Development of the Periodic Table. Chapter 5. Light and the EM Spectrum. Light

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1 Chapter 5 Periodic Table Song Periodicity and Atomic Structure Development of the Periodic Table Mid-1800 s, several scientists placed known elements in order based on different criteria. Mendeleev s and Meyer s versions, Modern Arrangement What s Going on Inside the Atom? Once scientists developed a logical order for the elements they began studying the structure and composition of individual atoms. They used substances interactions with light to explain the structure of atoms and develop a model to explain how atoms affected properties of light. In order to understand interactions, we must understand behavior of light. Light Light and the EM Spectrum EM waves Light is typically described as traveling in waves (similar to water); All electromagnetic (EM) waves (including light) are made of two components: electric and magnetic EM waves travel at the speed of light, c ( x 10 8 m/s 3.00 x 10 8 m/s) c = λν (Know these variables!) c = speed of light; λ (lambda) = wavelength (m, nm); ν (nu) = frequency (1/s, s -1, Hz) 1

2 EM Spectrum Different colors of light correspond to different wavelengths in the visible portion of the EM spectrum. Two wavelengths (λ) are shown below. Determine the frequency (ν) for each wave. Blue light Red light 1 nm = 1 x 10-9 m OR 1 x 10 9 nm = 1 m Describing Atoms Classical descriptions: Dalton: atoms are hard particles, all atoms of the same element are the same Thomson: atoms are divisible (electrons in atoms) Rutherford: positively charged nucleus New view of atomic behavior Planck: Blackbody radiation heat solids to red or white heat, matter did not emit energy continuously; in whole-number multiples of certain quantities Matter absorbs or emits energies in packets - quanta From Classical to Quantum Theory Quantum has come to mean small; originated from Planck s observation of quantized energy Einstein used this theory to observe metals reacting to different colors of light Photoelectric Effect: electrons are ejected from the surface of certain metals exposed to light at a certain minimum frequency Blue light (ν = 6.7 x Hz) causes Na to emit electrons, red light (ν = 4.0 x Hz) does not Wave-Particle Duality Based on photoelectric effect, light acts as a wave but also exists as a stream of particles called photons Energy of photons is proportional p to frequency, inversely proportional to wavelength h c E = hν = h = x J s λ J = kg m 2 / s 2 Photoelectric Effect Calculation Practice c = λν; E = hν 1) Which has a higher frequency: light from a red stoplight with a wavelength of 750 nm or a yellow light with a wavelength of 600 nm? 2) What is the wavelength of a radio station s waves transmitting at a frequency of MHz (megahertz)? (FM radio waves range from MHz.) 3) Red lights at traffic stops have wavelengths of about 650 nm. What is the frequency (in Hz) of this light? 4) Calculate the energy (in Joules) of a photon with a wavelength of 5.00 x 10 4 nm (infrared region). Worked Ex. 5.1, 5.4; Problems 5.1, 5.2, 5.3, 5.7, 5.8 Wavelike Properties of Matter de Broglie: If light can behave like a wave and a particle, then matter (i.e., electrons) can behave like a wave If an electron behaves like a standing wave, then it can only have specific wavelengths Can calculate wavelength for matter if we know its velocity (use v instead of c): λ = h / m v (This is the de Broglie equation.) h = Planck s constant, m = mass (electron s have constant mass: 9.11x10-31 kg), v = velocity (speed) Worked Ex. 5.5; Problem 5.9 2

3 Calculation Practice The energy of a photon is 5.87 x J. What is the frequency of the photon? What is the wavelength of an electron that travels at 34.7 m/s and has a mass of 9.11 x kg? A kg baseball is thrown at a velocity of 42.5 m/s. Calculate the wavelength of the baseball. How does the baseball s wavelength compare to the electron from the example above? Heisenberg Uncertainty Principle If electrons have wavelike properties, then we can t know both its position and velocity. In order to determine the position of an electron, we hit it with a photon of light, but this will change its position and velocity. Bohr s Model of the Atom Bohr s Model of the Atom Bohr sought to reconcile these views of the electron. Developed the planetary analogy of atoms. Electrons orbit around the nucleus like planets around the sun. Electrons travel in discrete, quantized circular orbits; like going up or down stairs. Each orbit has a specific energy associated with it, labeled as n = 1, 2, etc. Ground state is the lowest energy level for an atom (n = 1). Bohr s Model of the Atom Transitions between Energy Levels When an atom absorbs energy, an electron can jump from a lower energy level to a higher energy level. When an atom emits (releases) energy, an electron drops from a higher energy level to a lower energy level. This process sometimes gives off energy as visible light. Eng. Color H e - transitions 3

