Principles of Chemistry. The Development of Periodic Table. Mendeleev s Predictions. A Molecular Approach, 1 st Ed. Nivaldo Tro

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1 Principles of Chemistry A Molecular Approach, 1 st Ed. Nivaldo Tro Chapter 8 Periodic Properties of the Elements The Development of Periodic Table Dmitri Mendeleev( ) Arranged elements in a table form by atomic mass (modern PeriodicTable) Periodic law: When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically. put elements with similar properties in the same column used patterns to predict properties of undiscovered elements where atomic mass order did not fit other properties, he reordered by different properties I.e. Te and I Mendeleev s Predictions

2 Quantum Theory and Periodic table Mendeleev s periodic law- predicts what the properties of an element will be based on its position on the table Quantum theory explains why the periodic trends in the properties exist. Quantum theory describes the behaviour of electrons in atoms Helps us to understand the chemical behaviour Chemical bonding involve the transfer or sharing of electrons, and Spin quantum number(m s ): Electron Spin Spin is a fundamental property of all electrons All electrons have the same amount of spin The orientation of the electron spin is quantized with only two possibilities spin up or spin down Spin quantum number (m s ) describes how the electron spins on its axis. m s = +½ ( )or (-½ ( ). Opposite spins must cancel in an orbital(paired) Orbital Diagram Pauli Exclusion Principle Pauli Exclusion Principle- No two electrons in an atom can have the same four quantum numbers Each orbital can have a maximum of only two electrons with opposite spins only Degenerate orbital- orbitals of equal energy e.g. three p orbitals, five d orbitals or seven f orbitals The maximum number of electrons in each subshell Subshell Number of Orbitals Maximum Number of Electrons s (l = 0) 1 2 p (l = 1) 3 6 d (l =2) 5 10 f (l =3) 7 14

3 Electron Configurations Electron configuration- The distribution of electrons into the various orbitals in an atom in its ground state Ground State-of an atom is the lowest energy state The notation for a configuration Number designates the principal energy level. The letter designates the sublevel and type of orbital. The superscript designates the number of electrons in that sublevel. For example He (Z=2) has ground state electron configuration He = 1s 2 Orbital Diagram: We can represent an orbital as a box and the electrons in that orbital as arrows orbital with 2 electrons Sublevel Splitting in Multielectron Atoms Single electron atoms (H and H like)-the sublevels in each principal energy level of hydrogen all have the same energy (degenerate) Energy of hydrogen sublevels depends on principal quantum number n Multielectron atoms, the energies of the sublevels are split Energy of the sublevels depends on n and l caused by electron electron repulsion The lower the value of the quantum number l, the less energy the sublevel has. s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3) Penetrating and Shielding The effective nuclear charge is the positive charge an electron experiences from the nucleus minus any shielding effects from intervening electrons

4 Orbital energies for multielectron atoms Important Points 1. Because of penetration, sublevels within an energy level are not degenerate. 2. Penetration of the fourth and higher energy levels is so strong that 4s orbital lies lower in energy than the 3d orbitals and 5s orbital lies lower in energy than the 4d orbitals. 3. The energy difference between levels becomes smaller for higher energy levels (and can cause anomalous electron configurations for certain elements). Electron Configurations Every atom has an infinite number of possible electron configurations Ground state electron configuration:the configuration associated with the lowest energy level of the atom Other configurations correspond to excited states By filling orbital of lowest energy first you usually get the lowest energy (ground state) of the atom

