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Variation of Physical Properties Across a Menstruation

The physical backdrop of elements vary across a period, mostly as a part of bonding.

Learning Objectives

Draw the general variations in concrete properties beyond a row of the periodic table.

Key Takeaways

Primal Points

• Every bit yous move from left to right across a menstruation, the physical properties of the elements alter.
• Ane loose trend is the tendency for elemental states to go from solid to liquid to gas across a menstruation. In the farthermost cases, Groups one and 18, we see that Grouping-1 elements are all solids and Group-18 elements are all gases.
• Many of the changes in physical properties every bit you cantankerous a period are due to the nature of the bonding interactions that the elements undergo. The elements on the left side of a period tend to class more than ionic bonds, while those on the correct side grade more covalent bonds.

Key Terms

• boiling point: The temperature at which a liquid boils, with the vapor pressure equal to the given external pressure.
• melting point: The temperature at which the solid and liquid phases of a substance are in equilibrium; it is relatively insensitive to changes in pressure.

The periodic table of elements has a total of 118 entries. Elements are bundled in a series of rows (periods) in social club of atomic number so that those with similar backdrop appear in vertical columns. Elements in the aforementioned menses have the same number of electron shells; moving across a period (so progressing from grouping to grouping), elements proceeds electrons and protons and become less metallic. This system reflects the periodic recurrence of similar properties every bit the atomic number increases. For example, the alkali metals lie in i group (Group 1) and share similar properties, such as high reactivity and the tendency to lose one electron to arrive at a noble-gas electron configuration.

Modernistic quantum mechanics explains these periodic trends in properties in terms of electron shells. The filling of each shell corresponds to a row in the table.

In the s-block and p-cake of the periodic table, elements inside the same period mostly practice not exhibit trends and similarities in properties (vertical trends downwards groups are more significant). However, in the d-block, trends across periods go significant, and the f-cake elements show a loftier caste of similarity beyond periods (peculiarly the lanthanides).

If we examine the concrete land of each element, we notice that on the left side of the table, elements such as lithium and glucinium are metallic solids, whereas on the right, nitrogen, oxygen, fluorine, and neon are all gases. This is considering lithium and beryllium course metal solids, whereas the elements to the correct form covalent compounds with little intermolecular strength belongings them together. Therefore we tin say that, in general, elements tend to go from solids to liquids to gases equally we motion beyond a given period. However, this is non a strict trend.

Bonding

As yous motility across a period in the periodic table, the types of unremarkably encountered bonding interactions change. For example, at the beginning of Catamenia 2, elements such equally lithium and beryllium class only ionic bonds, in full general. Moving across the period, elements such every bit boron, carbon, nitrogen and oxygen tend to form covalent bonds. Fluorine tin can form ionic bonds with some elements, such as carbon and boron, and neon does not tend to form any bonds at all.

Melting Points of the Halides

Some other concrete property that varies across a period is the melting point of the corresponding halide. A halide is a binary compound, of which i part is a halogen atom and the other part is an element or radical that is less electronegative (or more electropositive) than the halogen, to make a fluoride, chloride, bromide, iodide, or astatide compound. Many salts are halides; the hal- syllable in halide and halite reflects this correlation. All Group 1 metals form halides that are white solids at room temperature.

The melting point is correlated to the strength of intermolecular bonds within the chemical element. Get-go, we must analyze compounds formed from elements from Groups i and ii (due east.g., sodium and magnesium). To develop an understanding of bonding in these compounds, nosotros focus on the halides of these elements. The physical properties of the chlorides of elements in Groups ane and 2 are very different compared to the chlorides of the elements in Groups 4, 5, and vi.

All of the alkali halides and element of group i earth halides are solids at room temperature and take melting points in the hundreds of degrees centigrade. For example, the melting point of sodium chloride (NaCl) is 808 °C. In contrast, the melting points of the not-metallic halides from Periods two and 3, such as CCl₄, PCl_iii, and SCl₂, are below 0 °C, so these materials are liquids at room temperature. Furthermore, all of these compounds have low humid points, typically in the range of fifty °C to fourscore °C.

