Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
WAVE NATURE OF ELECTRONS IN ATOMS Electrons have wave properties Can not specify exact location of a wave, or of an electron Electrons occupy orbitals in atoms size, shape, names Are quantized QUANTUM NUMBERS (1) Principal quantum number (n) values: 1, 2, 3, … average distance from nucleus: size energy of electrons in orbital (2) Azimuthal quantum number (A) values: 0, 1, 2, …, (n–1) shape of electron orbital (3) Magnetic quantum number values: −A, …, −1, 0, 1, …, A orientation of orbital in space (4) Spin quantum number two values: +½ and –½ (ms) (mA) Symbols used for A quantum numbers A value 0 1 2 3 subshell name s p d f no. of electrons n=1 n=2 n=2 n=3 A=0 A=0 A=1 A=0 1s 2s 2p 3s (mA = 0) (mA = 0) (mA = –1,0,1) (mA = 0) Table 6.2 n2 = number of states = number of orbitals in shell n 2 2 6 2 s orbitals First s orbital: 1s n=1 A=0 mA = 0 Second s orbital: 2s n=2 A=0 mA = 0 SHAPES OF ORBITALS s orbitals electron density or probability Ψ2 (1s) 0 r at nucleus, r = 0 2s is larger than 1s Size of s orbital increases as n increases Shape: spherical symmetry p and d orbitals First p orbitals: 2p n=2 same size & shape A=1 mA = –1, 0, 1 } 3 different orientations First d orbitals: 3d n=3 A=2 mA = –2, –1, 0, 1, 2 5 orbitals 5 orientations SHAPES OF ORBITALS p orbitals 2 lobes with node between 2p n=2 A=1 mA = –1, 0, 1 3 orbitals Because n same & A same, they have same size & shape They differ in orientation p orbitals are directional SHAPES OF ORBITALS d orbitals 3d n=3 A=2 mA = –2, –1, 0, 1, 2 5 orbitals with different orientations SHAPES OF ORBITALS sphere 1s dumbbell 2pz 3d yz 2px 2py 3d xy 3d xz clover-leaf & friend 3d x 2 -y2 3d z2 REVIEW ORBITALS – region of space with size, shape, characteristic energy Name s p d f Number Shape 1 spherical 3 dumbbell 5 5 shapes 7 ------ QUANTUM NUMBERS n A mA principal size azimuthal shape magnetic orientation FOURTH QUANTUM NUMBER Electron has magnetic moment, as if were spinning Experimental observations confirm ms = ½ or –½ Electrons have 4 quantum numbers n A mA ms defines orbital Since ms has only 2 values Therefore, max of 2 electrons per orbital Pauli Exclusion Principle STERN-GERLACH EXP. Silver (Ag) atoms have one unpaired electron. A beam of Ag atoms splits according to sign of electron spin in magnetic field PAULI EXCLUSION PRINCIPLE No two electrons in an atom can have the same four quantum numbers ( n A mA ms ) Electrons in the same orbital have the same values for the first three quantum numbers ( n A mA ) ms can have only two values: ½ or –½ Therefore, an orbital can hold only two electrons, and they must have opposite spins Subshell s (A = 0) p (A = 1) d (A = 2) f (A = 3) No. of Orbitals Max. No. of e– 1 2 3 6 5 10 7 14 ENERGIES OF ORBITALS 1 One electron cases: E ∝ n2 not dependent on A or mA Two or more electrons: E does depend on n and A (but not mA) Therefore: E2s ≠ E2p E3s ≠ E3p ≠ E3d Value of n determines shell Same n and A means same subshell …and…. same subshell means same energy ORBITAL FILLING SEQUENCE 1 1s 2s 3s 4s 5s 6s 3s 2p 3p 4p 5p 6p 2 3 4 5 3d 4d 4f 5d 5f 6d 6 3p 3d increasing energy Orbital size AND shape effect energy FORCES ACTING ON ORBITAL ELECTRONS Electrons in outer orbitals see nucleus and also the inner electrons Shielding, Screening s and p orbitals have different shapes Therefore, they experience shielding differently s orbital has density at the nucleus p orbital does not s electrons see more of nuclear charge, Z have lower energy, are more stable p electrons see less of Z have higher energy, are less stable ELECTRON CONFIGURATIONS Orbitals are filled by electrons in sequence determined by energy H He Li Be B C N O F Ne Na Mg 1s 1s2 ← filled shell 1s22s ← new row of PT 1s22s2 1s22s22p 1s22s22p2 1s22s22p3 1s22s22p4 1s22s22p5 