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1 RVCC Fall 2008 CHEM 103-1&2 – General Chemistry I Chapter 7: Electron Configuration and the Periodic Table Chemistry: The Molecular Science, 3 rd Ed.

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Presentation on theme: "1 RVCC Fall 2008 CHEM 103-1&2 – General Chemistry I Chapter 7: Electron Configuration and the Periodic Table Chemistry: The Molecular Science, 3 rd Ed."— Presentation transcript:

1 1 RVCC Fall 2008 CHEM 103-1&2 – General Chemistry I Chapter 7: Electron Configuration and the Periodic Table Chemistry: The Molecular Science, 3 rd Ed. by Moore, Stanitski, and Jurs

2 2 Electromagnetic Radiation Theories about the arrangement and energy of electrons in atoms are based on experimental studies of the interaction of matter with electromagnetic radiation. Electromagnetic radiation - consists of oscillating perpendicular electric and magnetic fields that travel through space with the same speed – 3.00 ×10 8 m/s or 186,000mi/s. Light, microwaves, x-rays, and TV and radio transmissions are different kinds of electromagnetic waves Electric field Magnetic field

3 3 Electromagnetic Spectrum energy (and frequency) decrease wavelength increases limit your exposure to X-rays (high energy) radio waves are harmless (low energy)

4 4 Electromagnetic Radiation - waves The wavelength,  lambda), is the distance between any two adjacent identical points of a wave (m, cm, nm). The frequency,  (nu), of a wave is the number of wavelengths that pass a fixed point in one second (Hz, s- 1, or cycles/s). The amplitude, A, is the height of the crest or the intensity of the radiation. The phase, φ, is the relative position of the wave.

5 5 Electromagnetic Radiation The frequency and the wavelength are inversely related to each other. The constant (c) is the speed of light.  c/ λ cycles/s = m/s / m/cycle c = 3.00 × 10 8 m/s (speed of light in a vacuum is a constant!) Both waves are traveling the same speed.

6 6 Example 1 Calculate the frequency of X-ray that has a wavelength of 8.21 nm.

7 7 Continuous Spectrum When a white light from an incandescent lamp passes through a narrow slit and then through a prism it gives a continuous spectrum (containing light of all wavelengths or frequencies). Prism CD / Diffraction glasses /slit

8 8 Planck’s Quantum Theory A heated object gives off radiation of shorter wavelengths with increasing temperatures. (Classical physics cannot explain.) Plank’s Theory: Matter emits radiation in ‘packets’ that have a minimum (threshold) energy. E quantum = h ν radiation Planck’s constant, h = × Js shorter wavelength higher energy Increasing filament T

9 9 The Photoelectric Effect current (# of ejected e - ) increasing E No current until minimum frequency or energy obtained. Increasing the intensity (amplitude) of the light, only increases the current.

10 10 The Photoelectric Effect E = hν The only energies an atom can have are 1h, 2h, 3h … Einstein postulated that light consists of photons (a particle!) of electromagnetic energy.

11 11 The “wave” and “particle” pictures of light should be regarded as complementary views of the same physical entity. Dual Nature of Light The equation E = h displays this duality; h says that energy only occurs in discrete steps or quanta as a “packets” or “particle” and is the frequency of the associated “wave.” There is experimental evidence for both: wave – diffraction, refraction… particle - photoelectric effect, line spectra…

12 12 Dual Nature of Light - particle What is the energy of one quantum of laser light that has a frequency of 4.57 × s -1 ? E = h = (6.626 × Js ) (4.57 × s -1 ) = 3.03 × J

13 13 Dual Nature of Light - wave What is the wavelength of one quantum of laser light that has an energy of 3.03 × J? E = h = 6.55 × m = 655 nm

14 Photons of light with sufficient energy can eject electrons from a gold surface. To do so requires photons with an energy equivalent to a 257nm wavelength. Will photons in the visible region ( nm) of the spectrum dislodge electrons from a gold surface? NO CALCULATION NECESSARY! 32. To eject electrons from the surface of potassium metal requires a minimum energy of 3.69 x J. When 600nm photons shine on a potassium surface, will they cause the photoelectric effect? = (6.626 × Js ) (2.998 × 10 8 ms -1 ) 600nm ·1m/10 -9 nm = 3.31 × J Example 2 (Ch. 7, #31, #32) Photoelectric Effect

15 15 Atomic Line Spectra The light emitted by a heated gas, such as hydrogen, results in a line spectrum-a spectrum showing only specific wavelengths of light (only certain energy values). Hydrogen atoms become energized they give off energy in the form of light.