Melting and boiling points of various halides
Halide	Melting Signal (ºC)	Boiling Signal (ºC)
LiCl	610º	1382º
BeCl2	405º	488º
CCl4	-23º	77º
NClthree	−40º	71º
OCltwo	−20º	4º
FCl	−154º	−101º
NaCl	808º	1465º
MgCltwo	714º	1418º
SiCl4	-68º	57º
PCl3	−91º	74º
SCl2	−122º	59º
Cl2	−102º	−35º
KCl	772º	1407º
CaCl	772º	> 1600º

The not-metal halide liquids are also electrical insulators and do non conduct electrical current. In contrast, when an alkali halide or alkaline earth halide melts, the resulting liquid is an splendid electrical conductor. This tells u.s. that these molten compounds consist of ions, whereas the non-metal halides practise not. This again demonstrates the type of bonding that these compounds exhibit: the left-most elements grade more ionic bonds, and the further-right elements tend to form more than covalent bonds.

Variation of Concrete Properties Within a Group

The physical properties (notably, melting and humid points) of the elements in a given group vary as you move downwardly the table.

Learning Objectives

Describe the full general trends of physical backdrop within a group on the periodic table.

Fundamental Takeaways

Cardinal Points

• The physical properties of elements depend in part on their valence electron configurations. As this configuration remains the same within a grouping, physical properties tend to remain somewhat consistent.
• The about notable within-group changes in physical backdrop occur in Groups 13, 14, and 15, where the elements at the tiptop are non-metallic, while the elements at the bottom are metals.
• The trends in humid and melting points vary from group to group, based on the type of not-bonding interactions holding the atoms together.

Central Terms

• physical property: Whatsoever holding that is measurable whose value describes a physical system's land.
• malleable: Able to be hammered into thin sheets; capable of being extended or shaped by beating with a hammer or by the pressure of rollers.
• ductile: Capable of being pulled or stretched into thin wire past mechanical force without breaking.

In chemistry, a group is a vertical column in the periodic table of the chemic elements. At that place are 18 groups in the standard periodic table, including the d-block elements but excluding the f-block elements. Each chemical element within a group has like physical or chemical properties because of its cantlet's outermost electron trounce (most chemical properties are dominated by the orbital location of the outermost electron).

Common Physical Backdrop

A physical belongings of a pure substance tin can be defined as anything that can exist observed without the identity of the substance changing. The observations commonly consist of some type of numerical measurement, although sometimes there is a more than qualitative (non-numerical) clarification of the belongings. Physical properties include such things equally:

• Color
• Brittleness
• Malleability
• Ductility
• Electrical electrical conductivity
• Density
• Magnetism
• Hardness
• Atomic number
• Specific heat
• Heat of vaporization
• Heat of fusion
• Crystalline configuration
• Melting temperature
• Boiling temperature
• Heat conductivity
• Vapor pressure
• Trend to deliquesce in various liquids

These are only a few of the measurable concrete properties.

Within a group of the periodic tabular array, each chemical element has the aforementioned valence electron configuration. For example, lithium, sodium, potassium, rubidium, cesium, and francium all have a single electron in an south orbital, whereas every chemical element in the group including fluorine has the valence electron configuration ns²np⁵, where n is the period. This means the elements of a group often showroom similar chemic reactivity, and there may exist similarities in physical properties as well.

Boiling and Melting Points

Before a discussion of the melting points of diverse elements, information technology should be noted that some elements exist in different forms. For example, pure carbon tin be as diamond, which has a very high melting point, or as graphite, whose melting indicate is still loftier but much lower than that of diamond.

Unlike groups exhibit different trends in boiling and melting points. For Groups 1 and two, the humid and melting points subtract as yous move down the grouping. For the transition metals, boiling and melting points mostly increase equally you lot move down the group, but they decrease for the zinc family. In the main group elements, the boron and carbon families (Groups thirteen and xiv) subtract in their boiling and melting points every bit you move downwards the group, whereas the nitrogen, oxygen, and fluorine families (Groups xv, 16, and 17) tend to increase in both. The noble gases (Grouping eighteen) decrease in their boiling and melting points down the group.

