1 Chemistry 11 Unit 8 Atoms and the Periodic Table Part IV Chemical Bonding
2 1. Trends in periodic table In the previous section we have studied the periodic table and some trends dictated by it. In particular, we have looked at the following trends: (1) Metallic character The metallic character of elements increases from top right to bottom left of the periodic table. (2) Reactivity For metals: increases down a group For non-metals: decreases down a group
3 These trends are inferred by a vast number of experimental observations. How do we explain them theoretically? There are 3 factors which are usually discussed when trying to explain these trends: (i) nuclear charge (ii) number of shells (iii) electron-electron repulsion
4 (i) Nuclear charge This charge refers to the total charges present at the nucleus. Since neutrons are electrically neutral, only protons, each of which has +1 charge, will contribute to this. Therefore, the nuclear charge of an atom is equal to the number of protons (or its atomic number). Increasing atomic number implies a larger nuclear charge. However, this is not the actual charge felt by the electrons surrounding the nucleus! It is due to the shielding effect of electrons!
5 Recall that electrons in an atom occupy orbitals which are organized in shells. The lower energy orbitals are located closer to the nucleus than the higher energy counterparts. Therefore, the inner-shell electrons block the attraction of the nucleus to the outer-shell electrons.
6 This reduction of nuclear attraction due to inner-shell electrons is called shielding effect or screening effect. The actual charge felt by the electrons under the influence of shielding is called effective nuclear charge. As implied, the shielding is strong if there are more inner-shell electrons. This results in a weaker attraction on the outer-shell electrons, making it easier to remove the electrons away from the atom.
7 There are 2 trends we can deduce regarding nuclear charge: (a) Across a period While atomic number increases, the number of innershell electrons remains the same. Hence, the effective nuclear charge increases. (b) Down a group Both atomic number and number of inner-shell electrons increases, thereby causing the effective nuclear charge to be very similar.
8 (ii) Number of shells As more electrons are present in an atom, more orbitals will be occupied. According to the Aufbau principle, electrons fill from the lowest energy orbital up. From the analysis of hydrogen atom, we know that the higher the principal quantum number n, the higher the orbital energy. E n 13.6Z2 n 2 This approximation works also for atoms bigger than hydrogen (with atomic number Z). ev
9 In the meantime, it can also be derived from the Bohr s model that orbital energy is inversely proportional to the radius of the orbital. E n 1 r n Combining these two, we can conclude that higher energy orbitals, whose n values are larger, are farther away from the nucleus. r n n 2 (The full derivation of this relation requires both classical mechanics and quantum mechanics.)
10 (iii) Electron-electron repulsion Recall that electrons are particles with one unit of negative charge. When they are put close to one another, they repel because of electrostatic force. For a poly-electron atom, the electrons do feel each other although they may or may not be in the same orbital, subshell or shell. This leads to a repulsion which keeps them separated as much as possible. The magnitude of repulsion increases with increasing number of electrons. That means, the higher the atomic number, the stronger the electron-electron repulsion.
11 Using these 3 factors, we will be able to predict the trends of the following quantities which are essential to understanding the nature of chemical bonds. (1) Atomic radius Going across a period, the 3 factors show the following trends: Nuclear charge: increases Number of shells: same Electron-electron repulsion: increases
12 What do they tell us? (1) Because of the increasing effective nuclear charge, the attraction between the nucleus and electrons is stronger, making the atoms smaller. (2) Since the number of shells remains unchanged, the size of the atoms remains similar. (3) More electrons in the atoms cause electrons to be farther away from one another, making the atoms larger. Consequently, the size of atoms gets smaller when going across a period.
13 Going down a group, the 3 factors show the following trends: Nuclear charge: increases Number of shells: increases Electron-electron repulsion: increases How do we interpret these trends? (1) While nuclear charge increases, the number of shells also increases; this results in a similar effective nuclear charge to the outer-shell electrons. (2) The increased electron-electron repulsion pushes the outer-shell electrons farther away.
