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Chapter 4 Electrons In Atoms.

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Presentation on theme: "Chapter 4 Electrons In Atoms."— Presentation transcript:

1 Chapter 4 Electrons In Atoms

2 Chapter 4 Section 1 New Atomic Model

3 Objectives Explain the mathematical relationship among speed, wavelength and frequency of electromagnetic radiation. Discuss the dual wave-particle nature of light. Describe the photoelectric effect. Describe the Bohr model of the atom.


5 Rutherford Model Was an improvement over previous models.
Helped to explain the positively charged nucleus. It did not explain where the atom’s negatively charged electrons are located in space around the nucleus.

6 Light and Electrons To begin to grasp the nature of electrons, examining the nature of light is necessary. We will begin by first introducing some properties of light. We will then see how these properties are related to the properties of the electron.

7 Properties of Light Light behaves as waves and has wave-like properties. Electromagnetic Radiation – a form of energy that exhibits wavelike behavior as it travels through space. Kinds of electromagnetic radiation include visible light, X rays, ultraviolet and infrared light, microwaves and radio waves.

8 Properties of Light Electromagnetic Spectrum – includes all forms of electromagnetic radiation.

9 Electromagnetic Spectrum

10 Properties of Light All forms of electromagnetic radiation move at a constant speed. 3.0 x 108 m/s This is considered the speed of light.

11 A significant feature of waves is its repetitive nature.
Waves can be characterized by two features: Wavelength (l) - the distance between corresponding points on adjacent waves. The units for wavelength are meters, centimeter and nanometer depending on the form of electromagnetic radiation.

12 Wavelength

13 Frequency (n) – defined as the number of waves that pass a given point in a specific time, usually one second (hertz - Hz).

14 Frequency and wavelength are related by the following equation:
c = ln c = speed of light l = wavelength = frequency

15 c = ln Because c is the same for all electromagnetic radiation, the product ln is a constant. is inversely proportional to n As the wavelength (l) of light increase, its frequency (n) decreases, and vice versa.


17 The Photoelectric Effect
Photoelectric Effect – refers to the emission of electrons from a metal when light shines on the metal. When light strikes a metal, no electrons were emitted if the light’s frequency was below a certain minimum. Wave theory of light predicted any frequency of light could eject an electron.

18 Photoelectric Effect Experiment

19 The Photoelectric Effect
The explanation for the photoelectric effect is attributed to German physicist Max Planck. Planck proposed that objects emit energy in small, specific amounts called quanta. Quantum – the minimum quantity of energy that can be lost or gained by an atom. E = hn

20 The Photoelectric Effect
E = hn E = energy n = frequency h = Planck’s constant – x J-s

21 The Photoelectric Effect
This energy can also be related to its wavelength by the following equations: E = hn and c = ln to get: E = hc l

22 The Photoelectric Effect
Albert Einstein expanded on Planck’s theory by explaining that electromagnetic radiation has a dual wave-particle nature. Light can also be thought of as a stream of particles. Each particle of light carries a quantum of energy.

23 The Photoelectric Effect
Einstein called these particles photons. Photon – a particle of electromagnetic radiation having zero mass and carrying a quantum of energy. The energy of a particular photon depends on the frequency of radiation: Ephoton = hn

24 The Photoelectric Effect
Summary: Light has both wave properties (l and n) and particle (photons) properties. In order for an electron to be ejected from a metal surface, the electron must be struck by a single photon possessing the minimum energy (frequency and wavelength).

25 Atom Line Emission Spectrum
Ground State – the lowest energy state of an atom. Excited State – A state in which an atom has a higher energy than it has in its ground state.

26 When an excited atom returns to its ground state, it gives off energy that it gained in the form of electromagnetic radiation. E2 Excited state energy Electromagnetic radiation Electric current E1 Ground state energy

27 When an electric current was passed through a tube containing hydrogen gas, a pink glow of light was emitted. When this pink emitted light was passed through a prism, it was separated into a series of specific wavelengths of visible light. The bands of light were part of what is known as hydrogen’s line-emission spectrum.

