Periodic Trends
Introduction
Electrons are held in an atom or ion by the
electrostatic attraction between the positively charged nucleus and the
negatively charged electrons. In multi-electron species, the electrons do not
experience the full positive charge of the nucleus due to shielding by
electrons which lie between the electron of interest and the nucleus. The
amount of positive charge that actually acts on an electron is called the effective nuclear charge.
Effective Nuclear Charge
The concept of effective nuclear charge (Z*) is
important to understanding periodic properties. The effective nuclear charge is
that portion of the total nuclear charge that a given electron in an atom
experiences. This is equal to the atomic number (Z) minus the amount (σ) that
other electrons in the atom shield the given atom from the nucleus.
Z* = Z-σ
Example: Lithium has three protons and an electron configuration of 1s22s1.
The electron in the 2s orbital is shielded from the full attraction of the
protons by the electrons of the 1s orbital (Figure 1). Thus, Z* felt by the 2s
electron should be one rather than three. However, lithium's 2s electron does
not behave as if it is experiencing exactly a +1 charge (Z* is actually about
1.3 charge units). This can be explained by the fact that the 2s orbital has
two maxima in its radial probability function (Figure 1), and the lesser maxima
penetrates within the maximum of the inner 1s electron. Although lithium's 2s
electron spends most of its time in the outer lobe of that orbital feeling a
nuclear charge of +1, some of the time it is inside the 1s orbital experiencing
the full nuclear charge of +3. Thus, Z* is somewhat greater than +1.
Slater's Rules for Determining σ
In 1930, J.S. Slater formulated the following set of
empirical rules for determining the values of the shielding constant σ.
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Slater's
Rules
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Example
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Example: Calculate Z* for a 4s and a 3d electron in
Zn
Determine the electron configuration for Zn
(1s)2(2s, 2p)8(3s, 3p)8(3d)10(4s)2
For a 4s electron:
Establish the screening constant for the 4s electron σ = (1 x 0.35) + (18 x 0.85) + (10x1.00) = 25.65
Calculate the effective nuclear charge
Z*= Z-σ = 30-25.65 = 4.35
Calculate the effective nuclear charge
From this example, you can see that
the 3d electrons experience a much greater positive charge than the 4s
electron and would be held more tightly. Thus, the 4s electrons will be the
first removed when Zn is ionized.
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PROBLEMS
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You can access a spreadsheet for calculating
effective nuclear charges here. If you need
assistance in using Excel for plotting data, try this tutorial.
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The Periodic Table
You have used the periodic table throughout your study
of chemistry. Read more about the periodic table
here. Mendeleev was one of the early chemists to recognize
that the properties of the elements were periodic in nature. Read from
Mendeleev's original publication. To get a flavor of what it would
be like to derive the periodic table, try this simulation. It is like
discovering the pieces of a jigsaw puzzle and putting them together.
Periodic Trends
In the remainder of this module, you will be analyzing
the periodic trends that exist among the elements. Start your investigation by
viewing this movie on periodic trends. To view this movie, you may need
to download and install the Real Player.
Atomic Radius
There are several ways to define the atomic radius of
an atom:
- covalent radius (rcov), the half-distance between the nuclei of two atoms joined in a covalent bond
- van der Waals radius (rvdw), the half-distance between the nuclei of two atoms of neighboring molecules
- metallic radius (rmet), for metallic elements, the half-distance between the nuclei of two neighboring atoms in the solid metal
The difference between covalent radius and van der
Waals radius is shown in Figure 2.
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PROBLEMS
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Ionization Energy
Recall that the ionization energy (actually the first
ionization energy) is the energy required to remove an electron from the
outermost occupied orbital of a gaseous atom.
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PROBLEMS
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Electron Affinity
Electron affinity is the change in energy that occurs
when an electron is added to a neutral, gaseous atom.
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PROBLEMS
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Plot the electron affinity for Main Group elements ( Group 1A-8A) versus
atomic number.Display your plot as both a normal graph and as a bar graph. Explain the
general trends that occur across a period and down a family using the concept
of effective nuclear charge. An anomaly to the general trend occurs in a
period in going from the Group 1A element to the Group 2A element and in
going from the Group 4A to Group 5A. An anomaly can also be found between
rows 2 and 3 in going down a period. Give an explanation for each of these
anomalies.
