Chemistry
Periodic Trends
Periodic trends refer to the variations in properties of elements as one moves across a row or down a column of the periodic table. These trends include atomic radius, ionization energy, electronegativity, and electron affinity. They provide valuable insights into the behavior and reactivity of elements, helping to predict their chemical and physical properties.
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11 Key excerpts on "Periodic Trends"
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
As described by the Periodic Law, there are trends in many properties of the elements in the periodic table which increase or decrease systematically as you move across a period or down a group. These Periodic Trends can be explained by the changes in electronic configuration of the elements within the periods or groups. Some important Periodic Trends are: atomic radius, ionization energy, electron affinity, and electronegativity. These observed trends give chemists a useful tool that allows them to predict an element's properties, chemical reactivities, and types of compounds that it will form from their position in the periodic table.Perhaps the first periodic trend to be observed is that of changing atomic radius . The atomic radius is the distance from the nucleus to the outermost occupied electron orbital in an atom. As shown in Fig. 2.14 , the atomic radii follow a fairly regular pattern of a minimum followed immediately by a maximum and then decreasing steadily to the next minimum as atomic number is increased. The maximum radii are found for the alkali metals in group 1; lithium (atomic number 3), sodium (atomic number 11), potassium (atomic number 19), rubidium (atomic number 37), cesium (atomic number 55), and francium (atomic number 87). These maxima are followed by a steady decrease in atomic radius with increasing atomic number until reaching a minimum. The minimum radii are for the noble gases in group 18; helium (atomic number 2), neon (atomic number 10), argon (atomic number 18), krypton (atomic number 36), xenon (atomic number 54), and radon (atomic number 86). So, the atomic radius decreases from left to right across a period.Fig. 2.14 Atomic radius in picometers as a function of atomic number for elements 1–95.The reason for this decreasing trend in atomic radii occurs because as atomic number increases proceeding from left to right across a period, the electrons are added to the same electron shell. This has little effect on increasing the atomic radius. But, as the electrons are added to the shell, protons are also added to the nucleus. Increasing the number of positively charged protons in the nucleus causes the attraction between the valence electrons and the nucleus to become stronger, pulling the valence shell closer to the nucleus. So, the atomic radius decreases across a period as protons are added to the nucleus.
- Martha Mackin(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
FIVE The periodic table OVERVIEW This chapter gives the history of the development of the periodic table. The properties of the elements are repeated in a regular way, so the elements can be arranged into a periodic table. The horizontal rows in the table are called periods; the vertical columns are called groups. Chapter 4 described how the electronic structure of atoms varies in a regular way with atomic number. This chapter shows the relationship between the periodic repetition of the properties of the elements and the periodic repetition of electronic structure. Each group in the periodic table has similar properties and similar arrangement of outer electrons. Each period is related to the filling of one or more energy levels. Elements may be identified as metals, nonmetals, semimetals, transition elements, inner tran-sition elements, or noble gases by their position in the periodic table. The properties of the elements within a group are similar but vary from each other in a regular way. The properties of elements in a period are different but each period shows similar changes in properties of the elements from left to right in the periodic table. These periodic varia-tions in properties are related to periodic changes in the structure of the atoms. Atomic size increases down a group and decreases from left to right across a period; Ionization potential and electron affinity decrease down a group and increase across a period. Metallic properties are related to low ionization potential, low electron affinity and large atomic size. Nonmetallic properties are related to high ionization potential, high electron affinity and small atomic size. Other periodic arrangements of the elements are described and illustrated. A brief survey of elements by periodic group is included.
