1 © 2009 Brooks/Cole - Cengage Chapter 6 Chem 1211 Class 14 Atomic Structure

2 © 2009 Brooks/Cole - Cengage Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS ( l ) ORBITALS (m l )

3 © 2009 Brooks/Cole - Cengage QUANTUM NUMBERS There is a hierarchy in quantum numbers: l n m l l A few different l can exist for the same n and a few different m l can exist for the same l nn determines the energy of the orbital l m lfor hydrogen atom the energy of the orbital does not depend on l and m l ll determines shape of the orbital m l lm l determines position of the orbital relative to others with the same l

4 © 2009 Brooks/Cole - Cengage Azimuthal Quantum Number ll determines the shape of the orbital for given n, l can be any integer from 0 to (n-1) every l has its proper (letter) name: orbital with l = 0 called s-orbital orbital with l = 1 called p-orbital orbital with l = 2 called d-orbital orbital with l = 3 called f-orbital

5 © 2009 Brooks/Cole - Cengage Magnetic Quantum Number m l l lm l determines position of the orbital relative to others with the same l and how many orbitals with the same l can exist. m l l to l, l, (2 l +1) m lm l is an integer from - l to l, so for given l, there are (2 l +1) orbitals with different m l How many? l =0 s-orbital1 l =1 p-orbital3 l =2 d-orbital5 l =3 f-orbital7

6 © 2009 Brooks/Cole - Cengage Types of Atomic Orbitals See Active Figure p x n, principal l, azimuthal m l, magnetic

7 © 2009 Brooks/Cole - Cengage Each orbital can be occupied by no more than 2 electrons! Because: No two electrons in atom can have the same set of all quantum numbers ( ) No two electrons in atom can have the same set of all quantum numbers (Pauli’s Exclusion Principle)And: There is a 4th quantum number, the electron spin quantum number, m s. Arrangement of Electrons in Atoms: restriction on number of residents! Wolfgang Pauli ( )

8 © 2009 Brooks/Cole - Cengage Electron Spin Quantum Number, m s Can be proven experimentally (Stern, Gerlach, 1926) that electron has a magnetic moment (acts as a small magnet).Can be proven experimentally (Stern, Gerlach, 1926) that electron has a magnetic moment (acts as a small magnet). It referred to as “spin.”It referred to as “spin.” Spin is quantized: it can be only along magnetic field or against magnetic field.Spin is quantized: it can be only along magnetic field or against magnetic field. Two spin directions are given byTwo spin directions are given by m s = +1/2 and -1/2. Can be proven experimentally (Stern, Gerlach, 1926) that electron has a magnetic moment (acts as a small magnet).Can be proven experimentally (Stern, Gerlach, 1926) that electron has a magnetic moment (acts as a small magnet). It referred to as “spin.”It referred to as “spin.” Spin is quantized: it can be only along magnetic field or against magnetic field.Spin is quantized: it can be only along magnetic field or against magnetic field. Two spin directions are given byTwo spin directions are given by m s = +1/2 and -1/2.

9 © 2009 Brooks/Cole - Cengage Electron Spin Quantum Number, m s If two electrons in atom have the opposite spins, the total spin is zero (spins are “paired”)If two electrons in atom have the opposite spins, the total spin is zero (spins are “paired”) Electron spin is responsible for magnetic properties of the materialsElectron spin is responsible for magnetic properties of the materials Materials containing unpaired spins are either paramagnetic or ferromagneticMaterials containing unpaired spins are either paramagnetic or ferromagnetic Materials containing only paired spins are diamagneticMaterials containing only paired spins are diamagnetic If two electrons in atom have the opposite spins, the total spin is zero (spins are “paired”)If two electrons in atom have the opposite spins, the total spin is zero (spins are “paired”) Electron spin is responsible for magnetic properties of the materialsElectron spin is responsible for magnetic properties of the materials Materials containing unpaired spins are either paramagnetic or ferromagneticMaterials containing unpaired spins are either paramagnetic or ferromagnetic Materials containing only paired spins are diamagneticMaterials containing only paired spins are diamagnetic

10 © 2009 Brooks/Cole - Cengage Electron Spin and Magnetism Diamagnetic : NOT attracted to a magnetic field (slightly repelled)Diamagnetic : NOT attracted to a magnetic field (slightly repelled) Paramagnetic : substance is attracted to a magnetic field.Paramagnetic : substance is attracted to a magnetic field. Ferromagnetic: substance is strongly attracted to a magnetic field.Ferromagnetic: substance is strongly attracted to a magnetic field. Substances with unpaired electrons are paramagnetic or ferromagnetic.Substances with unpaired electrons are paramagnetic or ferromagnetic.