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Electronegativity
Electronegativity is the ability of an atom in a molecule to draw electrons to itself.
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PROBLEMS
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Sketch a periodic table, indicating the trend of electronegativities from
lowest to highest. Relate these trends to effective nuclear charge and atomic
size.
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Other Trends
There are many properties for which you can
investigate periodic trends. There is a nice Excel spreadsheet that allows you
to look at the trends for various properties in 3-D. Open the spreadsheet and
view the plots.
If you have difficulty viewing the spreadsheet in
conjunction with this webpage, try downloading it (right clicking on the link will allow you to save
the file). This spreadsheet uses macros for accessing the plots. If you get a
message that the macros have not downloaded because of the security setting ,
reset the security setting for Excel using the tools menu
A. DEVELOPMENT OF PERIODIC SYSTEM
Enterprises grouping based on common elements of nature done for these elements are easy to learn.
1. Triade Dobereiner
In 1829, Johan Wolfgang Dobereiner study the properties of some of the elements are already known at the time. Dobereiner see any resemblance between some elements of nature and the elements are grouped according to similarities in character. Apparently each group consists of three elements of the so-called triade. If the elements in the triade compiled and sorted based on the similarity of its atomic mass, then the second element is an average of the properties and atomic masses of the first and third elements.
2. Octave Theory Newland
In 1864, John Alexander Reina Newland compiled a list of more numerous elements. Newland arrangement shows that if the elements arranged by atomic mass increases, then the first element has similar properties to the eighth element, the second element is similar to the elements of nature ninth, and so on. The discovery was declared a law Newland Newland Octave.
At Newland Octave compiled list, the elements of noble gases (He, Ne, Ar, Kr, Xe, and Rn) has not been found. Noble gases are found by Rayleigh and Ramsay in 1894. Elements of the noble gases argon gas was first discovered. Newland Octave law only applies to elements with low atomic mass.
3. Mendeleev Periodic System
In 1869, began the periodic table arranged. The periodic table is the work of two scientists, Dmitri Ivanovich Mendeleev from Russia and Julius Lothar Meyer from Germany. They work independently and produce a similar table at the same time. Mendeleev presents its results to the Russian Chemical Society in early 1869, and a new periodic table Meyer appeared in December 1869.
Mendeleev who first suggested the periodic table, it is considered as the inventor of the periodic table is often referred to as short-periodic system of elements. Mendeleev developed the periodic table by atomic mass increases and similar properties. Mendeleev periodic system was first published in the scientific journal Annalen der Chemie in 1871.
The important thing there is in the Mendeleev periodic system are as follows:
a. two adjacent elements, has a relative atomic mass difference of at least two or one unit;
b. there are empty boxes for undiscovered elements, such as 44, 68, 72, and 100;
c. can predict the properties of elements that have not been known as ekasilikon;
d. measurement error can be corrected relative atomic mass of some elements, such as Cr = 52.0 not 43.3.
a. Excess Mendeleev periodic system
1) Chemical and physical properties in a single class of similar elements and changed regularly.
2) the highest valence of an element equal to their group number.
3) to predict the properties of elements yet to be discovered at the time and has had a vacant place.
b. Mendeleev periodic system deficiencies
1) The length of the period is not the same and why not described.
2) Some elements are not ordered by their atomic mass increases, for example: Te (128) before I (127).
3) the mass difference of successive elements are not always 2, but the range between 1 and 4 so difficult to predict the mass of an element that has not been known.
4) Valence more than one element of the unpredictable group.
5) anomaly (deviation) elemental hydrogen from other elements not described.
4. Modern Periodic System
In 1914, Henry G. J. Moseley found that the order of the elements in the periodic table according to the atomic number of the element increases. Moseley managed to find errors in the Mendeleev periodic table, that there is an inverse element is located. Placement tellurium and iodine is not in accordance with the increase of relative atomic mass, were consistent with the increase in atomic number.