- Barron's Educational Series, Albert S. Tarendash(Authors)
- 2021(Publication Date)
- Barrons Educational Services(Publisher)
Pages 256–258 )9.6Concept: For Groups 1, 2, and 13–18 in the Periodic Table, elements within the same group have the same number of valence electrons (helium is an exception) and, therefore, have similar chemical properties. (Page 244 )Skill: Given the chemical formula of a compound (such as XCl or XCl2 ), determine the group of an element (X). (Pages 256–258 )9.7Concept: The succession of elements within the same group demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, and metallic (or nonmetallic) properties. (Pages 248–255 )Skill: Compare and contrast the properties of elements within a single group (Groups 1, 2, 13–18) in the Periodic Table. (Pages 248–255 )9.8Concept: The succession of elements across the same period demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, and metallic (or nonmetallic) properties. (Pages 248–255 )Skill: Compare and contrast the properties of elements across a period (for Groups 1, 2, 13–18) in the Periodic Table. (Pages 249–258 )9.9Concept: When an atom gains one or more electrons, its radius increases. When an atom loses one or more electrons, its radius decreases. (Page 247 )
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9.1 INTRODUCTION
The first question we need to ask is this: Why do we have a Periodic Table of the Elements? The term periodic means that a quantity repeats itself at regular intervals. For example, the motion of the Moon around Earth is periodic because, after a certain length of time (1 month), the motion repeats itself. As early as the Middle Ages, scientists recognized that elements could be differentiated by their physical and chemical properties. These properties are also periodic—if we list the elements properly.
eBook - PDFChemistry for Today
General, Organic, and Biochemistry
- Spencer Seager, Michael Slabaugh, Maren Hansen, , Spencer Seager, Spencer Seager, Michael Slabaugh, Maren Hansen(Authors)
- 2021(Publication Date)
- Cengage Learning EMEA(Publisher)
The horizontal rows in the table are called periods and are numbered from top to bottom. Thus, each element belongs to both a period and a group of the periodic table. periodic law A statement about the behavior of the elements when they are arranged in a specific order. In its present form, it is stated as follows: Elements with similar chemical properties occur at regular (periodic) intervals when the elements are arranged in order of increasing atomic numbers. group or family of the periodic table A vertical column of elements that have similar chemical properties. period of the periodic table A horizontal row of elements. Copyright 2022 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 72 Chapter 3 Example 3.1 Groups and Periods of the Periodic Table Phosphorus (P) is an essential element for the formation of bones and teeth. Chromium (Cr) is also an essential element which helps the body to utilize glucose. Identify the group and period to which each of the following belongs: a. P b. Cr c. element number 30 d. element number 53 Solution a. Phosphorus (P) is present in group VA(15) and period 3. b. Chromium (Cr) is in group VIB(6) and period 4. c. The element with atomic number 30 is zinc (Zn), which is found in group IIB(12) and period 4. d. Element number 53 is iodine (I), found in group VIIA(17) and period 5. ✔ LEARNING CHECK 3.1 Write the symbol for the element found in the following places of the periodic table: a.
- R.A. Mackay, W. Henderson(Authors)
- 2017(Publication Date)
- CRC Press(Publisher)
Table. 2.14 that these increase towards the right of the Periods and decrease down the Groups. In addition, they reflect the other, smaller, variations which have been remarked; for example, the changes in the Main Groups are more pronounced than in the Transition Groups, and the discontinuity in properties of the elements from gallium to bromine when compared with the rest of their respective Groups is reflected in their electronegativity values.8.5 Chemical behaviour and periodic position
The detailed chemistry of the elements is discussed in the succeeding chapters. In this section, the skeleton of the periodic properties is outlined to provide a framework for the more detailed account which follows.Those elements where the outermost electrons are in a new quantum level, after a rare gas configuration, normally react by losing these loosely bound electrons and forming cations. This mode of behaviour is typical of the elements of the lithium, beryllium and scandium Groups together with the lanthanide elements, which have the respective valency shell configurations, s1 , s2 and d1 s2 . All these elements, with the exception of beryllium itself, lose these outer electrons completely with the formation of cations; M+ in the lithium Group, M2+ in the beryllium Group, and M3+ for scandium, yttrium and the lanthanides.The elements of the boron, carbon, nitrogen, oxygen, and fluorine Groups, where the outermost electrons are in p orbitals, show more complicated behaviour.(a) Elements with electron configurations close to the rare gases can acquire electrons to form anions with complete rare gas shells. Thus the elements of the halogen Group all form X– ions, and we also find stable compounds containing O2− , S2− and N3−
eBook - PDFEngineering Chemistry
Fundamentals and Applications
- Shikha Agarwal(Author)
- 2019(Publication Date)
- Cambridge University Press(Publisher)