1s22s22p6 ← filled shell 1s22s22p63s ← new row of PT 1s22s22p63s2 HUND’S RULE When electrons are filling orbitals of equal energy, they go singly into orbitals before starting to double up. Electrons in partially-filled orbitals have the same spin. Carbon 1s2 2s2 2pz 6 electrons 2p2 2px 2py e- repel each other, go into different orbitals Each orbital unique region of space. ELECTRON CONFIGURATIONS Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn 1s22s22p6 ← filled shell (Ne)3s ← new row of PT (Ne)3s2 (Ne)3s23p (Ne)3s23p2 (Ne)3s23p3 (Ne)3s23p4 (Ne)3s23p5 (Ne)3s23p6 ← subshell filled (Ar)4s ← new row of PT (Ar)4s2 (Ar)4s23d (Ar)4s23d2 transition (Ar)4s23d3 metals (Ar)4s13d5 (Ar)4s23d5 (Ar)4s23d6 half or full d orbital 2 7 (Ar)4s 3d more stable than 2 8 (Ar)4s 3d filled s orbital 1 10 (Ar)4s 3d (Ar)4s23d10 PERIODIC TABLE APPRECIATION What first comes to mind when looking at the Periodic Table of the Elements? Marvel at the order and symmetry of the building blocks of our world. We are going to use what we have learned about atomic structure to understand that order. ELECTRON CONFIGURATIONS AND THE PERIODIC TABLE Electron configurations relate to the elements’ location in the periodic table Example: Group 1A Li [He]2s Na [Ne]3s K [Ar]4s Rb [Kr]5s Cs [Xe]6s ELECTRON CONFIGURATIONS AND THE PERIODIC TABLE Alkali metals have ns1 outer shell Li, Na, K, Rb, ... Halogens have np5 outer shell F, Cl, Br, I, ... Noble gases have filled outer shell np6 Ne, Ar, Kr, ... <><><><><><><><><><><><><><><><> Valence electrons determine chemistry Valence electrons are those outside the noble-gas core ION ELECTRON CONFIGURATIONS Metal atoms lose electrons to form cations with charge equal to the group number Example: Mg (2A) Æ Mg2+: [Ne] Nonmetals gain electrons to form anions with charge equal to the group number minus 8. Example: O (6A) Æ O2-: [Ne] Transition metals lose s electrons before d electrons EXAMPLE Fe [Ar]3d64s2 Fe2+ [Ar]3d6 Fe2+ + 2e– [Ar]3d6 Fe3+ + e– [Ar]3d5 ISOELECTRONIC SERIES Isoelectronic: same no. of electrons EXAMPLES O–2 F– Ne Na+ Mg2+ Al3+ 10 electrons each: 1s22s22p6 = [Ne] S2– Cl– Ar K+ Ca2+ 18 electrons each: [Ne]3s23p6 = [Ar] Chapter 6 review 6.25 Chapter 6 review 6.34 Chapter 6 review 6.60 TRENDS IN ATOMIC PROPERTIES: THE PERIODIC TABLE 9 Electron configurations determine organization of the periodic table 9 Next… properties of elements and their periodic behavior 9 Elemental properties determined by: – size (n) and shape (l) of orbitals – atomic number (nuclear charge) Atomic sizes Ionization energies Electron affinities ATOMIC SIZE Size of atom increases going down a group Why? As we go down a group, n increases. As n increases, orbital radius increases. <><><><><><><><><><><><><><><><> Size of atom decreases going from left to right along a period Why? Going across increases no. of protons and the nuclear charge Added outer electrons shielded ineffectively Effective nuclear charge increases, so the electrons are drawn closer ORBITAL SIZE INCREASES WITH n 2s 1s 1s 2s 3s 3s ORBITAL SIZE DECREASES WITH INCREASING Zeff Zeff = Z - S S = core electron screening charge S similar for elements in same period Zeff increases with Z in same period Na: [He]3s1 3s eNe core (10 e-) 2p Z = 11+ 2s lower shells smaller distance screening same shell similar distance little screening PERIODIC TREND IN ATOMIC RADII Fig. 7.6 IONIC SIZE Periodic trends same as for atoms Cation smaller than related atom Na+ Na 97 pm 154 pm why? Na: [Ne]3s1 Anion larger than related atom Cl– Cl 181 pm 99 pm Cl-: [Ne]2s22p6 why? ATOMIC SIZE AND ISOELECTRONIC SERIES Isoelectronic: same no. of electrons EXAMPLES O–2 F– Ne Na+ Mg2+ Al3+ 10 electrons each: 1s22s22p6 = [Ne] nuclear charge increases → size decreases → Ca2+ K+ Ar Cl– S2– what trends in nuclear charge and atomic or ionic size?