Tellurium has the atomic number 52 and iodine has an atomic number of 53. The modern periodic system can be regarded as a refinement Mendeleev periodic system. The modern periodic system, also known as the long form of the periodic system, is based on the atomic number increases, and similar properties. In the modern periodic system of horizontal rows are called periods and upright rows are called groups.
The number of periods in the periodic system there are 7 and marked with numbers:
• Period 1 is referred to as the period is very short and contains 2 elements
• Period 2 is referred to as short a period and contains 8 elements
• Period 3 is referred to as short a period and contains 8 elements
• Period 4 called the period length and contains 18 elements
• Period 5 called the period length and contains 18 elements
• Period of six called the period is very long and contains 32 elements, in this period there are elements of the Lanthanide element number 58 to number 71 and placed on the bottom
• Period 7 is referred to as the period is not yet complete because it may increase the number of elements that occupy it again, until now contains 24 elements. In this period there is a row of elements called Actinides, ie elements numbered 90 to 103 and the number is placed at the bottom.
The number of groups in the periodic system are 8 and marked with Roman numerals. There are two major categories, namely class A (main group) and class B (class transitions). Group B is located between groups IIA and IIIA groups.
The names of the classes to the element of class A
• Group IA alkali group called
• Group IIA alkaline earth group called
• Type IIIA called boron golonga
• Group IVA called carbon groups
• type is called a class of nitrogen VA
• VIA Group called the oxygen group
• Group VIIA group called halogens
• Group VIIIA group called noble gases
In the period 6 class IIIB there are 14 elements that are very similar, the lanthanide elements. In the period 7 also apply the same thing and called the actinide elements. Both series of elements called transition elements inside.
The elements of the lanthanide and actinide belonged IIIB, included in the
group because it has very similar properties.
B. NATURE OF METAL
Properties owned by the elements very much. In this discussion, we will only discuss some properties of the elements. Based on the nature kelogamannya, generally elements are divided into three categories: metals, non-metallic elements, and the element metalloid (semi-metal).
Metals we have encountered many around us, for example iron, aluminum, copper, silver, gold, and others. In general, the metal has physical properties, such as:
1. good heat conductor;
2. good conductor of electricity;
3. shiny metal surfaces;
4. can be beaten into thin plates;
5. can stretch when pulled.
The ability of the metal to stretch when pulled is called ductility. The ability to conduct electricity and stretch metal used to make the wire or cable. The ability to change shape if forged metal called maleabilitas. The ability of a metal to change shape if forged dimanfaatka to make various kinds of goods, such as machetes, knives, hoes, and others. The above properties are not owned by the non-metal elements (non-metallic).
When viewed from the electron configuration, metallic elements tend to lose electrons (having a small ionization energy), while non-metallic elements tend to capture electrons (it has a large ionization energy).
Thus, we can see the tendency of metals in the periodic system properties, ie in a group from top to bottom is getting bigger and in a period from left to right are getting smaller. If we look at the periodic table of elements, metal elements berletak on the left, while the non-metallic elements located on the right (see the periodic table of elements).
On the periodic table, the boundary between metallic elements and nonmetallic often depicted with thick ribbed diagonal ladder. The elements in the border area have a dual nature. For example, the metal beryllium (Be) and aluminum (Al), the metals have some nonmetallic properties, and is called amphoteric elements. The metal which is next door (in the periodic table) are Boron (B) and Silicon (Si) is a non-metallic element, which owns several properties of metals and metalloid elements called
Enterprises grouping based on common elements of nature done for these elements are easy to learn.
1. Triade Dobereiner
In 1829, Johan Wolfgang Dobereiner study the properties of some of the elements are already known at the time. Dobereiner see any resemblance between some elements of nature and the elements are grouped according to similarities in character. Apparently each group consists of three elements of the so-called triade. If the elements in the triade compiled and sorted based on the similarity of its atomic mass, then the second element is an average of the properties and atomic masses of the first and third elements.
2. Octave Theory Newland
In 1864, John Alexander Reina Newland compiled a list of more numerous elements. Newland arrangement shows that if the elements arranged by atomic mass increases, then the first element has similar properties to the eighth element, the second element is similar to the elements of nature ninth, and so on. The discovery was declared a law Newland Newland Octave.
