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 ( ) values: 0, 1, 2,, (n 1) shape of electron orbital (3) Magnetic quantum number (m ) values:,, 1, 0, 1,, orientation of orbital in space (4) Spin quantum number (m s ) two values: +½ and ½
Symbols used for quantum numbers value 0 1 2 3 subshell name s p d f no. of electrons n = 1 = 0 1s (m = 0) 2 n = 2 = 0 2s (m = 0) 2 n = 2 = 1 2p (m = 1,0,1) 6 n = 3 = 0 3s (m = 0) 2 Table 6.2 n 2 = number of states = number of orbitals in shell n
s orbitals First s orbital: 1s n = 1 = 0 m = 0 Second s orbital: 2s n = 2 = 0 m = 0
SHAPES OF ORBITALS s orbitals Ψ 2 (1s) electron density or probability 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 = 1 same size & shape m = 1, 0, 1 } 3 different orientations First d orbitals: 3d n = 3 = 2 m = 2, 1, 0, 1, 2 5 orbitals 5 orientations
SHAPES OF ORBITALS p orbitals 2 lobes with node between 2p n = 2 = 1 m = 1, 0, 1 3 orbitals Because n same & same, they have same size & shape They differ in orientation p orbitals are directional
SHAPES OF ORBITALS d orbitals 3d n = 3 = 2 m = 2, 1, 0, 1, 2 5 orbitals with different orientations
SHAPES OF ORBITALS 1s sphere dumbbell 2p z 2p x 2p y 3d yz 3dxz 3dxy 3d 2 2 3d 2 x -y z clover-leaf & friend
REVIEW ORBITALS region of space with size, shape, characteristic energy Name Number Shape s 1 spherical p 3 dumbbell d 5 5 shapes f 7 ------ QUANTUM NUMBERS n principal size azimuthal shape m magnetic orientation
FOURTH QUANTUM NUMBER Electron has magnetic moment, as if were spinning Experimental observations confirm m s = ½ or ½ Electrons have 4 quantum numbers n m m s defines orbital Since m s 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 m m s ) Electrons in the same orbital have the same values for the first three quantum numbers ( n m ) m s can have only two values: ½ or ½ Therefore, an orbital can hold only two electrons, and they must have opposite spins Subshell No. of Orbitals Max. No. of e s ( = 0) 1 2 p ( = 1) 3 6 d ( = 2) 5 10 f ( = 3) 7 14
ENERGIES OF ORBITALS One electron cases: E 1 n 2 not dependent on or m Two or more electrons: E does depend on n and (but not m ) Therefore: E 2s E 2p E 3s E 3p E 3d Value of n determines shell Same n and means same subshell and. same subshell means same energy
ORBITAL FILLING SEQUENCE 1 2 1s 3 4 2s 2p 5 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6 3s 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 1s He 1s 2 filled shell Li 1s 2 2s new row of PT Be 1s 2 2s 2 B 1s 2 2s 2 2p C 1s 2 2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s 2 2s 2 2p 6 filled shell Na 1s 2 2s 2 2p 6 3s new row of PT Mg 1s 2 2s 2 2p 6 3s 2
HUND S 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 6 electrons 1s 2 2s 2 2p 2 2p z 2p x 2p y e - repel each other, go into different orbitals Each orbital unique region of space.
ELECTRON CONFIGURATIONS Ne 1s 2 2s 2 2p 6 filled shell Na (Ne)3s new row of PT Mg (Ne)3s 2 Al (Ne)3s 2 3p Si (Ne)3s 2 3p 2 P (Ne)3s 2 3p 3 S (Ne)3s 2 3p 4 Cl (Ne)3s 2 3p 5 Ar (Ne)3s 2 3p 6 subshell filled K (Ar)4s new row of PT Ca (Ar)4s 2 Sc (Ar)4s 2 3d Ti (Ar)4s 2 3d 2 V (Ar)4s 2 3d 3 transition Cr (Ar)4s 1 3d 5 metals Mn (Ar)4s 2 3d 5 Fe (Ar)4s 2 3d 6 half or full d orbital Co (Ar)4s 2 3d 7 more stable than Ni (Ar)4s 2 3d 8 filled s orbital Cu (Ar)4s 1 3d 10 Zn (Ar)4s 2 3d 10
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 ns 1 outer shell Li, Na, K, Rb,... Halogens have np 5 outer shell F, Cl, Br, I,... Noble gases have filled outer shell np 6 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) Mg 2+ : [Ne] Nonmetals gain electrons to form anions with charge equal to the group number minus 8. Example: O (6A) O 2- : [Ne] Transition metals lose s electrons before d electrons EXAMPLE Fe Fe 2+ + 2e [Ar]3d 6 4s 2 [Ar]3d 6 Fe 2+ Fe 3+ + e [Ar]3d 6 [Ar]3d 5
ISOELECTRONIC SERIES Isoelectronic: same no. of electrons EXAMPLES O 2 F Ne Na + Mg 2+ Al 3+ 10 electrons each: 1s 2 2s 2 2p 6 = [Ne] S 2 Cl Ar K + Ca 2+ 18 electrons each: [Ne]3s 2 3p 6 = [Ar]
6.25 Chapter 6 review
6.34 Chapter 6 review
6.60 Chapter 6 review
TRENDS IN ATOMIC PROPERTIES: THE PERIODIC TABLE Electron configurations determine organization of the periodic table Next properties of elements and their periodic behavior 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 1s 2s 3s 1s 2s 3s
ORBITAL SIZE DECREASES WITH INCREASING Z eff Z eff = Z - S S = core electron screening charge S similar for elements in same period Z eff increases with Z in same period Na: [He]3s 1 3s e - Ne 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]3s 1 Anion larger than related atom Cl Cl 181 pm 99 pm why? Cl - : [Ne]2s 2 2p 6
ATOMIC SIZE AND ISOELECTRONIC SERIES Isoelectronic: same no. of electrons EXAMPLES O 2 F Ne Na + Mg 2+ Al 3+ 10 electrons each: 1s 2 2s 2 2p 6 = [Ne] nuclear charge increases size decreases Ca 2+ K + Ar Cl S 2 what trends in nuclear charge and atomic or ionic size?