11 © 2009 Brooks/Cole - Cengage Measuring Paramagnetism Paramagnetic : substance is attracted to a magnetic field. Substance has unpaired electrons. Diamagnetic : NOT attracted to a magnetic field See Active Figure 6.18

12 © 2009 Brooks/Cole - Cengage ferromagneticdiamagnetic paramagnetic magnetic field

13 © 2009 Brooks/Cole - Cengage n → shell1, 2, 3, 4,... l → subshell0, 1, 2,... n - 1 m l → orbital - l l m s → electron spin+1/2 and -1/2 n → shell1, 2, 3, 4,... l → subshell0, 1, 2,... n - 1 m l → orbital - l l m s → electron spin+1/2 and -1/2 QUANTUM NUMBERS Now there are four!

14 © 2009 Brooks/Cole - Cengage Magnetic Resonance Imaging (MRI) Protons and neutrons are also small magnets, means, they also have spin; Proton is the nucleus of hydrogen, energy transitions of its spin in strong magnetic field can be measured by Nuclear Magnetic Resonance (NMR); Concentration of the protons (i.e. water) in different parts of the sample can be measured by NMR - that is called Magnetic Resonance Imaging (MRI) It is powerful diagnostic tool, because living organism mostly consists of water!

15 © 2009 Brooks/Cole - Cengage Magnetic Resonance Imaging (MRI) X-Ray MRI

16 © 2009 Brooks/Cole - Cengage Chapter 7 Chem 1211 Atomic Electron Configurations and Periodic Table

17 © 2009 Brooks/Cole - Cengage Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS ( l ) ORBITALS (m l )

18 © 2009 Brooks/Cole - Cengage Each orbital can be assigned no more than 2 electrons! This is tied to the existence of a 4th quantum number, the electron spin quantum number, m s and… Arrangement of Electrons in Atoms

19 © 2009 Brooks/Cole - Cengage Pauli Exclusion Principle No two electrons in the same atom can have the same set of 4 quantum numbers. That is, each electron has a unique address.

20 © 2009 Brooks/Cole - Cengage Electrons in Atoms When n = 1, then l = 0 this shell has a single orbital (1s) to which two e- can be assigned. this shell has a single orbital (1s) to which two e- can be assigned. When n = 2, then l = 0, 1 2s orbital 2e- 2s orbital 2e- three 2p orbitals6e- three 2p orbitals6e- TOTAL = 8e- first shell second shell

21 © 2009 Brooks/Cole - Cengage Electrons in Atoms When n = 3, then l = 0, 1, 2 3s orbital 2e- 3s orbital 2e- three 3p orbitals6e- three 3p orbitals6e- five 3d orbitals10e- TOTAL = 18e- When n = 3, then l = 0, 1, 2 3s orbital 2e- 3s orbital 2e- three 3p orbitals6e- three 3p orbitals6e- five 3d orbitals10e- TOTAL = 18e- third shell

22 © 2009 Brooks/Cole - Cengage Electrons in Atoms When n = 4, then l = 0, 1, 2, 3 4s orbital 2e- 4s orbital 2e- three 4p orbitals6e- three 4p orbitals6e- five 4d orbitals10e- seven 4f orbitals14e- seven 4f orbitals14e- TOTAL = 32e- And many more! In general, electron capacity of n-th shell is 2n 2

23 © 2009 Brooks/Cole - Cengage ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY Length of period corresponds to shell capacity only for the 1 st two periods. WHY?

24 © 2009 Brooks/Cole - Cengage Assigning Electrons to Atoms Electrons generally assigned to orbitals of successively higher energy.Electrons generally assigned to orbitals of successively higher energy. For H atoms, E = - C(1/n 2 ). E depends only on n.For H atoms, E = - C(1/n 2 ). E depends only on n. For many-electron atoms, energy depends on both n and l.For many-electron atoms, energy depends on both n and l. See Active Figure 7.1 and Figure 7.2 See Active Figure 7.1 and Figure 7.2 Electrons generally assigned to orbitals of successively higher energy.Electrons generally assigned to orbitals of successively higher energy. For H atoms, E = - C(1/n 2 ). E depends only on n.For H atoms, E = - C(1/n 2 ). E depends only on n. For many-electron atoms, energy depends on both n and l.For many-electron atoms, energy depends on both n and l. See Active Figure 7.1 and Figure 7.2 See Active Figure 7.1 and Figure 7.2

25 © 2009 Brooks/Cole - Cengage Assigning Electrons to Subshells In H atom all subshells of same n have same energy.In H atom all subshells of same n have same energy. In many-electron atom:In many-electron atom: a) subshells increase in energy as value of n + l increases. b) for subshells of same n + l, subshell with lower n is lower in energy.

