Chapter 5

 The scale model shown is a physical model. However, not all models are physical. In fact, several theoretical models of the atom have been developed over the last few hundred years. You will learn about the currently accepted model of how electrons behave in atoms. 5.1

 The Development of Atomic Models ◦ What was inadequate about Rutherford’s atomic model? 5.1

 Rutherford’s atomic model could not explain the chemical properties of elements.  Rutherford’s atomic model could not explain why objects change color when heated. 5.1

 The timeline shows the development of atomic models from 1803 to

 The timeline shows the development of atomic models from 1913 to

 The Bohr Model ◦ What was the new proposal in the Bohr model of the atom?  Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus. 5.1

◦ Each possible electron orbit in Bohr’s model has a fixed energy.  The fixed energies an electron can have are called energy levels or shells. See periodic table  A quantum of energy is the amount of energy required to move an electron from one energy level (shell) to another energy level. 5.1

 Like the rungs of the strange ladder, the energy levels in an atom are not equally spaced.  The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level. 5.1

 The Quantum Mechanical Model ◦ What does the quantum mechanical (wave mechanical, electron cloud) model determine about the electrons in an atom? 5.1

◦ The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.  The exact location of an electron cannot be known with certainty because to do so you would have to interfere with it (Heisenberg Uncertainty Principle). 5.1

◦ Austrian physicist Erwin Schrödinger (1887–1961) used new theoretical calculations and results to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom.  The modern description of the electrons in atoms, the quantum mechanical model, comes from the mathematical solutions to the Schrödinger equation. 5.1

 The propeller blade has the same probability of being anywhere in the blurry region, but you cannot tell its location at any instant. The electron cloud of an atom can be compared to a spinning airplane propeller. 5.1

 In the quantum mechanical (electron cloud) model, the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. The cloud is more dense where the probability of finding the electron is high. 5.1

 Atomic Orbitals ◦ An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron. 5.1

 Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell- shaped. 5.1

 Four of the five d orbitals have the same shape but different orientations in space. 5.1

F orbitals are weird.

 If this rock were to tumble over, it would end up at a lower height. It would have less energy than before, but its position would be more stable. You will learn that energy and stability play an important role in determining how electrons are configured in an atom. 5.2

 Electron Configurations ◦ The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations.  Electron configurations are given in your periodic table  ex: Na  Two electrons are in the first energy level, eight in the second and 1 in the third.  Ex: Br  Two electrons in the first, eight in the second, eighteen in the third and seven in the fourth 5.2

 Neon advertising signs are formed from glass tubes bent in various shapes. An electric current passing through the gas in each glass tube makes the gas glow with its own characteristic color. You will learn why each gas glows with a specific color of light. 5.3

◦ Sunlight consists of light with a continuous range of wavelengths and frequencies.  When sunlight passes through a prism, the different frequencies separate into a spectrum of colors.  In the visible spectrum, red light has the longest wavelength and the lowest frequency. 5.3

 The electromagnetic spectrum consists of radiation over a broad band of wavelengths. 5.3

 Atomic Spectra ◦ What causes atomic emission spectra?  When atoms absorb energy, electrons move into higher energy levels. These electrons then lose energy by emitting light when they return to lower energy levels. It looks pretty 5.3

 A prism separates light into the colors it contains. When white light passes through a prism, it produces a rainbow of colors. 5.3

 When light from a helium lamp passes through a prism, discrete lines are produced. 5.3

 The frequencies of light emitted by an excited element separate into discrete lines to give the atomic emission spectrum of the element.  Note: this spectrum is UNIQUE to each element. 5.3 Mercury Nitrogen

 The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron.  Therefore each transition produces a line of a specific frequency in the spectrum. 5.3

 The three groups of lines in the hydrogen spectrum correspond to the transition of electrons from higher energy levels to lower energy levels. 5.3

 When electrons have their lowest possible energy, they are in the lowest levels available, and the atom is in the ground state. Ex: Na  When electrons absorb energy (heat, light, sound…) it raises them to a higher level, and the atom is in an excited state. Ex: Na  Note: all the electron configurations in your periodic table are ground state. 5.3