10/2/2013 5.1 The timeline shoes the development of atomic models from 1803 to 1911. 5.1 5.1 The timeline shows the development of atomic models from 1913 to 1932. The Development of Atomic Models What was inadequate about Rutherford’s atomic model? 1, 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. 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. Niels Bohr 7 October 1885 – 18 November 1962 1 10/2/2013 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. A quantum of energy is the amount of energy required to move an electron from one energy level 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. Lewis- The total number of electrons represented in a Lewis structure is equal to the sum of the numbers of valence electrons on each individual atom. Non-valence electrons are not represented in Lewis structures. Bohr The Bohr Model has an atom consisting of a small, positively-charged nucleus orbited by negatively-charged electrons. The Quantum Mechanical Model What does the quantum mechanical model determine about the electrons in an atom? 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. 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 2 10/2/2013 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 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. Different atomic orbitals are denoted by Atomic Orbitals How do sublevels of principal energy levels differ? An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron. Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found. letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped. 3 10/2/2013 Aufbau Principle Representation of s, p, d atomic orbitals Four of the five d orbitals have the same shape but different orientations in space. Hydrogen Energy Levels The s and p types of sublevel 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 4 10/2/2013 The number of electrons allowed in each of the first four energy levels are shown here. s p f f d f g g d g g 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 Hydrogen Energy Levels The s and p types of sublevel The numbers and kinds of atomic orbitals depend on the energy sublevel. 1. Rutherford's planetary model of the atom could not explain any properties of elements. the chemical properties of elements. the distribution of mass in an atom. the distribution of positive and negative charges in an atom. 5 10/2/2013 Bohr's model of the atom proposed that electrons are found embedded in a sphere of positive charge. in fixed positions surrounding the nucleus. in circular orbits at fixed distances from the nucleus. orbiting the nucleus in a single fixed circular path. 2. Electron Configurations What are the three rules for writing the electron configurations of elements? The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations. Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms. 3. What is the lowest-numbered principal energy level in which p orbitals are found? 1 2 3 4 Aufbau Principle According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram below, each box represents an atomic orbital. Aufbau Principle 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. 6 10/2/2013 Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped. First, determine how many electrons are in the atom. Iron has 26 electrons. Arrange the energy sublevels according to increasing energy: 1s 2s 2p 3s 3p 4s 3d … Fill each sublevel with electrons until you have used all the electrons in the atom: Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6 The sum of the superscripts equals the atomic number of iron (26) Electron Configurations and the Periodic Table • The periodic table can be used as a guide for electron configurations. • The period number is the value of n. • Groups 1A and 2A have the s-orbital filled. • Groups 3A - 8A have the p-orbital filled. • Groups 3B - 2B have the d-orbital filled. • The lanthanides and actinides have the f-orbital filled. The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. We can use the periodic table to predict which sublevel is being filled by a particular element. The sublevel is written followed by a superscript with the number of electrons in the sublevel. If the 2p sublevel contains 2 electrons, it is written 2p2 7 10/2/2013 Electron Configurations and the Periodic Table • The periodic table can be used as a guide for electron configurations. • The period number is the value of n. • Groups 1A and 2A have the s-orbital filled. • Groups 3A - 8A have the p-orbital filled. • Groups 3B - 2B have the d-orbital filled. • The lanthanides and actinides have the f-orbital filled. 5th 2nd 6th 1st Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped. The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. The sublevel is written followed by a superscript with the number of electrons in the sublevel. If the 2p sublevel contains 2 electrons, it is written 2p2 The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. The sublevel is written followed by a superscript with the number of electrons in the sublevel. If the 2p sublevel contains 2 electrons, it is written 2p2 8 10/2/2013 When we write the electron configuration of a positive ion, we remove one electron for each positive charge: Na → Na+ 1s2 2s2 2p6 3s1 → 1s2 2s2 2p6 When we write the electron configuration of a negative ion, we add one electron for each negative charge: O 1s2 2s2 2p4 → → O21s2 2s2 2p6 Electron Configurations and the Periodic Table Pauli Exclusion Principle According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons. To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired. Hund’s Rule Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. • The periodic table can be used as a guide for electron configurations. • The period number is the value of n. • Groups 1A and 2A have the s-orbital filled. • Groups 3A - 8A have the p-orbital filled. • Groups 3B - 2B have the d-orbital filled. • The lanthanides and actinides have the f-orbital filled. 9 10/2/2013 http://wine1.sb.fsu.edu/chm1045/notes/Struct/EConfig/Struct08.htm The two electrons in Helium represent the complete filling of the first electronic shell. Thus, the electrons in He are in a very stable configuration For Boron (5 electrons) the 5th electron must be placed in a 2p orbital because the 2s orbital is filled. Because the 2p orbitals are equal energy, it doesn't matter which 2p orbital is filled. Orbital Filling Diagram Ne 1s2 2s2 2p6 F 1s2 2s2 2p5 F- 1s2 2s2 2p6 Mg 1s2 2s2 2p6 3s2 Mg2+ 1s2 2s2 2p6 (10 electrons) (9 electrons) (10 electrons) (12 electrons) (10 electrons) Notice – different elements can have the same number of electrons 10 10/2/2013 11 10/2/2013 12 10/2/2013 13 10/2/2013 14 10/2/2013 15 10/2/2013 Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations. Exceptions to the aufbau principle are due to subtle electronelectron interactions in orbitals with very similar energies. Copper has an electron configuration that is an exception to the aufbau principle. Exceptional Electron Configurations Why do actual electron configurations for some elements differ from those assigned using the aufbau principle? 