CHAPTER 4. Arrangement of Electrons in Atoms

Transcription

1 CHAPTER 4 Arrangement of Electrons in Atoms

2 4.1 Part I Development of a New Atomic Model

3 4.1 Objectives 1. Explain the mathematical relationship among the speed, wavelength, and frequency of electromagnetic radiation. 2. Discuss the dual wave-particle nature of light. 3. Discuss the significance of the photoelectric effect and the line-emission spectrum of hydrogen to the development of the atomic model. 4. Describe the Bohr model of the hydrogen atom.

4 4.1 Development of a New Atomic Model Rutherford s model of the atom could not explain the distribution of electrons What prevented the negative electrons from being drawn into the positive nucleus? Studies of light in the early twentieth century actually led to a new atomic model Scientists discovered a relationship between light and electrons!

5 4.1 Wavelike Properties of Light Before 1900, scientists thought light behaved just as a wave Wave A disturbance that travels from one location to another location Transfer energy but not mass What is energy?

6 4.1 Energy What is energy? Ability to do work Work means a change in position, speed, state, or form of matter. Energy is the capacity to change matter. Electromagnetic radiation A form of energy that exhibits wavelike behavior as it travels through space Can travel through a vacuum, unlike other types of waves

7 4.1 Wavelike Properties of Light Together, all the forms of electromagnetic radiation form the electromagnetic spectrum

8 4.1 Wavelike Properties of Light Wavelength (λ) the distance between corresponding points on adjacent waves Frequency (v) the number of waves that pass a given point in a specific time, usually one second Measured in Hertz (Hz) = waves/s Wavelength and frequency are mathematically related

9 4.1 Wavelike Properties of Light The speed of electromagnetic radiation is constant Speed of light (c) c = 3.00x10 8 m/s = λv Inverse relationship between frequency and wavelength As the wavelength gets longer, the frequency gets lower As the wavelength gets shorter, the frequency gets higher Higher frequency = MORE energy Lower frequency = LESS energy

10 4.1 Frequency and Energy Continuous spectrum ROY G BIV

11 4.1 Sample Problem What is the frequency of light that has a wavelength of 310 nm? λ = 310 nm = 310x10-9 m c = λv 3.00x10 8 m/s = (310x10-9 m)*v v = 9.7 x Hz

12 4.1 Ionizing Radiation High frequency radiation with enough energy to free electrons from atoms or molecules, thereby ionizing them Can be either: 1. energetic subatomic particles, ions or atoms moving at relativistic speeds 2. electromagnetic waves on the high-energy end of the electromagnetic spectrum

13 4.1 Part II Development of a New Atomic Model

14 4.1 Objectives 1. Explain the mathematical relationship among the speed, wavelength, and frequency of electromagnetic radiation. 2. Discuss the dual wave-particle nature of light. 3. Discuss the significance of the photoelectric effect and the line-emission spectrum of hydrogen to the development of the atomic model. 4. Describe the Bohr model of the hydrogen atom.

15 4.1 New Discoveries In the early 1900s, scientists studying the interaction of light and matter discovered things that couldn t be explained by the wave theory of light 1. Photoelectric effect 2. Hydrogen-atom line-emission spectrum

16 4.1 The Photoelectric Effect Photoelectric effect Electrons are emitted from the surface of a metal when light shines on it Wave theory of light predicted that light of any frequency could supply enough energy to eject an electron, but that wasn t the case For a given metal, no electrons were emitted if the light s frequency was below a certain minimum, regardless of the light s intensity

17 4.1 Planck s Postulate This unexpected behavior of electrons was being studied by scientists around the 1900s Max Planck studied the emission of light by hot objects and made a revolutionary discovery in 1900 He concluded that the light energy is NOT emitted continuously, as would be expected if they were waves Instead, they were emitted in small, specific packets of energy called quanta Quantum of energy the minimum quantity of energy that can be lost or gained by an atom E=hv E = energy in joules of a quantum of radiation, h = Planck s constant (6.626x10-34 J*s), v = frequency in s 1

18 4.1 The Particle Theory of Light In 1905, Einstein expanded Planck s theory by introducing the idea that electromagnetic radiation has a dual wave-particle nature Photons Particle of electromagnetic radiation having 0 mass and carrying one quantum of energy Einstein said that electromagnetic radiation is absorbed by matter only in whole numbers photons In order for an electron to be ejected from a metal surface, it must be struck by a photon possessing a minimum amount of energy Since different metals bond electrons more or less tightly, they require different minimum frequencies

19 Dual Nature of Light So what is light, a particle or a wave!?? It is BOTH! 58&v=J1yIApZtLos DQH5x7svfg

20 4.1 Line-Emission Spectrum The lowest energy state of an atom is its ground state. Atoms can also exist at different excited states, where they have a higher potential energy than the ground state When an excited atom returns to a lower energy state or the ground state, it emits the energy as light Neon signs, fireworks

21 4.1 Line-Emission Spectrum When the emitted light is separated through a prism, it takes on unique banding patterns, NOT a continuous spectrum Each element emits different wavelengths of light, and has a unique emission-line spectrum