These phenomena can be understood in relation to the types of forces holding the elements together. For metallic species, the metallic bonding interaction (electron-sharing) becomes more difficult every bit the elements become larger (toward the lesser of the table), causing the forces holding them together to become weaker. Every bit you move right along the tabular array, withal, polarizability and van der Waals interactions predominate, and as larger atoms are more polarizable, they tend to showroom stronger intermolecular forces and therefore higher melting and humid points.

Metal Character

Metallic elements are shiny, normally grayness or argent in color, and conductive of oestrus and electricity. They are malleable (tin be hammered into sparse sheets) and ductile (tin exist stretched into wires). Some metals, such as sodium, are soft and can be cutting with a pocketknife. Others, such as iron, are very hard. Non-metallic atoms are dull and are poor conductors. They are brittle when solid, and many are gases at STP (standard temperature and pressure). Metals give away their valence electrons when bonding, whereas non-metals tend to accept electrons.

A metal and a non-Metallic: On the left is sodium, a very metallic chemical element (ductile, malleable, conducts electricity). On the correct is sulfur, a very non-metal element.

Metallic character increases from correct to left and from top to bottom on the tabular array. Non-metallic grapheme follows the opposite pattern. This is considering of the other trends: ionization free energy, electron affinity, and electronegativity. You lot will detect a jagged line running through the periodic table starting between boron and aluminum – this is the separation between metallic and non-metallic elements, with some elements close to the line exhibiting characteristics of each. The metals are toward the left and heart of the periodic table, in the southward, d, and f blocks. Poor metals and metalloids (somewhat metal, somewhat non-metal) are in the lower left of the p block. Not-metals are on the right of the tabular array.

Electron Configurations and Magnetic Properties of Ions

The electron configuration of a given chemical element can be predicted based on its location in the periodic table.

Learning Objectives

Predict the type of ions an element will form based on its position in the periodic table

Key Takeaways

Primal Points

• The electron configuration of an element dictates the element's backdrop in a chemic reaction. Electron configurations vary regularly along the periodic table.
• The Aufbau principle determines the electron configuration of an element. The principle states that the everyman- energy orbitals are filled first, followed successively by higher-energy orbitals.
• Magnetism can outcome from unpaired electrons in a given ion of an element, depending on the spin states of the electrons.

Key Terms

• electron configuration: The organisation of electrons in an atom, molecule, or other concrete structure, such as a crystal.

Blocks of the Periodic Table

The periodic tabular array does more than than just list the elements. The give-and-take "periodic" ways that within each row, or menstruation, the elements show a design of characteristics. This is because the elements are listed in part by their electron configuration.

Blocking in the periodic table: The periodic table can be cleaved into blocks, corresponding to the highest energy electrons.

The alkali metals and alkaline earth metals have one and two valence electrons (electrons in the outer trounce), respectively; because of this, they lose electrons to form bonds easily and so are very reactive. These elements comprise the s cake of the periodic tabular array. The p block, on the correct, contains common non-metals, such as chlorine and helium. The noble gases, in the column on the correct, nigh never react, since they have eight valence electrons forming a stable outer shell. The halogens, directly to the left of the noble gases, readily gain electrons and react with metals. The s and p blocks make upwardly the principal- group elements, also known as representative elements. The d block, which is the largest, consists of transition metals, such every bit copper, fe, and gold. The f block, on the lesser, contains rarer metals, including uranium. Elements in the aforementioned group or family have the same configuration of valence electrons, then they behave in chemically similar ways.

Periodic table of the elements: This image is colour-coded to show the south, p, d, and f blocks of the periodic tabular array.

Electron Configuration

In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule in diminutive or molecular orbitals. For instance, the electron configuration of the neon atom (Ne) is 1sⁱⁱ 2s² 2p^six. According to the laws of quantum mechanics, a sure energy is associated with each electron configuration. Under sure conditions, electrons can movement from one orbital to another by emission or absorption of a quantum of energy, in the form of a photon.

Knowledge of the electron configurations of unlike atoms is useful in understanding the structure of the periodic tabular array. The concept is also useful for describing the chemical bonds that hold atoms together. In bulk materials, this same idea helps explain the peculiar properties of lasers and semiconductors.