28 Hydrogen Atom Line Emission Spectrum

29 Why has the hydrogen atoms given off only specific wavelengths of light?
Scientists had expected to observe the emission of a continuous range of wavelengths of electromagnetic radiation, that is a continuous spectrum. Attempts to explain this observation led to a new theory of the atom call Quantum Theory.

30 Whenever an excited hydrogen atom falls back from an excited state to its ground state, it emits a photon of radiation. The energy of this photon is: Ephoton = hn This energy is equal to the difference in energy between the atom’s excited state (E2) and its ground state (E1). E2 – E1 = Ephoton = hn

31 Energy difference between ground and excited state

32 The fact that hydrogen atoms emit only specific wavelengths of light indicated that the energy differences between the atom’s energy states were fixed. This suggested that the electron of a hydrogen atom exists only in very specific energy states.

33 Bohr Model of the Hydrogen Atom
Niels Bohr, a Danish physicist explained the line spectrum of hydrogen in 1913. His model combined the concepts of Planck and Einstein. Ephoton = hn Bohr assumed the atom contained a nucleus and that the electrons circled the nucleus in circular orbits.


35 Bohr Model The three postulates of the Bohr model:
The electron in the hydrogen atom may only occupy orbits of certain radii that correspond to certain discrete energies. While an electron is in an allowed energy orbit, it does not radiate energy and it remains in that orbit without crashing into the nucleus.

36 Bohr Model 3) An electron may move from one energy state to another by absorbing or releasing energy. The energy needed is the difference between one energy level and another and is equal to a photon, Ephoton = hn

37 Bohr Model

38 Bohr Model

39 Bohr Model Ephoton = hn By knowing the wavelengths from the hydrogen atom line emission spectrum, Bohr could solve for the energy of the photon using the above equation. This energy (Ephoton) represents the difference in energy between the different orbits of the hydrogen atom.

40 Bohr Model While the Bohr model works well for hydrogen, it does have its limitations: It did not work well with atoms with more than one electron. It does not account for electron-electron repulsions. Additional electron-nucleus interactions present problems.

41 Classwork Section Review, page 97 Questions 1-5

42 Homework Page Questions 1, 6, 9, 31, 33 Collected for a grade

43 light experiments with various gases
Lab Demo light experiments with various gases

44 Electron Configurations
Chapter 4 Section 3 Electron Configurations

45 Objectives List the atomic orbitals of an atom.
List the total number of electrons needed to fully occupy each main energy level. State the Aufbau principle, the Pauli Exclusion principle and Hund’s rule. Write the electron configuration for any element.


47 Atomic Orbitals Quantum Mechanical Model
A more complex, highly mathematical model was developed to explain observations of atoms containing more than one electron. This model works for all the elements and not just for hydrogen as in the Bohr model.

48 Electronic Configuration – describes the arrangement of electrons in an atom.
Because atoms of different elements have different number of electrons, a distinct electron configuration exists for each element.

49 The electrons will assume arrangements that have the lowest possible energies.
Ground State Configuration – the lowest energy arrangement of the electrons for each element.

50 Atomic Orbitals Bohr Model – the orbit of the electron was circular around the nucleus. In the quantum mechanical model the simple circular orbit was replaced with 3D orbitals (electron clouds) of various shapes in which an electron is likely to be.

51 Atomic Orbitals There are four main atomic orbitals which describe the electron configuration of the elements: S orbital - spherical shape P orbital – dumbbell shape D orbital – clover shape F orbital – Too complex to discuss.

52 Periodic Table with Orbitals

53 S orbital - spherical shape

54 P orbital – dumbbell shape

55 D orbital – clover shape

56 s, p and d orbitals

57 Atomic Orbitals Energy levels of the three orbitals of interest:
S orbital – lowest energy P orbital – slightly higher in energy D orbital – higher in energy than P orbital

58 Electron Configuration Rules
The number of electrons in an atom is the same as the number of protons. So the periodic table will be of real value in determining electron configurations. To build up electron configurations for any particular atom, first energy levels of the orbitals are determined.


60 Electron Configuration Rules
The electrons are added to the orbitals one by one according to three basic rules: Aufbau Principle – An electron occupies the lowest energy orbital that can receive it. The orbital with the lowest energy is the 1s orbital. The one electron of hydrogen goes in this orbital.