16 16 Heated solid objects or molecules emit continuous spectra. Excited atomic gases emit line spectra. Each element has a unique pattern. Atomic Line Emission Spectra

17 17 The Bohr Model of the Hydrogen Atom In 1913, Niels Bohr using the work of Einstein and Planck, offered an explanation of the line spectra and applied new quantum theory to the simplest atom, hydrogen. Electrons can only occupy discrete positions about the nucleus. An ‘ah-ha!’ moment – quanta of light and structure of atom connected.

18 18 An electron can only inhabit specific energy levels in an atom. The levels represent orbits of increasing distance (n) from the nucleus. As n increases, the potential energy increases. E potential = -R H /n 2 LAB! The Bohr Model of the Hydrogen Atom

19 19 The Bohr Model of the Hydrogen Atom Electrons are typically in their ‘ground state’ or lowest energy configuration. When energy is absorbed, an electron moves from a low energy level to a higher energy level (n i =2 to n f =3,4,5,6).  E=+ This is called an ‘excited state’. Energy is emitted when an electron moves from a higher energy level back to a lower energy level (i.e. n=3 to n=2).  E=-

20 20 Example Calculate the wavelength of light that will be emitted when an electron in the hydrogen atom moves from… n = 6 (E = × J) to n = 2 (E = × J) Guess what?… the Bohr model makes accurate predictions only for a one electron atom (Hydrogen).

21 21 Schrödinger’s Equations (1926): Treats an electron as a standing wave. (If light can be a particle why can’t an electron be a wave?!) The mathematical description of the behavior of electrons as waves is called “wave mechanics” or “quantum mechanics”. wave functions Mathematical functions (wave functions,  ) predict the allowed energy states of an electron and the probability of finding that electron in a given region of space. Beyond the Bohr Model: Quantum Mechanics

22 22 Energy Levels and Orbitals We can no longer think of an electron as having a precise orbit in an atom. We can obtain the probability of finding an electron of a given energy and momentum at a given point around the nucleus. Orbital is a region in space around the nucleus where there is a high (90 %)probability of finding an electron The main differences between orbitals and Bohr’s orbits are: - orbital is three-dimesional – we talk about electron clouds - we can only determine the probability of finding an electron in an orbital.

23 23 An electron density (probability) map plots  2 for each point in space. Quantum Mechanical Model The H-atom ground- state orbital An orbital depicts the space where an electron is most likely (99% of time) located. The quantum number n represents the most probable distance of the electron form the nucleus.

24 24 Energy Levels and Orbitals A collection of orbitals within the same probable distance from the nucleus is called an electron shell or energy level Each shell has one or more subshells within it. Each subshell has one or more orbitals within it.

25 25 Quantum Numbers There are four quantum numbers in the wave equation that describes the position of an electron. n = principal quantum number (the orbitals distance from the nucleus) l = angular momentum quantum number (the shape of the orbital) m l = magnetic quantum number (the orientation of the orbital) m s = spin quantum number (limits two electrons per orbital) Every electron has a unique set (n, l, m l, m s ) of quantum numbers.

26 26 Quantum Numbers The principle quantum number, n, is a positive integer that indicates the most probable distance of an orbital from the nucleus. Rule: n = 1, 2, 3, … n=4 n=3 n=2 n=1 E pot.

