Chemistry 1, Chapter 2: Periodic Relationships Among Elements PDF

Chemistry 1 Chapter 2: Periodic relationship among elements See chapter 8 in the text book 1  Outlines I- Periodic table classification II-Periodic Variation in Physical Properties...

Chemistry 1 Chapter 2: Periodic relationship among elements See chapter 8 in the text book 1  Outlines I- Periodic table classification II-Periodic Variation in Physical Properties II.1- Effective nuclear charge II.2-Atomic Radius II.3- Ionic Radius II.4- Ionization Energy II.5-Electron Affinity II.6- Electronegativity 2 I- Periodic table classification I.1- Brief history about the periodic table : -John Newlands (1864): the classification was initially based on increasing atomic weight. In 1869 the Russian chemist Dmitri Mendeleev adopted the same classification of elements based on atomic weight which made possible the prediction of several elements that had not yet been discovered. However, this classification showed later some discrepancies. -Modern classification of the elements consists on the classification of elements by increasing atomic number Z. A chronological chart of the discovery of the elements. This classification allows also to order elements based on the To date, 118 elements have been identified. electron configuration of atoms. 3 I.2- Periodic Classification of the Elements 18 columns called ‘’groups’’ or "families". 7 rows "periods" 4  I.3- Categories in periodic table According to the type of subshell being filled, the elements can be divided into categories: 1- the representative elements, (also called main group elements ) are the elements in Groups 1A through 7A: have incompletely filled s or p subshells of the highest principal quantum number. 2- the noble gases (except He completely filled 1s) :have a completely filled p subshell 3-the transition metals are the elements in Groups 1B and 3B through 8B, which have incompletely filled d subshells or produce cation with incompletely filled d subshells. 4-The Group 2B elements are Zn, Cd, and Hg, which are neither representative elements nor transition metals. 5- the lanthanides and the actinides have incompletely filled f subshells Block p Column number Block d Atomic number Z Block s Block f 5 I.4- Electron configuration of elements per Groups (column) in periodic table The chemical reactivity of the elements is largely determined by their valence electrons, which are the outermost electrons. For the representative elements, the valence electrons are those in the highest occupied n shell. All nonvalence electrons in an atom are referred to as core electrons. Example core electrons valence electrons The electron configurations of the representative elements in a given group display the same number and type of valence electrons. The similarity of the valence electron configurations in the same group implies a resemblance in chemical behavior 6 I.4- Electron configuration of elements per Groups (column) in periodic table Column 1: The alkali metals : all have the valence electron configuration of ns1 and they all tend to lose one electron to form the unipositive cations Column 2:The alkaline earth metals all have the valence electron configuration of ns2 , and they all tend to lose two electrons to form the dipositive cations. Columns 3 to 11: In transition metals the valence electron configuration is (n-1)dx, nsy ( 1 ≤x ≤ 10 and 0 ≤y ≤ 2) Column 17: contains halogens whose valence electron configuration is: ns2 np5 Column 18: noble gases with an external electronic structure is: ns2np6, except for He (1 s2) The noble gas configuration corresponds to the saturation of the electron subshell. The saturation of the ns and np subshells with electrons induces in most cases a lack of chemical reactivity , a condition that often correlates with great stability. 7  The outer electrons ground-state configurations of the elements The International Union of Pure and Applied Chemistry (IUPAC) has recommended numbering the columns sequentially with Arabic numerals 1 through 18 8 I.4- Electron configuration of elements per Groups (column) in periodic table Do not generalize! The similarity of the valence electron configurations in the same group and chemical behavior could not be generalized. Examples: Elements in Group 14 all have the same outer electron configuration, ns2np2, but they display a difference in chemical properties such as: Carbon is a nonmetal, silicon and germanium are metalloids, and tin and lead are metals. The noble gases behave with the exception of krypton and xenon (which can react with some halogens) very similarly. Noble gas are chemically inert due to the completely filled ns and np subshells, a condition that often correlates with great stability. 9 I.4- Electron configuration of elements per Groups (column) in periodic table In general, elements can be classified in a Periodic Table based on three main groups: metals, nonmetals or metalloids. Metals are good conductors of heat and electricity, and are malleable and ductile. Most of the metals are solids at room temperature except for mercury, which is a liquid. Nonmetals are (usually) poor conductors of heat and electricity, and are not malleable or ductile. The majority of the elemental nonmetals are gases at room temperature (H, O, N, etc), while some others can be liquids (Bromine Br) or solids (Carbon). Metalloids are intermediate in their physical properties. They are more like the nonmetals, but under certain circumstances, several of them can be made to conduct electricity. This group includes elements with semiconducting property are extremely important in computers and other electronic devices. 