The idea of an electron configuration was first conceptualized under the Bohr model of the atom, and it is still common to speak of "shells" and "subshells" despite the advances in agreement of the quantum-mechanical nature of electrons.

Aufbau Principle

The Aufbau principle (from the German Aufbau, meaning "building up, construction;" also called the Aufbau rule or edifice-up principle) is used to determine the electron configuration of an atom, molecule, or ion. The principle postulates a hypothetical process in which an atom is "built up" by the progressive add-on of electrons. Every bit electrons are added, they assume their most stable positions (electron orbitals) with respect to the nucleus and the electrons that are already there.

According to the principle, electrons fill orbitals starting at the everyman available energy land earlier filling higher states (east.g., 1s earlier 2s). The number of electrons that can occupy each orbital is limited by the Pauli exclusion principle. If multiple orbitals of the same energy are available, Hund's rule states that unoccupied orbitals volition be filled before occupied orbitals are reused (by electrons having dissimilar spins).

Atomic orbitals ordered by increasing energy: Order in which orbitals are bundled by increasing energy according to the Madelung rule. Each diagonal red arrow corresponds to a unlike value of north + l.

Magnetism

Magnetism is a property of materials that respond to an practical magnetic field. Permanent magnets have persistent magnetic fields caused by ferromagnetism, the strongest and nigh familiar type of magnetism. However, all materials are influenced differently by the presence of a magnetic field. Some are attracted to a magnetic field (paramagnetism); others are repulsed by it (diamagnetism); still others have a much more complex relationship with an applied magnetic field (e.g., spin-glass behavior and antiferromagnetism). Substances that are negligibly affected past magnetic fields are considered non-magnetic, these are: copper, aluminum, gases, and plastic. Pure oxygen exhibits magnetic properties when cooled to a liquid country.

The magnetic backdrop of a given chemical element depend on the electron configuration of that element, which will change when the chemical element loses or gains an electron to form an ion. If the ionization of an element yields an ion with unpaired electrons, these electrons may align the sign of their spins in the presence of a magnetic field, making the material paramagnetic. If the spins tend to align spontaneously in the absence of a magnetic field, the resulting species is termed ferromagnetic.

Hierarchy for various types of magnetism: There are various types of magnetism identified to date that can be organized in a bureaucracy.

Applications of Magnetism

A lodestone, or loadstone, is a naturally magnetized slice of the mineral magnetite (Fe₃O₄). Ancient people first discovered the belongings of magnetism in lodestone. Pieces of lodestone, suspended and so they could plough, were the kickoff magnetic compasses, and their importance to early navigation is indicated by their very proper noun, which in Middle English means "course stone" or "leading stone." Lodestone is i of only 2 minerals that is institute naturally magnetized; the other, pyrrhotite, is only weakly magnetic.

Diminutive Radius

Atomic radii decrease from left to right across a menstruation and increase from top to bottom along a group.

Learning Objectives

Predict the relative atomic sizes of the elements based on the full general trends in atomic radii for the periodic table.

Primal Takeaways

Cardinal Points

• The atomic radius of a chemical element is a measure out of the size of its atoms, usually the mean or typical distance from the nucleus to the boundary of the surrounding cloud of electrons.
• Since the boundary is not a well-defined physical entity, in that location are various non-equivalent definitions of atomic radius.
• The periodic trends of the diminutive radii (and of diverse other chemical and physical properties of the elements) can be explained by the electron shell theory of the cantlet.

Key Terms

• breakthrough theory: A theory adult in early 20th century, according to which nuclear and radiation phenomena can be explained by assuming that energy but occurs in discrete amounts called quanta.
• electron beat: The collective states of all electrons in an cantlet having the same principal quantum number (visualized as an orbit in which the electrons movement).
• element of group 0: Any of the elements of Group 18 of the periodic tabular array, being monatomic and (with very limited exceptions) inert.

In chemistry, periodic trends are the tendencies of certain elemental characteristics to increase or decrease as one progresses along a row or column of the periodic tabular array of elements. The atomic radius is i such characteristic that trends across a period and downward a group of the periodic table.