14 (3) Increasing number of shells also means the outershell to be bigger. Consequently, the size of the atoms gets bigger going down a group.
15 Example: What is the largest atom in the periodic table? Francium (why?) Example: Which is larger: (a) F or B? B (why?) (b) Ba or Be? Ba (why?) (c) Sr or P? Sr (why?)
16 How about the trends of the sizes of ions? For cations: Since cation is formed by removing electrons from a neutral atom, cation has less electron-electron repulsion and net positive charge that pulls the electrons closer. Therefore, a cation is smaller than the corresponding atom.
17 For anion: Negative ion is obtained by adding electrons to a neutral atom. Because of this, there will be a net negative charge. Meanwhile, the electron-electron repulsion is enhanced, causing the electrons to be farther apart. Therefore, the ion becomes bigger than the atom it forms from.
18 For your reference:
19 An interesting aspect of ionic radii is to compare the sizes of isoelectronic species. For example, consider Ne (~0.76Å) (1) Na + > Mg 2+ > Al 3+ Due to increasing net positive charge (2) O 2- > F - Due to stronger repulsion between electrons (3) Na + > Ne Can t be compared as they are measured differently
20 (2) Ionization energy (IE) This is the energy required to remove an electron from a neutral atom. Depending on which electron is removed, IE can be divided into first IE, second IE, and so on. For example, First IE: X (g) X + (g) + e Second IE: X + (g) X 2+ (g) + e Remember energy is always necessary to break the attraction between the nucleus and electrons.
21 When going across a period: Since nuclear charge increases, electrons are bounded more tightly, rendering them more difficult to remove. The number of shells remains the same, so that the electrons are not farther away from the nucleus. There are more electrons in the outermost shell and the electron-electron repulsion is increased. Hence, the actual outcome is that the electrons are attracted more strongly to the nucleus, and ionization energy increases.
22 When going down a group: While nuclear charge is increased, the number of shells increases as well, making the effective nuclear charge essentially the same. But increasing the number of shells causes the electrons to be farther away from the nucleus. The increasing number of electrons leads to a bigger electron-electron repulsion. Consequently, the electrons become less bounded and the ionization energy decreases.
23 The experimental first ionization energies of the elements in the periodic table.
24 Important implications of the trend of IE: (a) Noble gas elements are very stable because their first IE are very high! This explains why they are chemically unreactive. (b) Indeed, all species having a noble gas configurations (i.e., fully filled s- and p-subshells) are stable. This accounts for the anomalies of the trends of first and second IE for some elements.
25 (c) The favorable noble gas configurations also explains why main group elements tend to form ions with characteristic charges. At these states, the ions are isoelectronic with the nearest noble gas. Element Ionic charge Electron configuration Oxygen -2 [He] 2s 2 2p 4 + 2e - [Ne] Fluorine -1 [He] 2s 2 2p 5 + e - [Ne] Neon 0 [He] 2s 2 2p 6 = [Ne] Sodium +1 [Ne] 3s 1 e - [Ne] Magnesium +2 [Ne] 3s 2 2e - [Ne]
26 Example: Which atom has the highest first ionization energy? Helium (why?) Example: Which has the lower first ionization energy: (a) F or B? B (why?) (b) Ba or Ba 2+? Ba (why?) (c) Sr or P? Sr (why?)
27 (3) Electronegativity This term refers to the ability of an atom to attract electrons in a chemical bond. In a chemical bond, two atoms are fighting for the electrons by electrostatic force. The more powerful one will drag the electrons closer to it. If its electronegativity is high enough, it could even seize the electrons completely and form ions. What can we predict about the trends of electronegativity using the 3 factors?
28 When going across a period: Since the nuclear charge increases while the number of shells remains the same, the effective nuclear charge increases, and it is easier to attract electrons from the atoms nearby. On the other hand, the added electrons will strengthen the electron-electron repulsion, causing the resulting ion to be not stable. Combining these factors, the electronegativity does increase across a period.