10 II-Periodic Variation in Physical Properties II.1- Effective nuclear charge II.2-Atomic Radius II.3- Ionic Radius II.4- Ionization Energy II.5-Electron Affinity II.6- Electronegativity 11 II-Periodic Variation in Physical Properties II.1- Effective Nuclear Charge The effective nuclear charge (Z eff ) is the nuclear charge felt by an electron when both the actual nuclear charge (Z) and the repulsive effects (shielding) of the other electrons are taken into account. 𝒁∗𝒆𝒇𝒇 = 𝑍 − ෍ σ𝑖→ 𝑗 Generally, the core electrons, which are closer to the nucleus than valence electrons, shield valence electrons much more than valence electrons shield one another. Trends within a period : As we go from the left to the right in a period, the number of core electrons remains constant while the nuclear charge increases. However, because the added electrons is a valence electrons (which do not shield each other well), the net effect of moving across the period is the increase of effective nuclear charge felt by the valence electrons. Trends within a group: Moving from top to bottom, we might expect the elements to have a similar effective nuclear charge as they all have similar total shielding by inner shells. 𝒁∗𝒆𝒇𝒇 remains almost constant. However, because the valence electrons are now added to increasingly large shells as n increases, the electrostatic attraction between the nucleus and the valence electrons decreases. 12 II-Periodic Variation in Physical Properties II.1- Effective Nuclear Charge Example: second period 13 II- Periodic Variation in Physical Properties II.2- Atomic Radius (AR) What’s an atomic radius? It represents one-half the distance between the two nuclei in two adjacent atoms or in a diatomic molecule. AR Trend within a period : the atomic radius decreases from the left to the right, because the effective nuclear charge Zeff increases, which induces that the added valence electron at each step is more strongly attracted by the nucleus than the one before (increase of the nucleus–electrons attraction force). Trend within a group (from the top to the bottom) : the atomic radius size increases with atomic number, because orbital size increases with the increasing principal quantum number n. Ex : Li to Cs 14 II- Periodic Variation in Physical Properties II.2- Atomic Radius (AR) Atomic radii (in picometers) of representative elements according to their positions in the periodic table. Note that there is no general agreement on the size of atomic radii. We focus only on the trends in atomic radii, not on their precise values. 15 II- Periodic Variation in Physical Properties II.3-Ionic Radius Generalities : Ions Derived from Representative Elements Atoms of most representative elements have the tendency to form stable ions with as outer electron configuration the one of noble-gas of ns2 np 6. When cations (Mn+) are formed one or more electrons are removed from the highest occupied n shell. When anions (Mn-) are formed one or more electrons are added to the highest partially filled n shell. Na+ and Ne are isoelectronic (same electron configuration ) H- and He are isoelectronic O2-, N3- ,F-, Na+, and Ne are isoelectronic 16  Cations Derived from Transition Metals Usually, the ns orbital is always filled before the (n-1)d orbitals (i.e: 4s before 3d). However, when a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n - 1)d orbitals (see chapter 1, part 2). The reason is that Electron-electron and electron-nucleus interactions in a neutral atom can be quite different from those in its ion in transition metals and that the (n - 1)d orbital is more stable than ns orbital in transition metal ions. Examples : Mn: [Ar] 4s2 3d5 Mn2+ : [Ar] 3d5 Fe: [Ar] 4s2 3d6 Fe2+ : [Ar] 3d6 Most transition metals can form more than one cation and that frequently the cations are not isoelectronic (same electron configuration) with the preceding noble gases. 17 II- Periodic Variation in Physical Properties II.3-Ionic Radius (IR) What is an ionic radius? Comparison of atomic radii with ionic radii It is the radius of a cation or an anion, which is the distance of the outermost shell of the ion and the nucleus. AR When an anion is formed, its size (or radius) increases, because the IR nuclear charge remains the same but the repulsion resulting from the additional electron(s) enlarges the domain of the electron cloud IR > AR When a cation is formed by removing one or more electrons from an AR atom electron-electron repulsion is reduced but the nuclear charge IR remains the same, so the electron cloud shrinks (get smaller), and the cation is smaller than the atom. IR < AR 18  II- Periodic Variation in Physical Properties II.3-Ionic Radius IR There is a parallel trends between atomic radii and ionic radii. For example, from top to bottom both the atomic radius and the ionic radius increase within a group due to the increase of principal quantum number n. Trend within periods : for ions derived from elements in different groups a size comparison of ionic radius is meaningful only if the ions are isoelectronic A/ When a cation is Isoelectronic to an anion, results that cation has smaller ionic radius. Example 1: Na+ and F- are isoelectronic so they have the same number of electrons IR (Na+) < IR (F-) because Na (Z =11) has more protons than F (Z =9). So, the larger effective nuclear charge of Na+ results in a smaller radius. 