8.1 Introduction The genesis of classification of elements dates back to early 1800s when the German chemist Johann Dobereiner made the first attempt to systematise the study of properties of elements and gave the Dobereiners law of triads. This was followed by the Newlands law of octaves. The first detailed classification of elements was proposed by Russian chemist Dmitri Mendeleev (1837–1907) and German chemist Lothar Meyer (1830–1895). Both these scientists worked independently and in 1869 proposed that when elements are arranged in the increasing order of their atomic weights, similarities appear in their physical and chemical properties at regular intervals. However, the Mendeleev’s periodic table had certain anomalies, which were addressed in due course of time. Gradual improvement saw the development of the modern periodic law (given by Henry Moseley) in which the elements are arranged in the order of increasing atomic numbers. The elements are arranged in the periodic table in the order of increasing atomic numbers. They are divided into 18 vertical columns called groups and seven horizontal rows termed as periods. According to the IUPAC recommendations, the groups are numbered from 1 to 18, replacing the older notations of groups IA…VIIA, IB…VIIB and zero. The seven periods have 2, 8, 8, 18, 18 and 32 elements, respectively. The seventh period is incomplete and will theoretically consist of 32 elements. The properties of the elements vary periodically due to periodic variation in their electronic configuration. In this chapter, we will study the variations in electronic configurations, atomic and ionic sizes, ionisation energies and other related properties. 8.2 Basic Concepts Before studying the periodicity in properties of elements, it is important to have knowledge of some basic terms and concepts.
eBook - PDF- David A. Ucko(Author)
- 2013(Publication Date)
- Academic Press(Publisher)
Elements of the same type are grouped together in the periodic table. Wit HltS' cause of this special property, silicon is used to make the microelec-tronic c h i p s used in calculators, digital watches, and many other devices. Figure 5-6 summarizes the locations in the periodic table of the different categories of elements. 5.5 The properties of the elements show certain trends that correspond to PERIODIC two progressive changes indicated in the periodic table: OPERTIES ^ ^ . , , 1 The steady increase in atomic number by one, from each element to the next across a period. 2 The repeated filling of outermost electron levels, from one period to the next. Consider the sizes of atoms of the elements shown in Figure 5-7. Notice that as we move across each period from left to right, the atomic size decreases. This decrease in atomic size results from change 1 in the list. As the atomic number increases, the positive charge of the nucleus increases (because protons are positively charged). This in-crease in charge causes electrons to be attracted more strongly to the nucleus. As a result, the atom shrinks. A second trend, resulting from change 2, is the increase in atomic size down each group. In this case the atoms become larger because electrons are being added to levels that are farther away from the nucleus. These two trends, the decrease in atomic size across a period and the increase in atomic size down a group, are largest for the main group elements, smaller for the transition elements, and smallest for the inner transition elements. For example, the difference in atomic 137 —ir-i i Radii decrease across periods 2A FIGURE 5-7 Sizes of atoms. Notice that the size increases down a group and decreases across a period (from left to right).
eBook - PDF- C N R Rao(Author)
- 2009(Publication Date)
- World Scientific(Publisher)
• have similar properties. • become more metallic in character down the group. In water solution, most ions of transition elements are coloured. Some of these elements exhibit more than one valency (Fe 2+ , Fe 3+ ; Mn 2+ , Mn 3+ , Mn 4+ ,…, Mn 7+ ). Some of the metals and their ions have catalytic properties. Iron, cobalt and nickel are magnetic. These are permanent magnets ( ferromagnetic ). Compounds of these and many other transition elements are attracted by a magnet. They are paramagnetic . Sc scandium Ti titanium V vanadium Cr chromium Mn manganese Fe iron Co cobalt Ni nickel Cu copper Zn zinc Properties of an element in a group can be predicted on the basis of the properties of another element in the same group. Elements and the periodic table 119 Periodicity in the modern periodic table is a function of the electronic configuration. 2.5 Periodic table and properties of elements Physical properties: • Melting point, boiling point and density of elements increase across a period until maximum values are reached. Then they decrease. Noble gases have low values. • Elements become less metallic across a period and more metallic down a group. • The atomic size and ionic size decrease across a period and increase down a group. Physical properties Electronegative and electropositive nature Redox properties Properties of compounds There is PERIODICITY of 120 Understanding Chemistry Electron affinities of a few elements (in electron volts) are given below: F 3.40 I 3.06 Cl 3.61 H 0.75 Br 3.36 O 1.46 H 13.6 C 11.3 Na 5.1 He 24.6 N 14.5 Mg 7.6 Li 5.4 O 13.6 Cl 13.0 Be 9.3 F 17.4 K 4.3 B 8.3 Ne 21.6 Rb 4.2 Electron affinity is the energy change that occurs when an atom accepts an electron. Atoms with high electron affinity readily become negative ions.