5 Electron Configurations for Multielectron Atoms Aufbau principle(build up)- Electron occupy orbitals so as to minimize the energy of the atom, therefore orbitals fill from lowest energy to highest Order of orbital filling 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f Pauli Exclusion principle- No two electrons in one atom can have the same four quantum numbers (max. 2 e- with opposite spin) Hund s rule-when orbitals of same energy are available, electron first occupy these orbitals singly with parallel spin rather than in pairs Once the orbitals of equal energy are half full, the electrons start to pair. Another way to remember the order of orbital filling Start by drawing a diagram with each energy shell on one row, and list the subshells (s, p, d, f) for that shell in order of energy (left to right). Draw arrows down through the diagonals, looping back to the next diagonal each time. Writing Electron Configuration of Atoms in Their Ground State Filling order s p d f The # of e- in an atom = Z (atomic number Full configuration: C (Z= 6) 6 electrons C: 1s 2 2s 2 2p 2 Inner electron configuration -A shorthand way of writing an electron configuration use the symbol of the previous noble gas(noble gas core) in brackets to represent all the inner electrons, then just write the last set e.g. Na= 11 electrons = 1s 2 2s 2 2p 6 3s 1 Na [Ne]3s 1

6 Electron Configurations (Z= 3-10) Electron configuration, Valence electrons, and the Periodic Table Valence electrons-the electrons in the highest principal energy level(outermost shell) e - s outside the noble gas core are called valence electrons. Important in chemical bonding Core electrons-electrons in lower energy shells (inner shells) Both chemical and physical properties of an atom, are determined by the number of valence electrons Sample Problem Write electron configuration for the following and identify the valence electron and core electron Al Z= 13 1s 2 2s 2 2p 6 3s 2 3p 1 Kr Z= 36 Rb Z= 37

7 Electron Configuration and the Periodic Table The periodic table is divided into four blocks corresponding to the filling of the four quantum sublevels (s,p, d, and f) The number of columns in each block is the maximum number of electrons that sublevel can hold For the main group elements (recall chapt.2) group # = Total #of valence electrons Period # of element = Principal quantum # of the outer shell Periodic table and electron configuration Periodic Table can be divided into four regions or blocks according to the orbitals being fil led 1A 2A 3A 4A 5A 6A 7A 8A P Ne

8 Periodic table and electron configuration s- block elements - Group 1A and 2A (filling of s orbital) p-block elements - Group 3A through 8A (filling of p orbital; ns orbitals are already filled) d-block elements- Transition metals (filling of (n-1)d orbitals) f-block elements- Inner transition metals also called Lanthanide and actinide (filling of (n-2)f orbital Sample Problem Which of the following orbital filling diagrams represent : 1) ground state or 2) excited state or 3) forbidden for Boron (Z= 5) 1s 2s 2p ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) ( ) Transition elements Overlap of principal energy levels- After Ar the next electron will enter into the lowest sublevel of 4th principal level 4s instead of the highest sub level of third principal level (3d) e.g. Ca (Z = 20) Ca [Ar], 4s 2 Elements with Z = (Sc to Zn) After 4s is filled, 3d will start Energy of 4s orbital < 3d orbital and so on 1s, 2s, 2p,3s, 3p, 4s, 3d At Z= 36 Kr, 4p is completely filled. 4p 6

9 Anomalous Electron Configurations Half filled and filled subshells have an unusual stability Leads to anomalies in electron configuration Sometimes s electron is promoted to d subshell (anamolous configuration) Results in lowering the total energy of the atom. Due to decrease in electron- electron repulsions Anomalies occur where the energy differences between subshells are small i.e. 4s and 3d or 5s and 4d Z > 40 Example: Give an abbreviated electronic configuration of Ag Solution: Ag is one of the elements that will transfer an electron from one sub-shell to another in order to lower the total energy of the atom. Therefore e- configuration of Ag [Kr] 5s 1 4d 10 Exceptions to the Aufbau principle Some of the transition metals have anomalous electron configurations in which the ns only partially fills before the (n-1)d or doesn t fill at all Therefore, their electron configurations must be found experimentally (XPS spectroscopy) Expected (Aufbau principle) Experimental Cr = [Ar]4s 2 3d 4 Cr = [Ar]4s 1 3d 5 Cu = [Ar]4s 2 3d 9 Cu = [Ar]4s 1 3d 10 Mo = [Kr]5s 2 4d 4 Mo = [Kr]5s 1 4d 5 Ru = [Kr]5s 2 4d 6 Ru = [Kr]5s 1 4d 7 Pd = [Kr]5s 2 4d 8 Pd = [Kr]5s 0 4d 10 Quantum mechanical Model and chemical properties Quantum mechanical model accounts for the chemical properties of the elements The chemical properties of elements are largely determined by their # of valence e - s Elements in the same column have similar chemical and physical properties? same # of valence e - s