29 When going down a group: The increased nuclear charge and orbital size cancel out the influence of one another, making the effective nuclear charge essentially the same. However, the increased electron-electron repulsion pushes the electrons away from each other. Since the nuclear attraction is approximately the same yet the atomic radius gets larger, the power of drawing electrons from neighboring atoms will be smaller. Consequently, electronegativity decreases down a group.
30 The idea of electronegativity is rather abstract. Is it possible to quantify or to develop a scale so that comparison can be made more clearly? There are 2 ways of solving this problem. (a) Pauling scale of electronegativity It was developed by Linus Pauling (1901-1994) A very famous chemist who contributed to the theory of chemical bonds He has received 2 Nobel prizes!
31 In this scale, fluorine is the most electronegative element while francium is least electronegative. They are given values of 4.0 and 0.7 respectively. Other elements will have electronegativity values between them.
32 (b) Electron affinity (EA) When an electron is added to an atom, due to electrostatic attraction, energy is released. This energy change is called electron affinity. X g + e X (g) H = E. A. Electron affinity is closely related to electronegativity because the larger the electronegativity, the stronger the attraction between the nucleus and the newly added electron, and the more the energy released.
33 The experimentally measured electron affinities of elements are shown below. The elements with * are believed to have a negative electron affinity.
34 There is a great similarity between the trends of IE and electronegativity. Both are increasing across a period and decreasing down a group. This can be accounted for by considering the fact that if an atom is capable of attracting strongly the electrons from a neighboring atom, the attraction to its own electrons will have to be strong as well. That s why in general, if an atom has a high electronegativity, it will also have a high IE.
35 Summary of trends that we have explored so far:
36 2. Chemical bonds (1) Infant stage of the theory of bonding Since the early 19 th century, chemists started to develop theories which explain the combinations of chemical atoms by means of their electrical characters. This led to the development of the theory of valency by Edward Frankland, F.A. Kekulé, A.S. Couper, Alexander Butlerov, and Hermann Kolbe around 1850 s, which stated that compounds were joined by the attraction of positive and negative poles.
37 Richard Abegg (1869-1910) was a pioneer of valence theory. He made use of the concepts of valence to account for the behaviors of elements in chemical reaction, and attributed the low reactivity of some elements to their stable shell configurations. In 1904, Richard Abegg postulated that for main group elements the difference between the maximum positive and negative valence (that is, the maximum and minimum oxidation states) of an element is often eight. This is called the Abegg s rule.
38 (2) Early valence bond theory The idea of the Abegg s rule stimulated the further investigation of the nature of the interaction that binds atoms together. In the year of 1916, Walther Kossel (1888-1956) proposed the theory of ionic chemical bonds while Gilbert Lewis (1875-1946) put forth the concepts of cubical atoms, covalent bonds, electron pairs and Lewis dot structures.
39 The ideas of Kossel and Lewis about chemical bonds were further developed by Irving Langmuir (1881-1957), and in 1919 he published his work in which he first described the sharing of electron pair between atoms as a covalent bond. Langmuir popularized the concepts of cubical octet atom and covalency in chemistry, and extend it to something which is called the octet rule nowadays.
40 (3) The Langmuir-Lewis theory There are several important elements in the Langmuir- Lewis theory: (i) Valence electrons play an essential role in chemical bonding. (ii) Electrons are transferred when a chemical bond is formed so that atoms can achieve a more stable electron configuration (usually Noble gas configurations). (iii) Depending on the extent of electron transfer, chemical bonds are either ionic or covalent.
41 Lewis proposed electron dot diagrams to show the valence electrons of atoms and ions. Each diagram consists of a chemical symbol representing the nucleus and core electrons, and dots representing the valence electrons.
42 (a) Ionic bonds If the electron transfer is complete between two atoms, a pair of ions will be formed. These ions are attracted by each other by electrostatic force. This kind of bonding is called ionic bond. Since electrostatic attraction is not fully directional, ions of opposite charges can be bounded together to produce ionic crystal, in which each cation is surrounded by anions and vice versa. The chemical composition of an ionic crystal is represented by a formula unit, which is the smallest collection of ions that is electrically neutral.