26 © 2009 Brooks/Cole - Cengage Electron Filling Order See Figure 7.2 Aufbau scheme: electrons occupy the position with the lowest possible energy Why is this order?

27 © 2009 Brooks/Cole - Cengage Effective Nuclear Charge, Z* Z* is the nuclear charge experienced by the outermost electrons. See Figure 7.3Z* is the nuclear charge experienced by the outermost electrons. See Figure 7.3 Explains why E(2s) < E(2p)Explains why E(2s) < E(2p) Z* increases across a period owing to incomplete shielding by inner electrons.Z* increases across a period owing to incomplete shielding by inner electrons. Estimate Z* = [ Z - (no. inner electrons) ]Estimate Z* = [ Z - (no. inner electrons) ] Charge felt by 2s e- in Li Z* = = 1Charge felt by 2s e- in Li Z* = = 1 Be Z* = = 2Be Z* = = 2 B Z* = = 3and so on!B Z* = = 3and so on!

28 © 2009 Brooks/Cole - Cengage Effective Nuclear Charge See Figure 7.3 Electron cloud for 1s electrons Z* is the nuclear charge experienced by the outermost electrons.

29 © 2009 Brooks/Cole - Cengage Effective Nuclear Charge See Figure 7.3 Z* is the nuclear charge experienced by the outermost electrons. probability distribution for 2p electron Shielding by inner electrons: s < p < d Effective charge for s > p > d Correspondingly, energy: s < p < d

30 © 2009 Brooks/Cole - Cengage Electron Filling Order See Figure 7.2 Aufbau scheme: electrons occupy the position with the lowest possible energy Now we understand the order!

31 © 2009 Brooks/Cole - Cengage Writing Atomic Electron Configurations 1 1 s value of n value of l no. of electrons spdf notation for H, atomic number = 1 Two ways of writing configs. One is called the spdf notation.

32 © 2009 Brooks/Cole - Cengage Writing Atomic Electron Configurations Two ways of writing configs. Other is called the orbital box notation. One electron has n = 1, l = 0, m l = 0, m s = + 1/2 Other electron has n = 1, l = 0, m l = 0, m s = - 1/2

33 © 2009 Brooks/Cole - Cengage See “Toolbox” in ChemNow for Electron Configuration tool.

34 © 2009 Brooks/Cole - Cengage Electron Configurations and the Periodic Table See Active Figure 7.4 characteristic elements

35 © 2009 Brooks/Cole - Cengage LithiumLithium Group 1A Atomic number = 3 1s 2 2s 1 → 3 total electrons

36 © 2009 Brooks/Cole - Cengage BerylliumBeryllium Group 2A Atomic number = 4 1s 2 2s 2 → 4 total electrons

37 © 2009 Brooks/Cole - Cengage BoronBoron Group 3A Atomic number = 5 1s 2 2s 2 2p 1 → 5 total electrons 5 total electrons

38 © 2009 Brooks/Cole - Cengage CarbonCarbon Group 4A Atomic number = 6 1s 2 2s 2 2p 2 → 6 total electrons 6 total electrons Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible (maximal spin configuration has the lowest energy).

39 © 2009 Brooks/Cole - Cengage NitrogenNitrogen Group 5A Atomic number = 7 1s 2 2s 2 2p 3 → 7 total electrons 7 total electrons

40 © 2009 Brooks/Cole - Cengage OxygenOxygen Group 6A Atomic number = 8 1s 2 2s 2 2p 4 → 8 total electrons 8 total electrons

41 © 2009 Brooks/Cole - Cengage FluorineFluorine Group 7A Atomic number = 9 1s 2 2s 2 2p 5 → 9 total electrons 9 total electrons

42 © 2009 Brooks/Cole - Cengage NeonNeon Group 8A Atomic number = 10 1s 2 2s 2 2p 6 → 10 total electrons 10 total electrons Note that we have reached the end of the 2nd period, and the 2nd shell is full!