advertisement

Warm-Up Lithium has an atomic weight of 6.941 g/mol. When 10.4115 g of lithium is heated, it emits an energy measured at 262,500 joules. What is the energy given off by one atom? Wave Theory Particle Theory Light has measureable characteristics of frequency (ν) and wavelength (λ) When matter is heated, it emits radiation, the wavelength distribution of the radiation depends on the temperature. Planck showed that energy can only be released by atoms in “chunks” of some minimum size (quantum). E=hν Properties of Light Wave Theory Particle Theory Photoelectric effect: light shining on a clean metal surface causes the surface to emit electrons, but only if the frequency of the light is greater than some minimum frequency. Einstein showed that light comes in small energy packets called photons. The photon must have enough energy (E=hν) to overcome the attractive forces holding the electron within the metal atom. If the photon does not have enough energy (does not have the minimum frequency), no electron will be emitted. Emission Lines – Bohr Model of the atom Properties of Light To explain why energy states of an atom are quantized, scientists had to change the way they viewed the nature of the electron. The photoelectric effect, and hydrogen emission lines showed that light could behave as both a wave and a particle. What does this mean for the electron?? 4-2 The Quantum Model of the Atom Louis de Broglie – 1924 Suggested that electrons were waves confined to the space around the nucleus the waves could only exist at certain frequencies that corresponded to specific energies—the quantized energies of Bohr’s orbits (where E = hν). 4-2 The Quantum Model of the Atom A beam of electrons passed through a crystal shows the pattern to the right. Why is there a pattern of light and dark rings? What would you expect to see if you passed light through a tiny slit? Video Some more about waves The pattern is due to wave Interference and Diffraction. Wave Interference Interference is when two waves overlap (superpose) to form a resultant wave of greater or lesser magnitude. Constructive Interference – Destructive Interference – peaks add together and troughs add together resulting in a peaks add to troughs resulting in a smaller (or no) wave larger wave Wave Interference diffraction involves a change in direction of waves as they pass through an opening or around an obstacle in their path. As the waves bend around the obstacle, they experience interference (constructive and destructive) Wave Diffraction Electron beam diffraction experiments showed the bending of an electron wave as it passes by the edge of an object, such as an atom in a crystal. The resulting interference pattern occurs when waves overlap, resulting in a reduction of energy in some areas and an increase of energy in others. Further Evidence Shows That The Electron Also Acts As A Wave Diffraction shows that light behaves like a wave, and electrons also behave like a wave. Take Home Message… Bohr showed that electron orbits are quantized (only exist at certain frequencies). If we think of the electron wave as being like a wave that is confined to a box of a specific distance, the wave can only exist stably if it is showing harmonics. Why only certain frequencies? What Does The Wave Have To Fit? There is a restriction on the wavelengths that the electron can have. In successive revolutions, the waves must be exactly in phase with each other. Why? Oh Joy, Another Formula The wavelength of any orbital can be related to size of the orbit. How? 2πr = nλ; for n = 1, 2, 3, 4… This IS Why It’s Quantized Since each orbital can only have a specific wavelength associated with it, the energy of that orbital is fixed by the equation E = hν. The value of n gives the first quantum number, which indicates the main energy level occupied by the electron. Think 3-D, Not 2 The Heisenberg Uncertainty Principle (1927) It is impossible to determine simultaneously both the position and velocity of an electron or any other particle. Video Werner Heisenberg Think of the waveparticle duality of matter as this: A quantum entity, such as a photon or an electron travels as a wave, but arrives as a particle. Wave-Particle Duality Together with the uncertainty principle, the wave equation laid the foundation for modern quantum theory. Quantum theory describes mathematically the wave properties of electrons and other very small particles. The Schrödinger Wave Equation (1926) The wave functions give only the probability of finding an electron at a given place around the nucleus. Solutions to the equations = Wave Functions The electrons do not travel around the nucleus in neat orbits, as Bohr had thought. Instead, the electrons exist in certain regions called orbitals. An orbital is a three dimensional region around the nucleus that indicates the probable location of an electron. Orbitals In order to completely describe orbitals, quantum numbers are used. Quantum numbers specify the properties of atomic orbitals and the properties of electrons in orbitals. The first 3 quantum numbers result from solutions to the Schrodinger equation. Atomic Orbitals and Quantum Numbers Indicates the main energy level occupied by the electron. Values of n are positive integers (1, 2, 3…) As n increases, the electron’s energy and average distance from the nucleus increases. Principal Quantum Number (n) Indicates the shape of the orbital. For each n the number of orbital shapes possible is equal to n. The values of l that are allowed are zero and all positive integers less than or equal to n-1. Angular Momentum Quantum Number (l) l=0 designates an s orbital l=1 designates a p orbital l=2 designates a d orbital l=3 designates an f orbital Angular Momentum Quantum Number (l) Indicates the orientation of an orbital around the nucleus. The values allowed are m = -l……0…….