27 27 Quantum Numbers The angular momentum quantum number, l, represents subshells in the principal levels, and defines the shape of the orbital. l = 0123 orbital type = spdf “stupid pigeons don’t fly” RULE:l = 0, 1, … n-1 If n=1, l=0 (one subshell) If n=2, l=0 or 1 (two subshells) If n=3, l=0 or 1 or 2 (three subshells) If n=4, l=0 or 1 or 2 or 3 (four subshells)

28 28 Quantum Numbers (s) (p) (s) (p) (d) (s) (p) (d) (f) Principle Shell Subshell(s)

29 29 Quantum Numbers The magnetic quantum number, m l defines the orientation of an orbital in the space around the nucleus of an atom. RULE: m l = -l …0... +l If l=0, (s subshell), m l = 0 (one orientation) If l=1 (p subshell), m l = -1, 0, 1 (three orientations) If l=2 (d subshell), m l = -2, -1, 0, 1, 2 (five orientations) If l=3 (f subshell), m l = -3, -2, -1, 0, 1, 2, 3 (seven orientations)

30 30 The s orbitals (spherical shape) Node l = 0 (s orbital) m l = 0 (one orientation) Node – a region in space where the probability of finding an electron is zero.

31 31 The p orbitals (dumbbell shape) Node l = 1 (p orbital) m l = -1, 0, 1 (three orientations)

32 32 The d orbitals (cloverleaf shape) If l = 2 (d orbital) m l = -2, -1, 0, 1, 2 (five orientations)

33 33 Quantum Numbers The spin magnetic quantum number, m s m s = + 1/2, - 1/2. defines the two possible spin orientations. Each orbital (s, p x, d xz …) can hold two electrons, preferred if occupying the same orbital (i.e. p z, ) or preferred if occupying different orbitals at the same energy or sublevel(p x and p y or d x2y2 and d z2 )

34 34 Quantum Numbers - Summary (s) (p) (s) (p) (d) (s) (p) (d) (f) =n distance shapeorientation to nucleus

35 35 Trends in Quantum Numbering n th shell has n subshells total # of orbitals in the n th shell is n 2 The number of orbitals in each subshell equals (2l +1)

36 36 Practice How many subshells are there in the electron shell with the principal quantum number n=4? 4 subshells If n=4, l=0, 1, 2, or 3 How many orbitals are in the n=4 shell? n 2 = 16 orbitals (1 in l=0, 3 in l=1, 5 in l=2, and 7 in l=4)

37 37 Practice Which of the following combination of quantum numbers is allowed? If not, why? 1. n = 2, l = 1, m l = 2, m s =+½ 2. n = 3, l = 2, m l = 0, m s =-½ 3. n = 1, l = 0, m l = 0, m s =1 4. n = 3, l = 3, m l = 2, m s =-½ 5. n = 2, l = 0, m l = 0, m s =+½

38 38 Houses on a Hill – the e- neighborhood about the nucleus bed m s =+½ m s =-½ n=1 l=0 (s) m l =0 n=2 l=1 (p) m l =-1,0,1 n=2 l=0 (s) m l = 0 House floor room n=3 l=2 (d) m l =-2,-1,0,1,2 n=3 l=1 (p) m l =-1,0,1 n=3 l=0 (s) m l = 0 n=4 l=2 m l =-1 m s =+½ nucleus

39 39 Electron Configurations Before we start putting electrons in their houses (ground state configurations), there are some rules….

40 40 No two electrons in an atom may have the same set of four quantum numbers. (n, l, m l, m s ) Or…only two electrons can occupy an orbital and they must have opposite spins. Pauli’s Exclusion Principle

41 41 Aufbau Principle Electrons fill the orbitals in order, from lowest energy to highest. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, … n=2, l=0, m l =0 Principle Shell order: 1s<2s<3s… Subshell order: 3s<3p<3d p x = p y = p z

42 42 Electron Configuration and the Periodic Table s block main group elements d block transition elements p block main group elements 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, …

43 43 Hund’s Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in a particular set of degenerate (same energy) orbitals. Orbital Notation 1s 2s 2p or C: 1s 2 2s 2 2p 2

44 44 Electron Configuration of Main Group Elements H1s 1 1s He1s 2 condensed notation, orbital notation

45 45 Electron Configuration of Main Group Elements Li1s 2 2s 1 2s 1s Be1s 2 2s 2 1s 2s core electrons (inner shells) valence electrons (outer shells, those in highest principal quantum number, n.)