At Newland Octave compiled list, the elements of noble gases (He, Ne, Ar, Kr, Xe, and Rn) has not been found. Noble gases are found by Rayleigh and Ramsay in 1894. Elements of the noble gases argon gas was first discovered. Newland Octave law only applies to elements with low atomic mass.
3. Mendeleev Periodic System
In 1869, began the periodic table arranged. The periodic table is the work of two scientists, Dmitri Ivanovich Mendeleev from Russia and Julius Lothar Meyer from Germany. They work independently and produce a similar table at the same time. Mendeleev presents its results to the Russian Chemical Society in early 1869, and a new periodic table Meyer appeared in December 1869.
Mendeleev who first suggested the periodic table, it is considered as the inventor of the periodic table is often referred to as short-periodic system of elements. Mendeleev developed the periodic table by atomic mass increases and similar properties. Mendeleev periodic system was first published in the scientific journal Annalen der Chemie in 1871.
The important thing there is in the Mendeleev periodic system are as follows:
a. two adjacent elements, has a relative atomic mass difference of at least two or one unit;
b. there are empty boxes for undiscovered elements, such as 44, 68, 72, and 100;
c. can predict the properties of elements that have not been known as ekasilikon;
d. measurement error can be corrected relative atomic mass of some elements, such as Cr = 52.0 not 43.3.
a. Excess Mendeleev periodic system
1) Chemical and physical properties in a single class of similar elements and changed regularly.
2) the highest valence of an element equal to their group number.
3) to predict the properties of elements yet to be discovered at the time and has had a vacant place.
b. Mendeleev periodic system deficiencies
1) The length of the period is not the same and why not described.
2) Some elements are not ordered by their atomic mass increases, for example: Te (128) before I (127).
3) the mass difference of successive elements are not always 2, but the range between 1 and 4 so difficult to predict the mass of an element that has not been known.
4) Valence more than one element of the unpredictable group.
5) anomaly (deviation) elemental hydrogen from other elements not described.
4. Modern Periodic System
In 1914, Henry G. J. Moseley found that the order of the elements in the periodic table according to the atomic number of the element increases. Moseley managed to find errors in the Mendeleev periodic table, that there is an inverse element is located. Placement tellurium and iodine is not in accordance with the increase of relative atomic mass, were consistent with the increase in atomic number.
Tellurium has the atomic number 52 and iodine has an atomic number of 53. The modern periodic system can be regarded as a refinement Mendeleev periodic system. The modern periodic system, also known as the long form of the periodic system, is based on the atomic number increases, and similar properties. In the modern periodic system of horizontal rows are called periods and upright rows are called groups.
The number of periods in the periodic system there are 7 and marked with numbers:
• Period 1 is referred to as the period is very short and contains 2 elements
• Period 2 is referred to as short a period and contains 8 elements
• Period 3 is referred to as short a period and contains 8 elements
• Period 4 called the period length and contains 18 elements
• Period 5 called the period length and contains 18 elements
• Period of six called the period is very long and contains 32 elements, in this period there are elements of the Lanthanide element number 58 to number 71 and placed on the bottom
• Period 7 is referred to as the period is not yet complete because it may increase the number of elements that occupy it again, until now contains 24 elements. In this period there is a row of elements called Actinides, ie elements numbered 90 to 103 and the number is placed at the bottom.
The number of groups in the periodic system are 8 and marked with Roman numerals. There are two major categories, namely class A (main group) and class B (class transitions). Group B is located between groups IIA and IIIA groups.
The names of the classes to the element of class A
• Group IA alkali group called
• Group IIA alkaline earth group called
• Type IIIA called boron golonga
• Group IVA called carbon groups
• type is called a class of nitrogen VA
• VIA Group called the oxygen group
• Group VIIA group called halogens
• Group VIIIA group called noble gases
In the period 6 class IIIB there are 14 elements that are very similar, the lanthanide elements. In the period 7 also apply the same thing and called the actinide elements. Both series of elements called transition elements inside.
The elements of the lanthanide and actinide belonged IIIB, included in the
group because it has very similar properties.