4 Visible Light White light we see consists of all colors in the visible spectrum. Use a prism (or CD) to break them up. cont. spect. white light Light given off by atoms doesn t necessarily correspond to all visible colors. Flame tests Emission Spectra Hydrogen Each element gives off unique spectrum Demo: Gas Discharge Tubes Each element has its own individual emission spectrum. This allowed scientists to identify elements in different minerals. Spectra of Elements: Emission Spectra of Elements Figure 7.8 Quantum Mechanical Model The Bohr model worked well for hydrogen, but failed for elements with more than one proton and one electron. Quantum Mechanics was developed (by Schrödinger in the 1920 s) to describe the motion of subatomic particles Did not attempt to describe position of particles; used mathematical equations to describe the probability of finding the particles The probability density (map of likely locations) is the electron cloud Atomic Orbitals Quantum Mechanics Movie Quantum Numbers The region of highest probability for finding an electron is an electron cloud. This region of high probability is called an atomic orbital. Each orbital holds at most 2 electrons. Modern atom.exe e- orbit vs e- cloud 23 There are 4 quantum number that describe the size, shape, and location of electrons We use these numbers to describe where electrons are found for an atom. Can also use the periodic table!!! The Principal Quantum Number, n describes distance of the electron from the nucleus; called shells n = 1, 2, 3, etc; larger number is farther from nucleus n corresponds to a row in the periodic table 24 4

5 Orbitals and Quantum Numbers Shapes of Orbitals The Angular Momentum Quantum Number, l In each row of the periodic table are different groups of orbitals with different shapes. These groups of orbitals are called subshells and labeled s, p, d, and f. s subshells are spherical (first two columns) p subshells are dumb-bell b shaped (last six columns) d subshells are intersecting dumb-bells (transition metals) The Magnetic Quantum Number, m l describes the orientation of the orbital with respect to x, y, and z axes s, p, and d orbitals have different shapes and therefore different orientations The Spin Quantum Number, m s describes the spin of an electron in an orbital (shown as up and down arrows in orbital diagrams) s orbitals are spherical; white rings are nodes (regions where an electron won t be found) 1 s orbital in a subshell s orbital Shapes of Orbitals 2p x orbital 2p y orbital 2p z orbital Shapes of Orbitals p orbitals are dumb-bells (2 lobes); node between lobes 3 p orbitals in a subshell p orbital d orbitals: intersecting dumb-bells (4 lobes); nodes between lobes 5 d orbitals in a subshell d orbital Shells and Subshells The first shell (row) has 1 subshell (s) s only 1 orbital An s subshell can hold at most 2 electrons The 2 nd shell (row) has 2 subshells (s and p) p setof3orbitals orbitals A p subshell can hold at most 6 electrons The 3 rd shell (row) has 3 subshells (s, p, and d) d set of 5 orbitals A d subshell can hold at most? electrons The 4 th shell (row) has 4 subshells (s, p, d, and f) f set of 7 orbitals What is the maximum number of electrons allowed in the f subshell?

6 Things that make you go hmmm. What is the maximum number of: electrons allowed in the 2p x orbital? subshells allowed in the 4 th shell? electrons allowed in the 3d subshell? electrons allowed in the 4d subshell? electrons allowed in the 3p subshell? electrons allowed in the 3 rd shell? Energies of Orbitals In hydrogen, all shells are equivalent in energy. Energies of Orbitals In many-electron models, the energy levels depend on the shell and subshell. Aufbau Principle Aufbau principle: start with the nucleus and empty orbitals, then build up the electron configuration using orbitals of increasing energy

7 37 38 Writing 39 Write electron configurations for the following atoms. H He Li Be B N O Ne Na Al S Ar K Sc Ti Zn Br Worked Ex. 5.10, 5.11; Problems 5.18, 5.19, 5.20 (Figure 5.17) Valence Electrons Electrons in the outermost shell. 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 5 Identify the valence electrons (v. e.) in the following configurations: 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 3 3s 2 1s 2 2s

8 Shorthand Notation Rather than writing out complete electron configurations, we can use the previously filled shell (noble gas) and show the valence electrons (v. e.): P: 1s 2 2s 2 2p 6 3s 2 3p 3 [Ne] 3s 2 3p 3 (5 v. e.) Write the shorthand notation for: Ca Cl Sr Fe Electron Configuration Some exceptions to the Aufbau order What are the expected electron configurations for Cr and Cu? Filled and half-filled filled d subshells seem to be especially stable. Cr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Also true for Mo and W Cu: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 e_config. Also true for Ag and Au 44 Orbital Diagrams Orbital diagrams are pictorial representations of electron configurations. Hund s Rule If two or more orbitals (i.e., a p or d orbital) with the same energy are available, one electron goes into each orbital until they have to pair up. Fighting sibling analogy For example, an atom with 2 p electrons: 1 electron will go into the first (p x ) orbital, the next electron will go into the second (p y ) orbital. 45 Pauli Exclusion Principle Pauli Exclusion Principle: no two electrons can have the same values of all 4 quantum numbers Describes what happens when electrons share an orbital. Only two electrons can occupy a single orbital and they must have opposite spin (i.e., the 4 th quantum number). The first electron is designated as positive spin (up arrow), the second electron in that orbital has negative spin (down arrow). 47 8