16 10/2/2013 1. Identify the element that corresponds to the following electron configuration: 1s22s22p5. 3. Spectral lines in a series become closer together as n increases because the F energy levels have similar values. Cl Ne energy levels become farther apart. atom is approaching ground state. O electrons are being emitted at a slower rate. 2. N. Write the electron configuration for the atom 1s22s22p5 1s22s22p3 1s22s1p2 1s22s22p1 3. The electron configurations for some elements differ from those predicted by the aufbau principle because the the lowest energy level is completely filled. none of the energy levels are completely filled. half-filled sublevels are less stable than filled energy levels. half-filled sublevels are more stable than some other arrangements. ` 17 10/2/2013 Ground state – atoms with their electrons in the lowest energy level possible; this lowest energy state is the most stable. Excited state – a higher energy state; electrons jump to higher energy levels by absorbing energy Atom is less stable in an excited state; it will release the extra energy to return to the ground state 103 . The configuration that corresponds to the lowest electronic energy is called the ground state. Any other configuration is an excited state. As an example, the ground state configuration of the sodium atom is 1s22s22p63s1 , as deduced from the Aufbau principle . The first excited state is obtained by promoting a 3s electron to the 3p orbital, to obtain the 1s22s22p63p1 configuration, abbreviated as the 3p level. Atoms can move from one configuration to another by absorbing or emitting energy 18 10/2/2013 He showed that the energy of a quantum is related to the frequency of emitted radiation by the equation Equantum=hv Where E is energy, h is Plancks constant and v is frequency Plancks constant is 6.626 x 10-34 J*s J is the symbol for the joule Check how as energy increases, v will increase—(why violet light has more energy than red light) The photoelectric effect is when electrons are emitted from a metals surface when a certain frequency of light shines on the surface. The mystery concerns the frequency, and therefore color of the light. Too low a frequency, no matter how long or intense it is will not eject electrons. As long as the frequency is greater than 1.14 x 10 15 Hz, electrons will be ejected. In 1900 Max Planck searched for an explanation of why light was emitted from heated objects. His study lead to a conclusion that matter can gain or lose energy only in small specific amounts called quanta—a quantum is the minimum amount of energy that can be lost or gained by an atom. Einstein proposed in 1905 that electromagnetic radiation had both waveleline and particlelike natures and called these bundles of energy photons.—particles of electromagnetic radiation with no mass that carries a quantum of energy. Einstein calculated that a photons energy depended on its frequency. E photon = hv There has to be a minimum value to cause the ejection of a photoelectron 19 10/2/2013 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. Light How are the wavelength and frequency of light related? The wavelength and frequency of light are inversely proportional to each other. The product of the frequency and wavelength always equals a constant (c), the speed of light. The amplitude of a wave is the wave’s height from zero to the crest. The wavelength, represented by λ (the Greek letter lambda), is the distance between the crests. The frequency, represented by ν (the Greek letter nu), is the number of wave cycles to pass a given point per unit of time. The SI unit of cycles per second is called a hertz (Hz). According to the wave model, light consists of electromagnetic waves. Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays. All electromagnetic waves travel in a vacuum at a speed of 2.998 × 108 m/s. 20 10/2/2013 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. The electromagnetic spectrum consists of radiation over a broad band of wavelengths. 21 10/2/2013 A prism separates light into the colors it contains. When white light passes through a prism, it produces a rainbow of colors. c/ λ =v When light from a helium lamp passes Practice problems 1-4 p121 H and p140—14 & 15 regs through a prism, discrete lines are produced. The frequencies of light emitted by an 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. element separate into discrete lines to give the atomic emission spectrum of the element. Mercury Nitrogen 22 10/2/2013 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. An Explanation of Atomic Spectra How are the frequencies of light an atom emits related to changes of electron energies? 5.3 In the Bohr model, the lone electron in the hydrogen atom can have only certain specific energies. When the electron has its lowest possible energy, the atom is in its ground state. Excitation of the electron by absorbing energy raises the atom from the ground state to an excited state. A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level. Quantum Mechanics How does quantum mechanics differ from classical mechanics? In 1905, Albert Einstein successfully explained experimental data by proposing that light could be described as quanta of energy. The quanta behave as if they were particles. Light quanta are called photons. In 1924, De Broglie developed an equation that predicts that all moving objects have wavelike behavior. 5.3 Today, the wavelike properties of beams 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. of electrons are useful in magnifying objects. The electrons in an electron microscope have much smaller wavelengths than visible light. This allows a much clearer enlarged image of a very small object, such as this mite. 23 10/2/2013 Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves. 2.40 109 Hz 2.40 1024 Hz 2.40 106 Hz 2.40 102 Hz The Heisenberg uncertainty principle states that it is impossible to know exactly both the velocity and the position of a particle at the same time. This limitation is critical in dealing with small particles such as electrons. This limitation does not matter for ordinary-sized object such as cars or airplanes. 1. Calculate the frequency of a radar wave with a wavelength of 125 mm. 2. The lines in the emission spectrum for an element are caused by the movement of electrons from lower to higher energy levels. the movement of electrons from higher to lower energy levels. the electron configuration in the ground state. the electron configuration of an atom. The Heisenberg Uncertainty Principle 24 10/2/2013 25