# Chapter 11 Atoms, Energy and Electron Configurations Objectives

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1 Objectives 1. To review Rutherford s model of the atom 2. To explore the nature of electromagnetic radiation 3. To see how atoms emit light

2 A. Rutherford s Atom.but there is a problem here!!

3 Using Rutherford s Model Atoms Should Collapse...then how are they so stable?

4 When were you last exposed to electromagnetic radiation?

5 B. Energy and Light What is the nature of electromagnetic radiation? How are the types of electromagnetic radiation different? Light can be modeled as a wave, like a wave in water

6 B. Energy and Light How are the types of light different? (m) ν (s -1 ) c = ν.λ Wavelength, λ (Greek letter lambda ) units - m Frequency, ν (Greek letter nu ) units - s -1 or Hertz (Hz) Amplitude, peak height relates to energy in wave Speed, c : velocity of light in a vacuum is 3x10 8 m.s -1

7 B. Energy and Light Electromagnetic radiation

8 Wavelength and Frequency of Visible Light Frequency in Terahertz (1THz = Hz) and Wavelength in Nanometers (1nm = 10-9 meters)

9 Light Sometimes Behaves in Un-wavelike Ways The Photoelectric Effect Light shining on a metal surface can cause electrons to be separated from their atoms Below a threshold frequency no electrons are emitted however high the intensity At the threshold frequency electrons start to be emitted At higher frequencies electrons have additional kinetic energy Photoelectric Effect Demo

10 B. Energy and Light Dual nature of light Two co-existing models Wave Photon packet of energy

11 B. Energy and Light Different photons (from light of different wavelengths) carry different amounts of energy. Energy of photon = h.ν (h is Planck s Constant )

12 C. Emission of Energy by Atoms Atoms can give off light They first must receive energy and become excited. The energy is released in the form of a photon.

13 Typical Colors From Flame Tests K + Na + Periodic Table of Fireworks Which flame represents a higher energy transition? Li + Cu 2+

14 Objectives 1. To understand how the emission spectrum of hydrogen demonstrates the quantized nature of energy 2. To learn about Bohr s model of the hydrogen atom 3. To understand how the electron s position is represented in the wave mechanical model

15 A. The Energy Levels of Hydrogen Only certain types of photons are produced when excited Hydrogen atoms release energy. Why?

16 A. The Energy Levels of Hydrogen Atomic states Excited state atom with excess energy Ground state atom in the lowest possible state When an H atom absorbs energy from an outside source it enters an excited state.

17 A. The Energy Levels of Hydrogen Energy level diagram Energy in the photon corresponds to the energy used by the atom to get to the excited state.

18 A. The Energy Levels of Hydrogen Only certain types of photons are produced when H atoms release energy. Why?

19 A. The Energy Levels of Hydrogen Quantized Energy Levels Since only certain energy changes occur the H atom must contain discrete energy levels.

20 B. The Niels Bohr Model of the Atom ( ) Bohr s model of the atom Quantized energy levels Electron moves in a circular orbit Electrons jump between levels by absorbing or emitting photons of a particular wavelength Able to mathematically explain the emission spectrum of Hydrogen (compare with Rutherford)

21 B. The Bohr Model of the Atom Bohr s model of the atom was not totally correct. Had difficulties with spectra of larger atoms Electrons do not move in a circular orbit and don t seem to behave like discrete particles all the time Maybe small particles, such as electrons, can also behave like waves having a dual nature (just like photons)?

22 C. The Wave Mechanical Model of the Atom (de Broglie and Schroedinger mid-1920 s) Orbitals Nothing like orbits Probability of finding the electron within a certain space Bohr- Schrodinger

23 Objectives 1. To learn about the shapes of the s, p and d orbitals 2. To review the energy levels and orbitals of the wave mechanical model of the atom 3. To learn about electron spin

24 A. The Hydrogen Orbitals Orbitals do not have sharp boundaries. 90% boundary

25 A. The Hydrogen Orbitals Hydrogen Energy Levels Hydrogen has discrete energy levels. Called principal energy levels (electron shells) Labeled with whole numbers Energy is related to 1/n 2 E n = E 1 /n 2 Energy levels are closer together the further they are from the nucleus

26 A. The Hydrogen Orbitals Hydrogen Energy Levels Each principal energy level is divided into sublevels. Labeled with numbers and letters Indicate the shape of the orbital