Periodic trends: A graphic showing overall periodic trends in the periodic table.

Meaning of the Atomic Radius

The atomic radius of a chemical element is a measure of the size of its atoms, unremarkably the hateful or typical distance from the nucleus to the purlieus of the surrounding deject of electrons. Since the boundary is not a well-defined concrete entity, there are diverse not-equivalent definitions of atomic radius.

Depending on context, the term atomic radius may apply only to isolated atoms, or also to atoms in condensed matter, covalently spring in molecules, or in ionized and excited states. The value of an atomic radius may exist obtained through experimental measurements or computed with theoretical models. Nether some definitions, the value of a radius may depend on the atom's state and context. For our purposes, we are generally looking at atoms in their elemental country.

Sizes of atoms and their ions in picometers (pm): Red numbers are ionic radii of cations, black numbers are for neutral species, and blue numbers are for anions.

Atomic radii vary in a predictable and explicable manner across the periodic tabular array. Radii more often than not decrease from left to correct along each catamenia (row) of the tabular array, from the brine metals to the noble gases; radii increment down each group (column). The radius increases sharply between the noble gas at the end of each period and the brine metal at the beginning of the next period. These trends of the atomic radii (and of diverse other chemical and physical properties of the elements) tin be explained by the electron shell theory of the cantlet. Radii measurements provided important evidence for the evolution and confirmation of quantum theory.

Explanation of the General Trends

The style diminutive radius varies with increasing diminutive number tin be explained by the arrangement of electrons in shells of fixed capacity. Shells closer to the nucleus—those with a smaller radius—are by and large filled first, since the negatively charged electrons are attracted by the positively charged protons in the nucleus. Every bit the atomic number increases forth a row of the periodic tabular array, additional electrons are added to the same, outermost trounce. The radius of this shell gradually contracts every bit the attraction between the additional electrons and the nucleus increases. In a noble gas, the outermost shell is completely filled. Therefore, the additional electron of next brine metallic (1 row downwards on the periodic tabular array) volition get into a new outer shell, accounting for the sudden increase in the diminutive radius.

Atomic number to radius graph: A chart showing the atomic radius relative to the diminutive number of the elements.

The increasing nuclear accuse is partly balanced by the increasing number of electrons, a phenomenon that is known as shielding; this explains why the size of atoms ordinarily increases down each column. Underlying causes of the periodic trends in diminutive radius too have an impact on other chemic and physical properties of the elements.

Ionic Radius

Similarly charged ions tend to decrease in size across a period (row) and increase in size downwardly a group (cavalcade).

Learning Objectives

Identify the general trends of the ionic radius size for the periodic table.

Primal Takeaways

Cardinal Points

• The ionic radius is the distance between the nucleus and the electron in the outermost trounce of an ion.
• When an atom loses an electron to grade a cation, the lost electron no longer contributes to shielding the other electrons from the charge of the nucleus; consequently, the other electrons are more than strongly attracted to the nucleus, and the radius of the atom gets smaller.
• When an electron is added to an atom, forming an anion, the added electron repels other electrons, resulting in an increase in the size of the cantlet.
• The tendency observed in size of ionic radii is due to shielding of the outermost electrons by the inner-shell electrons and then that the outer crush electrons do non "feel" the entire positive accuse of the nucleus.

Cardinal Terms

• cation: A positively charged ion, as opposed to an anion.
• ion: An atom or grouping of atoms bearing an electric charge, such as the sodium and chlorine atoms in a salt solution.
• anion: A negatively charged ion, as opposed to a cation.

In chemistry, periodic trends are the tendencies of certain elemental characteristics to increase or subtract forth a flow (row) or grouping (cavalcade) of the periodic table of elements. Ionic radius (r _ion) is the radius of an ion, regardless of whether information technology is an anion or a cation. Although neither atoms nor ions have sharp boundaries, it is useful to treat them as if they are hard spheres with radii. In this way, the sum of ionic radii of a cation and an anion can give us the altitude between the ions in a crystal lattice. Ionic radii are typically given in units of either picometers (pm) or Angstroms (Å), with 1 Å = 100 pm. Typical values range from 30 pm (0.3 Å) to over 200 pm (2 Å).