43 Since ionic crystal is made of a gigantic number of ions, it is not a single molecule. It is usually referred to as an ionic lattice. Ionic bond is strong; therefore, ionic crystals usually have high melting points. When dissolved, ionic compounds become mobile ions and they can conduct electricity.
44 (b) Covalent bonds If electrons are not fully transferred but instead shared between two nuclei, then a covalent bond will be formed. Note that in a covalent bond, electrons are attracted by both nuclei, although they may stay more time closer to one than the other due to the difference in their ability to capture electrons.
45 The pair of electrons involved in the formation of covalent bond is called bonding pair. Depending on the number of bonding pairs present, a chemical bond may be classified into Bond order # bonding pair Type of bond 1 1 Single 2 2 Double 3 3 Triple Covalent bonds are strong, but they are intramolecular (within the molecule). Therefore, covalent compounds show a wide range of melting points.
46 How do we decide how many bonding pairs will be formed in a covalent bond? This is dictated by the octet rule which states that atoms in groups 14 to 17 tend to form covalent bonds until they are surrounded by 8 valence electrons. The pair of electrons not involved in bonding is called lone pair. Bonding pair in Lewis dot structure can be represented by a line.
47 But there are some exceptions: (1) Hydrogen does not fulfil the octet rule when forming covalent bonds. This is due to the fact that the 1s orbital can maximum hold 2 electrons. (2) Many heavy p-block elements can have more than 8 valence electrons when forming covalent compounds by making use of their empty d-orbitals which are low in energy. It is called the expansion of octet. Less than an octet Expanded octet
48 (c) Polar covalent bonds As discussed, the sharing of electrons is not necessarily even between the two nuclei. Electrons will be leaning toward the nucleus which is more attractive. What is the parameter that controls how much electrons are transferred and shared?
49 The type of bond is based on the difference in electronegativity of the two atoms. Recall this value indicates how powerful an atom is to seize electrons. The following is a general guideline. Type of bond Difference in electronegativity Covalent < 0.5 Polar covalent 0.5 1.6 Ionic > 2.0 For the difference of 1.6 2.0, if the bond involves a metal, the bond is ionic; if only non-metals are involved, the bond is polar covalent.
50 Example: Determine the types of bonds present in the following compounds: (i) NaCl, (ii) CCl 4, (iii) F 2. (i) NaCl: E.N. = 0.9 3.0 = 2.1 Therefore ionic (ii) CCl 4 : E.N. = 2.5 3.0 = 0.5 Therefore polar covalent (iii) F 2 : E.N. = 4.0 4.0 = 0.0 Therefore covalent
51 (d) Intermolecular forces When two atoms with different electronegativity are bonded covalently, the bonding pair is not shared equally. The more electronegative one drags the pair closer to it, causing the result bond to be polar. In a polar bond, because of an uneven distribution of electrons, there exist a partial positive charge and a partial negative charge on the more electron-deficient and the more electron-rich ends respectively.
52 The presence of partial charges on a bond forms a dipole. Dipoles can attract each other in the same way as ions. The resulting interaction, called dipole-dipole interaction, binds polar molecules together. Remember that dipole-dipole interaction is not a true bond, and is weak in general. Dipole-dipole force is denoted by a dotted line.
53 Adding the dipole of each polar bond in a molecule yields the overall molecular dipole. (Remember that dipole is a vector quantity!) It means that molecules that have polar covalent bonds do not necessarily possess molecular dipole. It depends on the 3D geometry of the molecule. Molecules having molecular dipole are polar, while those without it are non-polar.
54 A special type of dipole-dipole interactions is called hydrogen-bond in which the H atom bonded to an electronegative atom (e.g. O) in one molecule is attracted intermolecularly by a lone pair of an electronegative atom (e.g. F, O or N) of another molecule.