+l Magnetic Quantum Number (m) Only two possible values, +½ or -½ Indicates the two fundamental spin states of an electron in an orbital. A single orbital can hold a maximum of two electrons, which must have opposite spins. Spin Quantum Number Improves upon Bohr’s model because it describes the arrangement of electrons in atoms other than hydrogen. The arrangement of electrons in an atom is known as the atom’s electron configuration. The Quantum Model Electrons in atoms tend to assume arrangements that have the lowest possible energies - ground state electron configuration The Quantum Model Some background info… To simplify things, we show electrons in orbits around the nucleus (it’s really 3-D orbitals). There can be a maximum of 8 electrons in the outer shell!!! ◦ But only the noble gases (Group 18) are full (ex. Neon) Rules Governing Electron Configurations The energy levels of the orbitals are determined and electrons are added to the orbitals one by one according to 3 basic rules: 1. Aufbau principle 2. Pauli exclusion principle 3. Hund’s rule Aufbau (“Build-up”) Principle An electron occupies the lowest-energy orbital that can receive it. Note: Starting with the third main energy level (n = 3), the energies of the sublevels begin to overlap (see diagram on p. 105). Pauli Exclusion Principle (1926) No two electrons in the same atom can have the same set of four quantum numbers. Thus electron pairs in orbitals must be of opposite spin. Wolfgang Pauli: Hund’s Rule Notice how the “3p” electrons are all up arrows before you start filling with down Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin. The rule above minimizes electronelectron repulsion. Orbitals are like clouds showing the most probable region of electron locations. The sizes and shapes of electron clouds depend on the energies of the electrons that occupy them. Orbitals are like clouds… n (principle quantum number) – the main energy level (shell or orbit) l (angular momentum QN) – the shape of each sublevel m (magnetic QN) – how many orbitals are in each sublevel Spin QN – direction of spin for an electron in an orbital Quantum Numbers – Like GPS Coordinates for an electron Orbital Notation 2. Electron-Configuration Notation 3. Noble-Gas Notation 4. By the quantum numbers associated with each electron (you will NOT be held responsible for this on a test). 1. Terminology that you need to know follows: Four Ways of Representing Electron Configurations Highest Occupied Energy Level The electron-containing main energy level with the highest principal quantum number. AKA: The outermost shell/orbit Inner-shell Electrons Electrons that are not in the highest occupied energy level. AKA: Everything that is NOT in the outer shell/orbit Orbital Notation An unoccupied orbital is represented by a line. The lines are labeled with the sublevel (principal quantum number and orbital shape). ◦ ◦ ◦ ◦ Why 1 line for s-orbitals? Why 3 lines for p orbitals? Why 5 lines for d orbitals? Why 7 lines for f orbitals? Draw and label the sublevels (see Fig 4-16 p. 105) Make sure your energy levels are spaced appropriately! Add electrons to the orbitals one by one following the 3 rules. Orbital Notation Draw the orbital notation for Boron First, determine how many electrons Boron has – 5 Next draw & label the sublevels 2p 2s 1s Orbital Notation - Example Draw the orbital notation for Nitrogen First, determine how many electrons Nitrogen has – 7 Next draw & label the sublevels 2p 2s 1s Orbital Notation - Example Draw the orbital notation for Oxygen First, determine how many electrons Oxygen has – 8 Next draw and label the sublevels 2p 2s 1s Orbital Notation - Example Draw the orbital notation for Magnesium First, determine how many electrons magnesium has – 12 Next draw and label the sublevels 3s 2p 2s Your turn…1s Don’t have to draw the lines and arrows The number of electrons in a sublevel is shown by adding a superscript to the sublevel designation Electron Configuration Notation Draw the electron configuration notation for Boron 2p 2p 2p 1s22s22p1 2s 1s Orbital Notation for Boron Electron Configuration Notation Draw the electron configuration notation for Nitrogen 2p 2p 2p 1s22s22p3 2s 1s Orbital Notation for Nitrogen Electron Configuration Notation Draw the electron configuration notation for Oxygen 2p 2p 2p 1s22s22p4 2s 1s Orbital Notation for Oxygen Electron Configuration Notation Draw the electron configuration notation for Magnesium 3s 2p 2p 2p 1s22s22p63s2 2s 1s Orbital Notation for Magnesium Electron Configuration Notation WARNING!!! The electron-configuration notation on your periodic table in the book is NOT in the proper order!!! Do not copy from the table, use the “yellow brick road”!!! The electron configuration for an atom is 1s22s22p63s23p5 How many electrons does it have? What is the atomic number? 17 What element is it? Chlorine In the 3rd main energy level, how many p orbitals are filled? 2 How many unpaired electrons does this atom have? 1 Example 17 Noble Gas Configuration The outer main energy level is fully occupied, in most cases, by eight electrons (sometimes called “completing the octet”). Which noble gas does not complete the octet? Noble Gas Notation This is a shorthand method to show electron configurations, usually for atoms in the 3rd period and beyond. You show the nearest previous noble gas as the core, and then only denote the outermost electrons. Weird Variations Chromium (Cr) expect [Ar]4s23d4, is [Ar]4s13d5 Copper (Cu) expect [Ar]4s23d9, is [Ar]4s13d10 The reason is that mixes of half and/or fully filled orbitals are more stable (in a lower-energy state), than a combination of fully or half filled orbital with a partially filled orbital. Assignment Electron-configuration worksheet