Trends in Ionic Radii

Ions may be larger or smaller than the neutral atom, depending on the ion's charge. When an atom loses an electron to form a cation, the lost electron no longer contributes to shielding the other electrons from the charge of the nucleus; consequently, the other electrons are more strongly attracted to the nucleus, and the radius of the atom gets smaller. Similarly, when an electron is added to an atom, forming an anion, the added electron repels other electrons, resulting in an increase in the size of the atom.

The ionic radius is non a fixed property of a given ion; rather, it varies with coordination number, spin land, and other parameters. For our purposes, nosotros are considering the ions to be as close to their ground state as possible. However, ionic radius values are sufficiently transferable to allow periodic trends to exist recognized.

Sizes of atoms and their ions: Relative sizes of atoms and ions. The neutral atoms are colored grey, cations crimson, and anions blue.

Equally with other types of atomic radii, ionic radii increase upon descending a group and decrease going beyond a period. Note that this only applies if the elements are the aforementioned type of ion, either cations or anions. For case, while neutral lithium is larger than neutral fluorine, the lithium cation is much smaller than the fluorine anion, due to the lithium cation having a different highest energy shell.

Ionization Energy

The ionization energy tends to increase every bit one moves from left to right across a given period or up a group in the periodic table.

Learning Objectives

Recognize the general periodic trends in ionization energy.

Key Takeaways

Cardinal Points

• The ionization free energy is the energy required to remove an electron from its orbital around an atom to a point where it is no longer associated with that cantlet.
• The ionization energy of an chemical element increases as one moves beyond a period in the periodic table considering the electrons are held tighter by the college effective nuclear accuse.
• The ionization energy of the elements increases every bit ane moves up a given group because the electrons are held in lower-energy orbitals, closer to the nucleus and therefore are more tightly jump (harder to remove).

Key Terms

• ionization energy: The free energy needed to remove an electron from an atom or molecule to infinity.

Periodic Trends in the Ionization Energy

The ionization energy of a chemic species (i.e., an cantlet or molecule ) is the free energy required to remove electrons from gaseous atoms or ions. This property is also referred to as the ionization potentia and is measured in volts. In chemistry, it ofttimes refers to ane mole of a substance (tooth ionization energy or enthalpy) and is reported in kJ/mol. In diminutive physics, the ionization free energy is typically measured in the unit of measurement electron volt (eV). Large atoms or molecules have depression ionization free energy, while small molecules tend to have higher ionization energies.

The ionization energy is different for electrons of different atomic or molecular orbitals. More generally, the nth ionization energy is the energy required to strip off the nth electron afterwards the showtime n-1 electrons have been removed. It is considered a measure of the tendency of an atom or ion to give up an electron or the strength of the electron binding. The greater the ionization energy, the more than hard it is to remove an electron. The ionization energy may be an indicator of the reactivity of an element. Elements with a low ionization energy tend to be reducing agents and form cations, which in turn combine with anions to grade salts.

Ionization energy: This graph shows the beginning ionization energy of the elements in electron volts.

Moving left to right inside a period or upwardly inside a grouping, the first ionization energy generally increases. Equally the diminutive radius decreases, it becomes harder to remove an electron that is closer to a more positively charged nucleus. Conversely, as one progresses down a group on the periodic table, the ionization energy will likely subtract since the valence electrons are farther away from the nucleus and experience greater shielding. They experience a weaker allure to the positive charge of the nucleus. Ionization energy increases from left to right in a menstruation and decreases from top to bottom in a group.

Rationale for the Periodic Trends in Ionization Energy

The ionization energy of an element increases as one moves across a period in the periodic table because the electrons are held tighter past the college effective nuclear charge. This is considering additional electrons in the same shell do not substantially contribute to shielding each other from the nucleus, yet an increase in atomic number corresponds to an increment in the number of protons in the nucleus.

The ionization energy of the elements increases as 1 moves up a given grouping considering the electrons are held in lower-energy orbitals, closer to the nucleus and thus more tightly bound (harder to remove).