19 II- Periodic Variation in Physical Properties II.3-Ionic Radius IR B/ For isoelectronic anions, we find that the radius increases as we go from ions with uninegative charge to those with dinegative charge,etc : IR(X-) < IR(Y2-) < IR(Z3-) Example 2 : N3-, O2- and F- are isoelectronic Because the increasing trend in nuclear charge F (Z =9) > O (Z =8) > N (Z =7) implies the increasing of attraction effect applied by nucleus on electrons -> which results in decrease of ionic radius : IR (N3-)> IR (O2-) > IR (F-) C/ For isoelectronic cations, radii of tripositive ions are smaller than those of dipositive which in turn are smaller than unipositive ions IR(X3+) < IR(Y2+) < IR(Z1+) Example 3: Al3+, Mg2+ and Na+ are isoelectronic IR (Na+)> IR (Mg2+) > IR (Al3+) because for the same number of electrons, Z (Al)> Z(Mg) > Z(Na), so the attraction force is also increasing following the same trend than the nuclear charge. 20 II- Periodic Variation in Physical Properties II.3-Ionic Radius IR Ionic Radius increases The radii (in picometers) of ions of familiar elements arranged according to the elements’ positions in the periodic table. 21 II- Periodic Variation in Physical Properties II.3-Ionic Radius IR Application : For each of the following pairs, indicate which one of the two species is larger: (a) S2- or Cl-(b)Mg2+ or Ca2+ ; (c) Fe2+ or Fe3+ Strategy In comparing ionic radii, it is useful to classify the ions into three categories: (1) isoelectronic ions, (2) ions that carry the same charges and are generated from atoms of the same periodic group, (3) ions carry different charges but are generated from the same atom. In case (1), ions carrying a greater negative charge are always larger; in case (2), ions from atoms having a greater atomic number are always larger; in case (3), ions having a smaller positive charge are always larger. Solution (a) S2- and Cl- are isoelectronic anions, both containing 18 electrons. Because S2- has 16 protons and Cl- has 17, the smaller attraction exerted by the nucleus on the electrons results in a larger S2- ion. (b) Both Mg and Ca belong to the alkaline earth metals. Thus, Ca2+ ion is larger than Mg2+ because Ca’s valence electrons are in a larger shell (n = 4) than are Mg’s (n = 3). (c) Both ions have the same nuclear charge, but Fe2+ has one more electron (24 electrons compared to 23 electrons for Fe3+ ) and hence greater electron-electron repulsion. The radius of Fe2+ is larger. 22 II- Periodic Variation in Physical Properties II.4- Ionization energy Definition : Ionization energy is the minimum energy (in kJ/mol) required to remove an electron from a gaseous atom in its ground state. Ionization energy is the amount of energy in kilojoules needed to strip 1 mole of electrons from 1 mole of gaseous atoms. A (g) A+(g) + 1é 1st ionisation IE 1 A+(g) A2+ (g) + 1é 2nd ionisation IE2 A2+(g) A3+ (g) + 1é 3rd ionisation IE3 The higher the ionization energy, the more difficult it is to remove the electron. For a many-electron atom, the amount of energy required to remove the first electron from the atom in its ground state is called the first ionization energy. 23 II- Periodic Variation in Physical Properties II.4- Ionization energy When an electron is removed from an atom, the repulsion among the remaining electrons decreases. Because the nuclear charge remains constant, more energy is needed to remove another electron from the positively charged ion. Thus, ionization energies always increase in the following order: IE 1 < IE 2 < IE 3 By convention, energy absorbed by atoms (or ions) in the ionization process has a positive value So IE > 0. 24  II- Periodic Variation in Physical Properties II.4- Ionization energy IE 1 < IE 2 < IE 3 … The evolution of nth Ionization Energies (kJ/mol) of the First 20 Elements 25 II- Periodic Variation in Physical Properties Period 1 Period 2 Period 3 II.4- Ionization energy a) In a period, when the atomic number Z increases (from the left to the right): Apart from small irregularities, the first ionization energies of elements in a period increase due to the increase in effective nuclear charge (as in the case of atomic radii variation). Charge effect A larger effective nuclear charge means a more tightly held valence electron, and hence a higher first ionization energy. Variation of the first ionization energy with atomic number. 