10 Noble Gas Electron Configuration Quantum-mechanical calculations show that an atom with eight valence electrons should be unreactive. The noble gases have eight valence electrons (ns 2 np 6 ) very stable configuration Except for He, which has only two electrons Properties- noble gases are especially nonreactive He and Ne are practically inert Everyone Wants to Be Like a Noble Gas! The alkali and alkaline earth metals (Group IA & IIA) have one and two more electron than the previous noble gas Very reactive metals Valence shell configuration ns 1 and ns 2 In their reactions, They tend to lose their extra electrons, resulting in the same electron configuration as a noble gas e.g. Na + and Ca 2+ The halogens (Group7A)- valence shell e-configuration (ns 2 np 5 ) All are one electron short of the next noble gas configuration Most reactive nonmetals- form anion with 1 charge with nonmetals, they tend to share electrons Trend in Atomic Radius Main Group van der Waals radius (nonbonding atomic radius)- One half the distance between adjacent nuclei in the atomic solid Covalent radius (bonding atomic radius) One-half the distance between the nuclei of two identical covalently bonded atoms Approximate bond length = sum of the atomic radii of the two covalently bonded atoms

11 Periodic Trends in size of atoms- Main Group The general trends in the atomic radii of main group elements in the periodic table Atomic radius increases down a column (or group) The principle quantum # n increases resulting in larger orbitals effective nuclear charge fairly close Atomic radius decreases across period (from left to right) adding electrons to same valence shell effective nuclear charge increases valence shell held closer Trends in Atomic Radius Factors determining the size of an atom Two factors determine the size of an atom 1. The principal quantum number, n-as the value of n increases the size of the outer orbital increases (resulting in larger atoms) 2.The effective nuclear charge (Z eff )- The positive charge an electron experiences from the nucleus minus any shielding effects from intervening electrons

12 Atomic Radii and the transition elements In general the atomic radii of the transition elements do not follow the same trend as main group elements increase in size down the group Atomic radii of transition metals stay roughly constant across each row. Why? The # of electrons in the outermost principal energy level (highest n value) is nearly constant valence shell ns 2, not the d electrons effective nuclear charge on the ns 2 electrons approximately the same Electron Configuration and Ion Charge Ions are formed by loss (cation + ve) or gain (anion -ve) of electrons by a neutral atom the charge on an ion is predictable based on its position on the periodic table. Electron configuration of the ions is the same as the nearest noble gas. Isoelectronic - species with same number of electrons. Ne and Al 3+ Ions of the main- group elements The common monatomic ions found in compounds of the main group elements fall into three catagories 1. Cations of group IA-IIIA: Positively charged ions The ion charges = group # 2. Cations of group IIIA-VA (e - config. ns 2 ): The ion charges = Group # e.g. Tl + Sn 2+, Pb 2+ and Bi Anions of Groups VA VIIA:The Ion charges = group# - 8 Cations form when the atom loses electrons from the valence shell. Al (Z=13) = 1s 2 2s 2 2p 6 3s 2 3p 1 Al 3+ ion = 1s 2 2s 2 2p 6

13 Electron Configurations of Ions Transition metals- The s electrons are first in with the atoms and first out with the cations. Lose electrons from the valence shell first, which is not the last sublevel to fill according to the Aufbau sequence For example, zinc generally loses two electrons from its 4s sublevel to adopt a pseudo -noble gas configuration [Ar]4s 2 3d 10 [Ar]3d 10 Magnetic Properties of Transition Metal Atoms and Ions Only atoms with unpaired electrons exhibit magnetic susceptibility A paramagnetic substance is one that is weakly attracted by a magnetic field, usually the result of unpaired electrons A diamagnetic substance is not attracted by a magnetic field generally because it has only paired electrons slightly repelled by a magnetic field Zn atom and Zn 2+ ions are diamagnetic. Zn [Ar]4s 2 3d 10 Zn 2+ [Ar]3d 1 Example 8.6 Write the electron configuration and determine whether the Fe atom and Fe 3+ ion are paramagnetic or diamagnetic. previous noble gas = Ar 18 e Fe atom = [Ar]4s 2 3d 6 unpaired electrons paramagnetic 4s 3d Fe 3+ ion = [Ar]4s 0 3d 5 unpaired electrons paramagnetic 4s 3d