55 Non-polar molecules do not experience dipoledipole attraction as they are lacking molecular dipoles. However, they are still weakly attracted by one another by means of a temporary dipole caused by the instantaneous partial charges resulting from an unsymmetrical distribution of electrons.
56 The nature of this instantaneous dipole attraction was first discussed by Fritz London (1900-1954) in 1930 using quantum mechanics. This London force exists between any molecules (either polar or non-polar) when they are in close proximity, even when there are ionic or covalent bonds. London force is the weakest among all intermolecular forces. London force < Dipole dipole < Hydrogen bond
57 Since the London force arises from the temporary dipole, in general, the more electrons an atom/molecule has, the stronger the temporary dipole, and the stronger the London force. Hence, the strength of London force depends on the number of electrons. This is in turn related to the atomic number and mass of an atom.
58 London dispersion force accounts for the low but not zero melting points of non-polar substances such as noble gases, alkanes, etc. That also explains why the boiling point of alkanes increases with increasing chain length.
59 3. Lewis structures of molecules As introduced by G. Lewis in 1916, Lewis structures, or electron dot structures, are diagrams that show the connections of atoms in a molecule and lone pairs of electrons that may exist in the molecule. Lewis structures are used to depict the structures of covalent compounds of main-group elements, and coordination compounds of main-group and transition elements, although sometimes they can be used to describe the structures of ionic compounds.
60 The conventions of drawing Lewis structures: (1) Each atom is represented by its chemical symbol. (2) Valence electrons are represented by dots. (3) A bond pair of electrons between atoms is described by a solid line. (4) A lone pair of electrons on an atom is represented by a pair of dots, and is placed beside the atom. (5) A dative bond (i.e., a pair of electrons donated by one atom only) is represented by an arrow.
61 (a) For ionic compounds To construct the Lewis structure of an ionic compound, draw the dot structure for each of the ions involved including the ion charge. Then put the ions besides one another. Usually metal ion is placed at the middle, surrounded by the nonmetal ions.
62 Exercise: Draw Lewis structures for the following ionic salts. (a) KBr (b) MgCl 2 (c) Li 2 S (d) K 3 P
63 (b) For covalent compounds The procedures of drawing Lewis structures of covalent compounds are as follows: Step 1: Count the total number of valence electrons for the molecules and adjust according to the charge on the molecule. Step 2: Propose the connectivity of atoms in the molecule based on chemical intuition. Place 2 electrons for each bond between two atoms. Step 3: Place the remaining valence electrons on the atoms around the central atom(s) according to the octet.
64 Step 4: If there are still remaining electrons, put them in pairs on the central atom(s). Step 5: If the central atom has less than 8 electrons, try to move a lone pair on a neighboring atom to become a bond pair between them. Repeat this step if necessary. For example: Draw the Lewis structures for the following molecules. (a) H 2 (b) Cl 2 (c) HCl (d) H 2 O (e) NH 3 (f) CH 4
65 Answers:
66 Let s try something harder! Exercise: Draw the Lewis structures for the following molecules. (a) CO 2 (b) N 2 (c) C 2 H 3 Br (d) C 2 HI (e) HCN
67 Answers:
68 As discussed before (refer to p.47), some molecules do not obey the octet rule. These molecules contain either electron-deficient atoms or the elements from the third or fourth row. Electron-deficient elements refer to the elements from Group 2 and Group 13 (e.g. Be, B and Al). Be can share 4 electrons while B and Al can share 6 electrons in maximum.
69 Elements in the 3 rd and 4 th rows have accessible, empty d-orbitals which can hold electrons. Therefore, they can expand their octet and form more than four covalent bonds. Their Lewis structures can be drawn in the same way except adding more bond pairs to the central atom.
70 If a molecule cannot be expressed uniquely by one Lewis structure, then resonance structures will exist. They are alternate Lewis diagrams for the same molecule. e.g. NO 3 - e.g. HPO 3 2-