46 46 B1s 2 2s 2 2p 1 C1s 2 2s 2 2p 2 1s2s 1s2s 2p x 2p y 2p z 2p x 2p y 2p z N1s 2 2s 2 2p 3 1s2s2p x 2p y 2p z list the 4 quantum numbers for each pair of electrons list the 4 quantum numbers for each pair of electrons list the 4 quantum numbers for each pair of electrons

47 47 O1s 2 2s 2 2p 4 (=8) 1s2s2p x 2p y 2p z F1s 2 2s 2 2p 5 (=9) 1s2s2p x 2p y 2p z Ne1s 2 2s 2 2p 6 (=10) 1s2s2p x 2p y 2p z Neon has a closed shell configuration. Neon is stable and inert.

48 48 Noble Gas Notation P 1s 2 2s 2 2p 6 3s 2 3p 3 P[Ne]3s 2 3p 3 Ne

49 49 Table 7-5, p.295 Electron Configurations of the First Ten Elements

50 50 e e e e e e e e Electron Configurations of next eight Elements DO NOW: Ca (Appendix D)

51 51 Block identities show where successive e - add. Note: d “steps down”, f “steps down” again. Atom Electron Configurations Main group s block 2 s 4 s 5 s 6 s 7 s 3 s 1s1s 5 f 4 f 6d6d 4d4d 3d3d 5d5d 6d6d 5d5d 4d4d 3d3d 4p 5p 6p 7p 3p 2p 1s1s Lanthanides and actinides f block Transition elements d block Main group p block 57 La 89 Ac 20 Ca Aufbau Diagram and the Periodic Table

52 52 Electron Configuration of Transition Elements: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 or [Ar] 4s 2 3d 1 4s3d xy 3d yz 3d xz (n – 1)d orbitals are filled after ns orbitals and before np prbitals Z = 21 Sc

53 53 Electron Configuration of Transition Elements 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 or [Ar] 4s 2 3d 2 4s3d xy 3d yz 3d xz Hund’s Rule Z = 22 Ti

54 54 Cr [Ar] 4s 1 3d 5 4s3d xy 3d yz 3d xz From spectroscopy we know that: Half-filled d subshell, and half-filled s, is more energetically stable. Cr [Ar] 4s 2 3d 4 4s3d xy 3d yz 3d xz Expected:

55 55 Electron Configuration of Transition Elements 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 or [Ar] 4s 2 3d 6 4s3d xy 3d yz 3d xz Z = 26 Fe

56 56 Cu [Ar] 4s 1 3d 10 4s3d xy 3d yz 3d xz Filled d subshell, and half-filled s, is more energetically stable. Cu [Ar] 4s 2 3d 9 4s3d xy 3d yz 3d xz

57 57 Sc 3d 1 4s 2 Y 4d 1 5s 2 La 5d 1 6s 2 Ti 3d 2 4s 2 Zr 4d 2 5s 2 Hf 5d 2 6s 2 V 3d 3 4s 2 Nb 4d 4 5s 1 Ta 5d 3 6s 2 Cr 3d 5 3d 5 4s 1 Mo 4d 5 4d 5 5s 1 W 5d 4 6s 2 Mn 3d 5 3d 5 4s 2 Tc 4d 5 4d 5 5s 2 Re 5d 5 5d 5 6s 2 Fe 3d 6 4s 2 Ru 4d 7 5s 1 Os 5d 6 6s 2 Co 3d 7 4s 2 Rh 4d 8 5s 1 Ir 5d 7 6s 2 Ni 3d 8 4s 2 Pd 4d 10 Pt 5d 9 6s 1 Cu 3d 10 3d 10 4s 1 Ag 4d 10 4d 10 5s 1 Au 5d 10 5d 10 4s 1 Zn 3d 10 3d 10 4s 2 Cd 4d 10 4d 10 5s 2 Hg 5d 10 5d 10 6s 2 ½ filledfilled Note: ½ filled and filled shells have extra stability Electron Configurations of Transition Metals

58 58 Electron Configurations 4s 2 3d 10 4p 5 4s3d xy 3d yz 3d xz 4p x 4p y 4p z Br [Ar]

59 59 Electron Configurations After Z=57, the f-block fills 4f 5f

60 60 Valence electrons reside in the outermost shell of an atom. For the main group atoms, this includes ns+np electrons. Where n is the highest principal number. For transition metals, this includes ns+np+(n-1)d electrons. Valence electrons –Valence electrons are primarily involved in chemical reactions. –Elements within a given group have the same “valence shell configuration.” –This accounts for the similarity of the chemical properties among groups of elements. F [He]2s 2 2p 5

61 61 Valence electrons Electron configuration of halogens – group 7A F [He]2s 2 2p 5 Cl [Ne]3s 2 3p 5 Br [Ar]4s 2 3d 10 4p 5 I [Kr] 5s 2 4d 10 5p 5 Each halogen has 7 valence electrons Filled d-orbitals are NOT valence e-.