B. NATURE OF METAL
Properties owned by the elements very much. In this discussion, we will only discuss some properties of the elements. Based on the nature kelogamannya, generally elements are divided into three categories: metals, non-metallic elements, and the element metalloid (semi-metal).
Metals we have encountered many around us, for example iron, aluminum, copper, silver, gold, and others. In general, the metal has physical properties, such as:
1. good heat conductor;
2. good conductor of electricity;
3. shiny metal surfaces;
4. can be beaten into thin plates;
5. can stretch when pulled.
The ability of the metal to stretch when pulled is called ductility. The ability to conduct electricity and stretch metal used to make the wire or cable. The ability to change shape if forged metal called maleabilitas. The ability of a metal to change shape if forged dimanfaatka to make various kinds of goods, such as machetes, knives, hoes, and others. The above properties are not owned by the non-metal elements (non-metallic).
When viewed from the electron configuration, metallic elements tend to lose electrons (having a small ionization energy), while non-metallic elements tend to capture electrons (it has a large ionization energy).
Thus, we can see the tendency of metals in the periodic system properties, ie in a group from top to bottom is getting bigger and in a period from left to right are getting smaller. If we look at the periodic table of elements, metal elements berletak on the left, while the non-metallic elements located on the right (see the periodic table of elements).
On the periodic table, the boundary between metallic elements and nonmetallic often depicted with thick ribbed diagonal ladder. The elements in the border area have a dual nature. For example, the metal beryllium (Be) and aluminum (Al), the metals have some nonmetallic properties, and is called amphoteric elements. The metal which is next door (in the periodic table) are Boron (B) and Silicon (Si) is a non-metallic element, which owns several properties of metals and metalloid elements called
Before getting into these trends, we should engage a quick review and establish some terminology. As seen in the previous section on the octet rule, atoms tend to lose or gain electrons in order to attain a full valence shell and the stability a full valence shell imparts. Because electrons are negatively charged, an atom becomes positively or negatively charged as it loses or gains an electron, respectively. Any atom or group of atoms with a net charge (whether positive or negative) is called an ion. A positively charged ion is a cation while a negatively charged ion is an anion.
Now we are ready to discuss the periodic trends of atomic size, ionization energy, electron affinity, and electronnegativity.
Atomic Size (Atomic Radius)
The atomic size of an atom, also called the atomic radius, refers to the distance between an atom's nucleus and its valence electrons. Remember, the closer an electron is to the nucleus, the lower its energy and the more tightly it is held.Moving Across a Period
Moving from left to right across a period, the atomic radius decreases. The nucleus of the atom gains protons moving from left to right, increasing the positive charge of the nucleus and increasing the attractive force of the nucleus upon the electrons. True, electrons are also added as the elements move from left to right across a period, but these electrons reside in the same energy shell and do not offer increased shielding.Moving Down a Group
The atomic radius increases moving down a group. Once again protons are added moving down a group, but so are new energy shells of electrons. The new energy shells provide shielding, allowing the valence electrons to experience only a minimal amount of the protons' positive charge.Cations and Anions
Cations and anions do not actually represent a periodic trend in terms of atomic radius, but they do affect atomic radius, and so we will discuss them here.A cation is positively charged, meaning that it is an atom that has lost an electron or electrons. The positive charge of the nucleus is thus distributed over a smaller number of electrons and electron-electron repulsion is decreased, meaning that the electrons are held more tightly and the atomic radius is smaller than in the normal neutral atom. Anions, conversely, are negatively charged ions: atoms that have gained electrons. In anions, electron-electron repulsion increases and the positive charge of the nucleus is distributed over a large number of electrons. Anions have a greater atomic radius than the neutral atom from which they derive.