eBook - PDF- Morris Hein, Susan Arena, Cary Willard(Authors)
- 2016(Publication Date)
- Wiley(Publisher)
In Period 3, electrons are found in levels 1, 2, 3, and so on. Elements that behave in a similar manner are found in groups or families. These form the vertical columns on the periodic table. Two systems exist for numbering the groups. In one system, the columns are numbered from left to right using the numbers 1–18. The A groups are known as the representative elements. The B groups are called the transition elements. In this book we will focus on the representative elements. The groups (columns) of the periodic table often have family names. For example, the group on the far right side of the periodic table (He, Ne, Ar, Kr, Xe, and Rn) is called the noble gases. Group 1A(1) is called the alkali metals, Group 2A(2) the alkaline earth metals, and Group 7A(17) the halogens. How is the structure of the periodic table related to the atomic structures of the elements? We’ve just seen that the periods of the periodic table are associated with the energy level of the outermost electrons of the atoms in that period. Look at the valence electron configurations of the elements we have just examined (FIGURE 10.15). Do you see a pattern? The valence electron configuration for the elements in each column is the same. The chemical behavior and properties of elements in a particular family must therefore be associated with the electron configuration of the elements. The number for the principal energy level is different. This is expected since each new period is associ- ated with a different energy level for the valence electrons. The electron configurations for elements beyond these first 18 become long and tedious to write. We often abbreviate the electron configuration using the following notation: Na [Ne]3s 1 Look carefully at Figure 10.15 and you will see that the p orbitals are full at the noble gases.
eBook - PDF- Brian W. Pfennig(Author)
- 2021(Publication Date)
- Wiley(Publisher)
In the modern era, it is almost impossible to believe that there was ever a time when the periodic nature of the elements was not yet fully understood. The first primitive form of the periodic table was developed in 1829 by Johann Döbereiner, who noticed a repetitive pattern in the chemical properties of some of the elements, lumping them together in groups of three which he called triads. The elements chlorine, bromine, and iodine, for instance, formed a triad because they all reacted in a similar way with the element sodium. In 1865, the Englishman John Newlands recognized that if the elements were listed in order of their increasing atomic weights, they could be arranged into what he called the “law of octaves” (or groups of eight) according to their chemical properties. However, it was not until 1869, when the Russian schoolteacher Dmitri Mendeleev gave a presentation to the Russian Chemical Society entitled “A Dependence Principles of Inorganic Chemistry, Second Edition. Brian W. Pfennig. © 2022 John Wiley & Sons, Inc. Published 2022 by John Wiley & Sons, Inc. Companion website: www.wiley.com/go/pfennig/inorgchem2 between the Properties of the Atomic Weights of the Elements” that the idea of Periodic Trends among the various elements really began to take hold. Mendeleev had written out the physical properties of all the known elements on cards and arranged them into groups based on their trends in a popular card game at the time called “Patience.” Mendeleev’s blueprint for the periodicity of the elements ultimately became the standard because he left gaps in the table for elements that had not yet been discovered. When these elements were eventually found, their properties fit neatly into place, cementing the periodic table as one of the most valuable predictive tools in the history of chemistry.
eBook - PDF- Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
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, resulting in a more negative EA. FIGURE 3.35 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. 3.6 The Periodic Table LEARNING OBJECTIVES By the end of this section, you will be able to: • State the periodic law and explain the organization of elements in the periodic table • Predict the general properties of elements based on their location within the periodic table • Identify metals, nonmetals, and metalloids by their properties and/or location on the periodic table As early chemists worked to purify ores and discovered more elements, they realized that various elements could be grouped together by their similar chemical behaviors. One such grouping includes lithium (Li), sodium (Na), and potassium (K): These elements all are shiny, conduct heat and electricity well, and have similar chemical properties.
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