organization of the Periodic Table Dmitri Mendeleev was the first to organize the elements by their periodic properties. In 1871 he arranged the
Name Class Date The Periodic Table ELECTRONS AND THE STRUCTURE OF ATOMS 6.1 Organizing the Elements Essential Understanding Although Dmitri Mendeleev is often credited as the father of the periodic table,
Trends of the Periodic Table Diary Trends are patterns of behaviors that atoms on the periodic table of elements follow. Trends hold true most of the time, but there are exceptions, or blips, where the