27 Orbital Designations

28 Orbitals Define the Periodic Table Orbitals are Energy Levels

29 A. The Hydrogen Orbitals Hydrogen Energy Levels The s and p types of sublevel

30 Representation of s, p, d atomic orbitals

31 A. The Hydrogen Orbitals Hydrogen Orbitals Why does an H atom have so many orbitals and only 1 electron? An orbital is a potential space for an electron. Atoms can have many potential orbitals. s, p, d, f orbitals named for sharp, principal, diffuse and fundamental lines on spectra. Further orbitals designated alphabetically

32 s p d d Orbitals f f f g g g g

33 B. The Wave Mechanical Model: Further Development Electron Spin Close examination of spectra revealed doublets Need one more property to determine how electrons are arranged Spin electron modeled as a spinning like a top Spin is the basis of magnetism

34 B. The Wave Mechanical Model: Further Development Pauli Exclusion Principle Pauli Exclusion Principle (Wolfgang Pauli 1925) - an atomic orbital can hold a maximum of 2 electrons and those 2 electrons must have opposite spins When an orbital contains two electrons (of opposite spin) it is said to be full What are the four descriptors that define an energy level / electron s position in an atom?

35 Objectives 1. To understand how the principal energy levels fill with electrons in atoms beyond hydrogen 2. To learn about valence electrons and core electrons 3. To learn about the electron configurations of atoms 4. To understand the general trends in properties in the periodic table

36 A. Electron Arrangements in the First 18 Atoms on the Periodic Table H atom Electron configuration electron arrangement 1s 1 Orbital diagram orbital is represented as a box with a designation according to its sublevel. Contains arrow(s) to represent electrons (spin)

37 A. Electron Arrangements in the First 18 Atoms on the Periodic Table He atom Electron configuration 1s 2 Orbital diagram

38 A. Electron Arrangements in the First 18 Atoms on the Periodic Table Li atom Electron configuration 1s 2 2s 1 Orbital diagram Write the electron configuration and orbital diagrams for Boron, Nitrogen, Fluorine and Argon

39 A. Electron Arrangements in the First 18 Atoms on the Periodic Table Write the full electron configuration of Neon and Sulfur Draw an orbital diagram for Magnesium and Chlorine

40 Order of Filling of Orbitals Atoms fill their orbitals in the order of their energies:

41 B. Electron Configurations and the Periodic Table Orbital filling and the periodic table Dynamic Periodic Table PT 2a PT 2b

42 B. Electron Configurations and the Periodic Table Electron configurations for K through Kr

43 B. Electron Configurations and the Periodic Table If there were more elements.

44 A. Electron Arrangements in the First 18 Atoms on the Periodic Table Classifying Electrons Valence electrons electrons in the outermost (highest) principal energy level of an atom Core electrons inner electrons Elements with the same valence electron arrangement (same group) show very similar chemical behavior.

45 B. Electron Configurations and the Periodic Table

46 Using a Noble Gas Shorthand We can abbreviate electron configurations by using the configuration of the previous noble gas to cover the first part of the list of orbitals Mg is 1s 2 2s 2 2p 6 3s 2 or [Ne] 3s 2 The noble gas portion is the equivalent to the group of core electrons Use the Noble Gas shorthand to show the electron configurations of Carbon, Chlorine and Zirconium

47 C. Atomic Properties and the Periodic Table Metals and Nonmetals Metals tend to lose electrons to form positive ions. Nonmetals tend to gain electrons to form negative ions.

48 C. Atomic Properties and the Periodic Table Atomic Size Size tends to increase down a column. Size tends to decrease across a row. Effects of Shielding of outer electrons by inner orbitals (close to scale) Larger

49 C. Atomic Properties and the Periodic Table Ionization Energies Ionization Energy energy (ΔH) required to remove an electron from an individual atom (gas) Tends to decrease down a column Tends to increase across a row Changes in an opposite direction to atomic size

50 Ionization Energies

51 Electron Affinity Electron Affinity is defined as ΔH for the process: X (g) + e - = X (g) - ΔH = Electron Affinity Larger

52 Electronegativity Ionization Energy and Electron Affinity are combined to give Electronegativity a measure of how well atoms compete for electrons in a bond Larger

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