Ionization Energy and Electron Affinity
The process of gaining or losing an electron requires energy. There are two common ways to measure this energy change: ionization energy and electron affinity.Ionization Energy
The ionization energy is the energy it takes to fully remove an electron from the atom. When several electrons are removed from an atom, the energy that it takes to remove the first electron is called the first ionization energy, the energy it takes to remove the second electron is the second ionization energy, and so on. In general, the second ionization energy is greater than first ionization energy. This is because the first electron removed feels the effect of shielding by the second electron and is therefore less strongly attracted to the nucleus. If a particular ionization energy follows a previous electron loss that emptied a subshell, the next ionization energy will take a rather large leap, rather than follow its normal gently increasing trend. This fact helps to show that just as electrons are more stable when they have a full valence shell, they are also relatively more stable when they at least have a full subshell.Ionization Energy Across a Period
Ionization energy predictably increases moving across the periodic table from left to right. Just as we described in the case of atomic size, moving from left to right, the number of protons increases. The electrons also increase in number, but without adding new shells or shielding. From left to right, the electrons therefore become more tightly held meaning it takes more energy to pry them loose. This fact gives a physical basis to the octet rule, which states that elements with few valence electrons (those on the left of the periodic table) readily give those electrons up in order to attain a full octet within their inner shells, while those with many valence electrons tend to gain electrons. The electrons on the left tend to lose electrons since their ionization energy is so low (it takes such little energy to remove an electron) while those on the right tend to gain electrons since their nucleus has a powerful positive force and their ionization energy is high. Note that ionization energy does show a sensitivity to the filling of subshells; in moving from group 12 to group 13 for example, after the d shell has been filled, ionization energy actually drops. In general, though, the trend is of increasing ionziation energy from left to right.
Ionization Energy Down a Group
Ionization energy decreases moving down a group for the same reason atomic size increases: electrons add new shells creating extra shielding that supersedes the addition of protons. The atomic radius increases, as does the energy of the valence electrons. This means it takes less energy to remove an electron, which is what ionization energy measures.Electron Affinity
An atom's electron affinity is the energy change in an atom when that atom gains an electron. The sign of the electron affinity can be confusing. When an atom gains an electron and becomes more stable, its potential energy decreases: upon gaining an electron the atom gives off energy and the electron affinity is negative. When an atom becomes less stable upon gaining an electron, its potential energy increases, which implies that the atom gains energy as it acquires the electron. In such a case, the atom's electron affinity is positive. An atom with a negative electron affinity is far more likely to gain electrons.Electron Affinities Across a Period
Electron affinities becoming increasingly negative from left to right. Just as in ionization energy, this trend conforms to and helps explain the octet rule. The octet rule states that atoms with close to full valence shells will tend to gain electrons. Such atoms are located on the right of the periodic table and have very negative electron affinities, meaning they give off a great deal of energy upon gaining an electron and become more stable. Be careful, though: the nobel gases, located in the extreme right hand column of the periodic table do not conform to this trend. Noble gases have full valence shells, are very stable, and do not want to add more electrons: noble gas electron affinities are positive. Similarly, atoms with full subshells also have more positive electron affinities (are less attractive of electrons) than the elements around them.Electron Affinities Down a Group
Electron affinities change little moving down a group, though they do generally become slightly more positive (less attractive toward electrons). The biggest exception to this rule are the third period elements, which often have more negative electron affinities than the corresponding elements in the second period. For this reason, Chlorine, Cl, (group VIIa and period 3) has the most negative electron affinity.
Electronegativity
Electronegativity refers to the ability of an atom to attract the electrons of another atom to it when those two atoms are associated through a bond. Electronegativity is based on an atom's ionization energy and electron affinity. For that reason, electronegativity follows similar trends as its two constituent measures.Electronegativity generally increases moving across a period and decreases moving down a group. Flourine (F), in group VIIa and period 2, is the most powerfully electronegative of the elements. Electronegativity plays a very large role in the processes of Chemical Bonding.
BalasHapusin my article, it is said that the metal oxidation state in accordance with the valence electrons, eg goL IA = +1 and gol IIA = + 2, but why was not the case in non-metals, such as gol VIIA = -1
if you ask me,why metal groups he was its +1 oxaidation state of the element is a means to let go of the outher electrons to be like the class of the element class 7a 8a whereasits oxidation state of -1 because she need 1 electron to become like the noble gases.
BalasHapusI think group VIIA has seven valence electrons, so the halogen group 1 electron capture to stable.
BalasHapusI think that the class 1A or class IIA tend to charge, in order to achieve stability so the group is following the valence electrons, and VIIa padagolongan not like metal because he tends to capture electrons
BalasHapus