62 62 Lewis Electron-Dot Symbols A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element. Na..... Si... : P :. :. S Mg... Al.. : : Cl. : Ar : : : : Group IGroup IIGroup VIIGroup VIII Group VI Group IVGroup VGroup III Note that the group number indicates the number of valence electrons. Na..... Si... : P :. :. S Mg... Al.. : : Cl. : Ar : : : : Na..... Si... : P :. :. S Mg... Al.. : : Cl. : Ar : : : : [Ne]3s 1 [Ne]3s 2 3p 1 [Ne]3s 2 3p 3 [Ne]3s 2 3p 5 [Ne]3s 2 [Ne]3s 2 3p 2 [Ne]3s 2 3p 4 [Ne]3s 2 3p 6

63 63 Ions and Electron Configuration Group VIII Na..... Si... : P :. :. S Mg... Al.. : : Cl. : Ar : : : : Group IGroup IIGroup VIIGroup VIII Group VI Group IVGroup VGroup III Na..... Si... : P :. :. S Mg... Al.. : : Cl. : Ar : : : : Na..... Si... : P :. :. S Mg... Al.. : : Cl. : Ar : : : : Na..... Si... : P :. :. S Mg... Al.. : : Cl. : Ar : : : : Na..... Si... : P :. :. S Mg... Al.. : : Cl. : Ar : : : : Ne : : : : gain electrons to be like [Ar] anions (-) lose electrons to be like [Ne] cations (+) Chapter –Valence electrons are primarily involved in chemical reactions. XXXX X X

64 Isoelectronic = same number and configuration of electrons. The most stable ion is isoelectronic with the nearest noble gas. [Ne] O -2 F -1 Na +1 Mg +2 Al +3 [Ar] S -2 Cl -1 K +1 Ca Ion Electron Configurations

65 65 Ion Electron Configurations Negative ion: add one e - for each “-” Positive ion: remove one e - for each “+” (16 + 2) = 18 e - S[Ne] 3s 2 3p 4 [Ar] S 2- [Ne] 3s 2 3p 6 or[Ar] (13 - 3) = 10 e - Al [Ne] 3s 2 3p 1 [Ne] Al +3 [Ne] S 2- Al +3

66 66 Electron Configuration of Ions Cations: ` Li 1s 2 2s 1 [He]2s 1 Mg 1s 2 2s 2 2p 6 3s 2 [Ne]3s 2 Li + 1s 2 [He] Mg 2+ 1s 2 2s 2 2p 6 [Ne] Everybody wants to be a noble gas…or, at least, isoelectronic with one. X X X X - 1e - - 2e -

67 67 Electron configuration of ions Anions: O 1s 2 2s 2 2p 4 [He]2s 2 2p 4 F 1s 2 2s 2 2p 5 [He]2s 2 2p 5 P 1s 2 2s 2 2p 6 3s 2 3p 3 [Ne]3s 2 3p 3 O 2- 1s 2 2s 2 2p 6 [Ne] F - 1s 2 2s 2 2p 6 [Ne] P 3- 1s 2 2s 2 2p 6 3s 2 3p 6 [Ar] +2e - +1e - +3e -

68 68 Transition Metal Ions (n-1)dadded last (n-1)d e- are added last, BUT… nslost first BUT… ns e - are lost first. → Fe 2+ [Ar] 3d 6 Mn [Ar] 4s 2 3d 5 →Mn 2+ [Ar] 3d 5 → Fe 3+ [Ar] 3d 5 → Mn 4+ [Ar] 3d 3 → Mn 7+ [Ar] Fe [Ar] 4s 2 3d 6

69 69 Practice Write the electron configuration for Te. Write the electron configuration for Ag and Ag +1. Write the electron configuration for the chlorine ion. Write the electron configuration for Ba 2+. Write the electron configuration for Os. PPUnit11.ppt


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