14 Periodic Properties The periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. We will look at the following periodic properties: Atomic or ionic radius Ionization energy Electron affinity Metallic character Trends in Ionic Radius Ion size increases down the group and decreases across a row Why? higher valence shell = larger size Cations are smaller than the neutral atom Na [Ne] 3s 1 and Na + [Ne] Anions are bigger than the neutral atom Why? cations smaller than anions for isoelectronic species-species with same number of electrons. larger positive charge = smaller cation larger negative charge = larger anion

15 Order the following sets by size (smallest to largest). Zr 4+, Ti 4+, Hf 4+ same column and charge;therefore, Ti 4+ < Zr 4+ < Hf 4+ Na +, Mg 2+, F -, Ne isoelectronic; therefore, Mg 2+ < Na + < Ne < F - I -, Br -, Ga 3+, In + Ga 3+ < In + < Br - < I - Periodic Trends: Ionization Energy Ionization Energy-Minimum energy needed to remove an electron from an atom or ion in the gaseous state Units (ev) 1 ev = 96.5 kj mol -1 Endothermic process First ionization energy (IE 1 ) energy required to remove electron from neutral atom M(g) M + (g) + 1 e - IE 1 = kj mol -1 Second IE 2 energy required to remove an electron from 1+ ion M + (g) M 2+ (g) + 1 e - IE 2 = kj mol -1 General Trends in First Ionization Energy For main group elements: Ionization energy decreases down the group valence electrons farther from nucleus experience less Z eff Ionization energy generally increases across the period effective nuclear charge increases and electrons are held tightly by the nucleus

16 First ionization Energy versus At. # Irregularities in the Trend Which is easier to remove, an electron from B or Be? Why? Be(Z = 4) and B(Z= 5) Which is easier to remove an electron from, N or O? Why? N(Z =7) and O(Z=8) Trends in Successive Ionization Energies Removal of each successive electron costs more energy. Regular increase in energy for each successive v. e - Large increase in energy when start removing core electrons

17 Trends in Successive Ionization Energies Electron Affinity The electron affinity is the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion For a chlorine atom, the first electron affinity is illustrated by: Cl(g) + 1 e - Cl - (g) EA= -349 kj mol -1 Defined as exothermic (-), but may actually be endothermic (+) Some alkali earth metals and all noble gases are endothermic. Why? Trends in Electron Affinity The more energy that is released, the larger the electron affinity of the atom. The more negative the number, the larger the EA, and the more stable the negative ion that is formed Most groups in P.T do not exhibit any definite trend in EA Generally EA become more negative (adding an e- become more exothermic) as you move to the right across a period Highest EA in period = halogen

18 Properties of Metals and Nonmetals Metals malleable and ductile shiny, lustrous, reflect light conduct heat and electricity most oxides are basic and ionic form cations in solution lose electrons in reactions oxidized Nonmetals brittle in solid state dull, nonreflective solid surface electrical and thermal insulators most oxides are acidic and molecular form anions and polyatomic anions gain electrons in reactions reduced General Trends in Metallic character Metallic character The tendency to lose electrons decreases from left to right across a period. Metals are found at the left side and middle of the period Table increases down the column. Nonmetals are found at the top right of the P.T

19 Operational Skills Applying the Pauli exclusion principle and applying Hund s rule to write electron configuration Writing electron configurations of atoms and ions using the Aufbau principle Determining the electron configuration using the period and group numbers Writing orbital diagrams Identifying valence electrons and core electrons Applying periodic trends to predict ion size, relative ionization energies and metallic character

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