22 4.1 Line-Emission Spectrum WHY don t atoms emit a continuous spectrum? When an excited H electron falls to its ground state, it emits a photon of radiation whose energy is equal to the difference in energy between the two states Since H atoms only emit specific frequencies of light, the energy differences are fixed Therefore, electrons of a H atom exist only in specific energy states

23 4.1 Bohr Model of the Hydrogen Atom In 1913, Niels Bohr solved the puzzle of the hydrogen atom spectrum by linking electrons to photon emission Electrons circle the nucleus only in allowed paths, or orbits Electrons cannot exist in the empty spaces between orbits The orbits closer to the nucleus are lower energy orbits Like the rungs of a ladder Model only worked for hydrogen Electrons can absorb photons of energy to move up energy levels or emit photons when moving down energy levels

24 4.2 The Quantum Model of the Atom

25 4.2 Objectives 1. Discuss Louis de Broglie s role in the development of the quantum model of the atom. 2. Compare and contrast the Bohr model and the quantum model of the atom. 3. Explain how the Heisenberg uncertainty principle and the Schrödinger wave equation led to the idea of atomic orbitals. 4. List the four quantum numbers and describe their significance. 5. Relate the number of sublevels corresponding to each of an atom s main energy levels, the number of orbitals per sublevel, and the number of orbitals per main energy level.

26 4.2 The Quantum Model of the Atom In 1924, Louis De Broglie stated that electrons could also behave as both a wave and a particle He said the behavior of electrons in Bohr s orbits was like that of waves confined to space around the nucleus Other experiments confirmed that electrons can also be bent and interfere with each other, just like waves

27 4.2 Even Stranger If electrons are both waves and particles, where are they in the atom? Heisenberg s Uncertainty Principle In 1927, Werner Heisenberg stated that it is impossible to determine both the position and velocity of an electron, because it is changed by the observation of it Like a mouse and a flashlight Schrodinger s Wave Equation In 1926, Erwin Schrodinger developed an equation describing quantum mechanical behavior, treats electrons as waves

28 4.2 Quantum Theory The Heisenberg uncertainty principle and Schrodinger wave equation laid the foundation for modern quantum theory Quantum theory Mathematically describes the wave properties of electrons and other very small particles Quantum theory states that you can only determine the probability of finding an electron at a given place around the nucleus, they exist in orbitals Orbital A 3D region around the nucleus that indicates the PROBABLE location of an electron

29 4.2 Quantum Mechanical Model The quantum model of the atom expands on Bohr s model Properties of atomic orbitals and properties of electrons in orbitals are described with 4 quantum numbers: 1. Main energy level 2. Shape of the orbital 3. Orientation of the orbital 4. Spin of the electron

30 4.2 Principal Quantum Number Indicates the main energy level occupied by the electron Can only be positive integers n=1,2,3,4,5,6,7 More than one electron can have the same n value Electrons with the same n are in the same electron shell

31 4.2 Angular Momentum Quantum # 4 different sublevels - indicates the shape of the orbital Except for n=1, each level has sublevels with different shapes Orbitals are designated by nl, such as 1s, 2p etc Letter s p d f spherical dumbbell

32 4.2 Magnetic Quantum # Multiple orbitals can have the same shape, but different orientation of an orbital around the nucleus 3-D orientation using the x, y & z axis Sublevel # of orbitals s 1 p 3 d 5 f 7

33 4.2 Spin Quantum Number The easiest quantum number! It has only two values + ½, -½ Indicates one of 2 fundamental spin states of an electron A single orbital can hold a maximum of two electrons, but they must have opposite spin states Indicated with arrows, one up, one down

34 4.2 Quantum Numbers Summary

35 4.3 Electron Configuration

36 4.3 Objectives 1. List the total number of electrons needed to fully occupy each main energy level. 2. State the Aufbau principle, the Pauli exclusion principle, and Hund s rule. 3. Describe the electron configurations for the atoms of any element using orbital notation, electronconfiguration notation, and, when appropriate, noblegas notation.

37 4.3 Electron Configuration Rules Each element has a different number of electrons Therefore, elements have unique electron configurations Electrons arrange themselves at the lowest possible energy, called the ground-state electron configuration There are three rules to build ground-state electron configurations

38 4.3 Aufbau Principle An electron occupies the lowest-energy orbital that can receive it The two diagrams on the right represent the order that orbitals are filled. Write the order that orbitals are filled: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d

39 4.3 Pauli Exclusion Principle No two electrons in the same atom can have the same set of four quantum numbers Therefore if two electrons occupy the same orbital, they must have opposite spin states

40 4.3 Hund s Rule Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron All electrons in singly occupied orbitals must have the same spin state

41 4.3 Representing Electron Configurations Three methods 1. Orbital notation uses lines and arrows 2. Electron configuration notation written shorthand

42 4.3 Noble Gas Configuration Write the electron configuration for Ne. [Ne] = 1s 2 2s 2 2p 6 Can use the noble gas electron configurations to shorten the notation for elements above the second period Put the name of the noble gas in brackets to represent inner-shell electrons, only show the outer energy level electrons

43 4.3 Practice Problems Write the complete electron configuration for the element with atomic number 25. Identify the element. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 manganese

44 4.3 Practice Problems Write the complete electron-configuration notation and the noble-gas notation for iodine, I. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5 [Kr] 5s 2 4d 10 5p 5