61 Electron Configuration Rules
The 2s orbital is the next highest in energy, then the 2p orbitals. The numbers 1,2,3 etc. refer to the row of the periodic table the atom is located in. As can be seen on the diagram there is only 1-s orbital, 3-p orbitals and 5-d orbitals. These refer to their orientation in space.

62 Electron Configuration Rules
Note on the energy level diagram that the 4s orbital is lower in energy than the 3d orbital. Therefore, the 4s orbital is filled before any electrons enter the 3d orbitals.

63 Aufbau Principle

64 Electron Configuration Rules
2) Pauli Exclusion Principle – no more than two electrons may be present in an orbital and their spins must be paired. This rule basically states no two atoms can have the same electron configurations.

65 Electron Configuration Rules
3) Hund’s Rule – orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. The spins of these electrons must be opposite. This rule is because similarly charged electrons want to be as far away as possible. 2p orbital

66 Energy level diagram for oxygen

67 Electronic Configurations
Electronic configurations are important in chemistry: To predict what type of bonding will occur with a particular element and which electrons are being used in the bonding. Helps explain the properties of elements.


69 Electronic Configurations
While energy level diagrams are very useful they are bulky to work with. Electron configuration notations are simpler and give the same information.

70 Electronic Configurations
Electron configuration notations eliminate the lines and arrows of the diagrams. Instead the number of electrons in an energy level is shown by adding a superscript to the energy level designation. Example: hydrogen - 1S1

71 Electronic Configurations
Example: hydrogen - 1S1 The large 1 indicates hydrogen is in the first row of the periodic table. The S indicates the electron is in the s orbital. The superscript 1 indicates that there is one electron in the 1S orbital.

72 Electronic Configurations
Example: helium - 1S2 The superscript 2 indicates that there are two electrons electron in the 1S orbital. Problem: Give the electron configuration of boron and explain how the electrons are arranged.

73 Elements of the Second Period
In the first period elements, hydrogen and helium, electrons occupy the first energy level – 1s. After the 1s orbital is filled, the next electron occupies the 2s orbital – Aufbau principle. Lithium has an electron configuration of 1s22s1

74 Aufbau Principle

75 Classwork Page Problems 1 – 2 Page 116 – Problem 1

76 Elements of the Second Period
Highest Occupied Level – is the electron containing main energy level with the largest number. In the case of lithium that is the 2s level. Inner Shell Electrons – The electrons which are in the levels below the highest occupied level. In the case of lithium that is the 1s level.

77 Elements of the Second Period

78 Elements of the Second Period
When you get to neon (Ne) all the 2s and 2p orbitals are full. Octet Rule – when all of the sublevels (s and p orbitals) of the highest occupied level is filled with eight electrons. All the elements in the last column of the periodic table obey the octet rule.

79 Noble Gases Neon is a member of the Group 18 elements (last column).
These elements include neon, argon, krypton, xenon and radon). These elements are known as the noble gases.

80 Elements of the Third Period

81 Elements of the Third Period
To simplify sodium’s notation, the symbol for neon, enclosed in brackets, is used to represent the complete neon configuration. [Ne] = 1s22s22p6 So the electron configuration for sodium can be written: [Ne]3s1 This is the noble gas configuration

82 Elements of the Fourth Period
With the 4s level full (calcium), the 4p and 3d sublevels are next available. Referring to the Aufbau diagram of energy levels, the 3d sublevel is lower in energy than the 4p sublevel. There are five 3d orbitals that hold a total of 10 electrons. Elements range from Sc to Zn.

83 Elements of the Fourth Period

84 Elements of the Fourth Period

85 Elements of the Fifth Period
Elements in the fifth period start with the 5s orbital. 5s d p

86 Periodic Table with Orbitals

87 Problem Write both the electron configuration and noble gas configuration for iron (Fe). How many electron containing orbitals are in an atom of iron? How many are filled? How many unpaired electrons are there in an atom of iron?

88 Classwork Page 115 Practice Problems 1 – 3 Page 116
Practice Problems 1 and 2

89 Homework Page 118 Problems 22, 24, 25, 30 and 37 Due: For grade

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