26 II- Periodic Variation in Physical Properties II.4- Ionization energy b) In a column (group), when the atomic IE number Z increases (from the top to the bottom): - The number n of electronic shells increases - The distance nucleus-valence electron increases (atomic radius increases) IE Distance effect - Attraction force nucleus-valence electrons decreases (electron are more free) First ionization energy decreases within a group from the top to the bottom 27 II- Periodic Variation in Physical Properties II.4- Ionization energy Comparison of IE between groups of elements: Alkali metals group have the lowest first ionization energies: these metals have one valence electron (the outermost electron configuration is ns1), which is effectively shielded by the completely filled inner shells. Consequently, it is energetically easy to remove an electron from the atom of an alkali metal to form a unipositive ion (Li+, Na+, K+) Alkaline earth metals have higher first ionization energies than the alkali metals do. The alkaline earth metals have two valence electrons which do not shield each other well, the effective nuclear charge for an alkaline earth metal atom is larger than that for the preceding alkali metal (attraction forces between nucleus and electrons increases as Zeff increases). Metals have relatively low ionization energies compared to nonmetals. The ionization energies of the metalloids generally fall between those of metals and nonmetals. The difference in ionization energies suggests why metals always form cations and nonmetals form anions in ionic compounds. 28 II- Periodic Variation in Physical Properties II.5- Electron affinity EA Another property that greatly influences the chemical behavior of atoms is their ability to accept one or more electrons. This property is called electron affinity, which is defined as the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. The electron affinity of an atom can be defined by this process: ΔE(attach) EA = −ΔE(attach) Example : Gaseous fluorine atom accepts an electron ΔE(attach) ΔE(attach)= - 328 kJ/mol The electron affinity of fluorine is therefore assigned a value of EA = −ΔE(attach) = + 328 kJ/mol 29  II- Periodic Variation in Physical Properties II.5- Electron affinity Thus, a large positive electron affinity means that the negative ion is very stable (that is, the atom has a great tendency to accept an electron), just as a high ionization energy of an atom means that the electron in the atom is very stable. Trend of electron affinity in periodic table : The overall trend is an increase in the tendency to accept electrons (electron affinity values become more positive) from left to right across a period. EA Electron affinity varies approximatively according the same trend than ionization energy. EA Unlike ionization energy (always positive), electron affinity can be positive or negative. Electron affinity is positive if the reaction is exothermic and negative if the reaction is endothermic. 30 II- Periodic Variation in Physical Properties II.5- Electron affinity ▪ The halogens (Group 7A) have the highest electron affinity values. Noble gases have high effective nuclear charge, they have extremely low electron affinities (or zero). The reason is that an electron added to an atom with an ns2np6 configuration has to enter an (n + 1)s orbital, where it is well shielded by the core electrons and will only be very weakly attracted by the nucleus. This analysis also explains why species with complete valence shells tend to be chemically stable. ▪ The electron affinities of metals are generally lower than those of nonmetals. The values vary little within a given group. 31 II- Periodic Variation in Physical Properties II.6- Electronegativity Electronegativity, is the ability of an atom to attract toward itself the electrons in a chemical bond. Elements with high electronegativity have a greater tendency to attract electrons than do elements with low electronegativity. As we might expect, electronegativity is related to electron affinity and ionization energy. Mulliken scale The electronegativity of an element in the Mulliken scale is equal to the arithmetic average of the first ionization energy, IE1, and electron affinity EA (𝐼𝐸1 +𝐸𝐴) 𝐸𝑁 = 2 Thus, an atom such as fluorine, which has a high electron affinity (tends to pick up electrons easily) and a high ionization energy (does not lose electrons easily), has a high electronegativity. 32  Electronegativity and electron affinity are related but different concepts. Both indicate the tendency of an atom to attract electrons. However, electron affinity refers to an isolated atom’s attraction for an additional electron, whereas electronegativity signifies the ability of an atom in a chemical bond (with another atom) to attract the shared electrons. Furthermore, electron affinity is an experimentally measurable quantity, whereas electronegativity is an estimated number that cannot be measured. 33 II- Periodic Variation in Physical Properties II.6- Electronegativity Electronegativity decreases within a group from the top to the bottom and increase within a period from the the left to the right. The most electronegative elements are : the halogens, oxygen, nitrogen, and sulfur, which are found in the upper right-hand corner of the periodic table, and the least electronegative elements. Note that the transition metals do not follow these trends 34 II- Periodic Variation in Physical Properties II.6- Electronegativity Variation of electronegativity with atomic number. The halogens have the highest electronegativities, and the alkali metals the lowest. 35