Based on these two principles, the easiest element to ionize is francium and the hardest to ionize is helium.

Periodic trends in ionization energy – YouTube: This video explains the periodic trends in ionization energy….periodicity. The periodic tabular array is bundled in a mode to prove trends in the characteristics of the elements.

Electron Affinity

The electron affinity of the elements more often than not increases across a catamenia and sometimes decreases downwardly a group in the periodic table.

Learning Objectives

Recognize the full general periodic trends for electron analogousness.

Key Takeaways

Primal Points

• The electron affinity of an cantlet or molecule is the propensity for that particle to gain an electron. This is an exothermic process for all non-element of group 0 elements.
• There are general trends in electron affinity across and downward the periodic table of elements. Electron affinity generally increases across a period in the periodic table and sometimes decreases downwards a grouping.
• These trends are non necessarily universal.
• The chemical rationale for changes in electron affinity beyond the periodic table is the increased effective nuclear charge across a period and upwardly a group.

Key Terms

• electron analogousness: The electron affinity of an atom or molecule is defined as the amount of free energy released when an electron is added to a neutral atom or molecule to form a negative ion.
• electronegativity: The trend of an atom or molecule to attract electrons to itself.

The electron affinity (Due east_ea) of a neutral atom or molecule is defined equally the amount of free energy released when an electron is added to it to grade a negative ion, equally demonstrated by the following equation:

[latex]X(thou) + e^- \rightarrow X^{-}(g)[/latex]

Electron affinity is measured for atoms and molecules in the gaseous state just, since in the solid or liquid states their energy levels would be changed by contact with other atoms or molecules. Robert S. Mulliken used a list of electron affinities to develop an electronegativity calibration for atoms by finding the average of the electron affinity and ionization potential. A molecule or atom that has a more than positive electron analogousness value is often called an electron acceptor; i with a less positive electron affinity is called an electron donor. Together they may undergo accuse-transfer reactions.

To utilise electron affinities properly, it is essential to keep rail of the sign. For any reaction that releases free energy, the alter in energy (ΔE) has a negative value, and the reaction is called an exothermic process. Electron capture for nearly all non-noble gas atoms involves the release of energy and therefore is an exothermic procedure.

Periodic Properties: Part 4, Ionic Charges, Ionization Free energy, Electron Affinity – YouTube: We conclude our discussion of periodic properties by wrapping up the prediction of ionic charges of the transition metals, ionization energies, and electron affinity.

Confusion may arise in mistaking E_ea for ΔE. The numbers listed in tables of East_ea are all positive because they are magnitudes; the values of Eastward_ea in a table of electron affinities all signal the amount of energy released when an electron is added to an element. Because the release of energy is always an exothermic upshot, these all correspond to negative values of ΔE (indicating an exothermic process).

Periodic Trends in Electron Analogousness

Although E_ea varies greatly across the periodic tabular array, some patterns emerge. By and large, nonmetals take more positive E_ea than metals. Atoms, such as Group vii elements, whose anions are more stable than neutral atoms take a higher E_ea. The electron affinities of the noble gases have not been conclusively measured, and so they may or may not take slightly negative values. Chlorine has the highest Due east_ea while mercury has the lowest.

East_ea by and large increases across a period (row) in the periodic table, due to the filling of the valence beat out of the cantlet. For instance, within the same period, a Group-17 atom releases more energy than a Grouping-ane atom upon gaining an electron because the added electron creates a filled valence shell and therefore is more stable.

A trend of decreasing E_ea downwards the groups in the periodic tabular array would be expected, since the boosted electron is entering an orbital farther abroad from the nucleus. Since this electron is further away, it should exist less attracted to the nucleus and release less energy when added. Nonetheless, this trend applies only to Group-1 atoms. Electron affinity follows the tendency of electronegativity: fluorine (F) has a higher electron affinity than oxygen (O), and and then on.

The trends noted here are very similar to those in ionization energy and change for like (though opposing) reasons.

Electron affinities in the periodic table: This table shows the electron affinities in kJ/mol for the elements in the periodic table.

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