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Electrons in Atoms Chapter 5 ? Introduction • The atom evolved over the years as a result of new information that was collected. • The quantum mechanical model of the atom helped explain how electrons traveled around the nucleus and their energy. • Light and energy are related to each other and the movement of electrons. • Electrons are positioned in a particular way around the nucleus. Models of the Atom (Section 5.1) • The Development of Atomic Models • The Bohr Model • The Quantum Mechanical Model • Atomic Orbitals I.) The Development of Atomic Models • The atom ~ 1803 • Dalton’s atom • Small indivisible Deficiency: No subatomic structures • Represents the smallest part of matter that still has unique properties • The atom ~ 1897 • Thomson’s atom Deficiency: No nucleus and no movement of electrons • negatively charged electrons embedded on a positive charged sphere Deficiency: Could not explain chemical properties of elements such as… • The atom ~ 1911 • Rutherford’s atom • Small, dense, positively charged nucleus with electrons orbiting around it Hydrogen Helium Why heated objects glow red, then yellow, then white. Lithium Why elements give characteristic colors when heated. II.) The Bohr Model • A student of Rutherford’s • Believed Rutherford’s model needed to be improved to account for new discoveries. • Considered the simplest atom, the hydrogen atom. Niels Bohr Danish Physicist (1885 – 1962) • Bohr also proposed electrons orbit around nucleus • Electrons orbit in fixed energy levels. • Electrons can move up and down energy levels. • Energy is involved in this movement Quantized Energy Quantum of Energy: The amount of energy required to move an electron from one energy level to another. • The Bohr atom gave results that agreed with experiments for the hydrogen atom only. How does this energy relate to light and color? Light • Dual Nature of Light: Light can act like waves, and as straight line particles. • Light is one type of electromagnetic radiation (em), which is a form of energy that has wavelike behavior • Other types of em radiation are: x-rays, uv, infrared, microwaves & radio, and together they all form the Electromagnetic Spectrum Energy and Light The Wave The Electromagnetic Wave • The speed of all em waves through a vacuum (space) and through air is: Speed of light (c) = 3.0 x 108 m/s • The length of each individual wave is known as it’s wavelength () which is the length between corresponding points on adjacent waves, usually measured in nanometres (1nm = 10-9m) • The frequency (f or ) is how many waves pass a particular point in a second and is measured in waves/second = Hertz (Hz) The Light Wave Relationship Between Frequency & Wavelength The relationship between them is as follows: c=fx • What is the wavelength & color of light that has a frequency of 6 x 1014 Hz (1/s)? f = 6 x 1014 Hz (1/s) c = 3 x 108 m/s =? c=fx = c = 3 x 108 m/s = 5 x 10-7 m f 6 x 1014 1/s = 500 x 10-9 m = 500 nm ( = green ) Energy, Wavelength, & Frequency • When an object gets hot, it emits energy in small specific amounts called quanta. • A quantum is the minimum amount of energy that can be lost or gained by an atom. • The frequency determines the Energy… The frequency determines the Energy by: E = h.f or E = h (c/λ) h = 6.626 x 10-34Js (Planck’s constant) What is the Energy of Green light with a frequency of 6 x 1014 Hz ? f= 6 x 1014 Hz (1/s) E = h.f = 6.626 x 10-34 J.s (6 x 1014 Hz) = 3.98 x 10-19 J The Particle Nature of Light • • • • It was Einstein that proposed that light acted like a stream of particles called photons. A photon is a particle of em radiation that has zero mass and carrying a quantum of energy There must be a minimum amount of energy to eject electrons from a metal Using E = h.f, Planck realized that the em radiation providing the energy, must be of a certain frequency. Different metals need different minimum amounts of energy, and therefore frequencies • When the electron falls back to its ground state or a lower energy state, it emits a photon of radiation with a specific amount of energy, and not continuous amounts of energy. • Therefore, specific frequencies of light are emitted when excited hydrogen “cools” down. • When the H atom gets excited, the e- jumps to another specific orbit and not somewhere in between – like going up a ladder • When the e- falls back, it loses energy in the form of a photon (a bundle of light energy) Convert circular orbits to lines Energy Levels of an Atom The Absorption & Emission Spectrum Emission Spectrum: The pattern of discrete lines resulting from the frequencies of light emitted by an element that has been energized. • The emission spectra represents the movement of an electron from higher energy levels down to lower energy levels. • This spectrum is for hydrogen. Atomic Spectra Hydrogen Helium Mercury Argon Iodine Neon Explanation of the Atomic Spectra • Each discrete line in an emission spectrum corresponds to one exact frequency of light emitted by the atom. • Emission spectrum, like a person’s fingerprint, can be used to ID an element. • As we saw earlier, the light emitted by an electron moving from a higher to a lower energy level has a frequency and wavelength directly proportional to the energy change of the electron. The Bohr/ Rutherford Planetary Orbital Model • Based on describing paths of moving electrons as you would describe the path of large moving objects. What was to be done? • Data collected and theoretical calculations showed that electrons did not move in this way for larger atoms. III.) The Quantum Mechanical Model of the Atom • Comes from mathematical solution to a complex equation. • Based on probabilities. • Electrons occupy certain spaces around nucleus. • Used theoretical calculations and “new” experimental results to devise a mathematical equation describing the behavior of the electron in a hydrogen atom. • The quantum mechanical model of the atom comes from the mathematical solution of this equation. • Solution restricts the energy of electrons to certain energy levels. Erwin Schrodinger Austrian Physicist 1887 - 1961 • Solution does not involve an exact path the electrons travel around the nucleus. Summarizing the Findings of Quantum Mechanics • Determines the allowed energies an electron can have & how likely it is to find the electron in various locations around the nucleus. • The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented by a fuzzy cloud. The cloud is more dense where the probability of finding an electron is high. The cloud is less dense where the probability of finding an electron is low. Summary of Atomic Orbitals • Defined as a region of space in which there is a high probability of finding an electron. • Each energy level may possess several orbitals with different shapes and at different energy levels. • These energy levels are deemed “sublevels” and corresponds to an orbital. The Atomic Orbitals The s Orbital (1shape) Each energy sublevel corresponds to an orbital of a different shape. The p Orbitals (3 shapes) The d orbitals (5 shapes) Question is: Why can’t e-’s be in an orbit between the specific Energy levels? French scientist Louis DeBroglie pointed out that the electron orbits acted like the behavior of waves. (i.e. you can only have a certain amount of waves in a given distance/space – not ½ of a wave) If you can only have specific # of waves, then based on the formula c = f x you can only have specific f’s, which in turn translates into specific Energy levels. Each energy level has energy sublevels that correspond to an atomic orbital. Principle Energy Level Number of Sublevels Type of Sublevel n=1 1 1s (1 orbital) N=2 2 2s (1 orbital) 2p (3 orbitals) N=3 3 N=4 4 3s (1 orbital) 3p (3 orbitals) 3d (5 orbitals) 4s (1 orbital) 4p (3 orbitals) 4d (5 orbitals) 4f (7 orbitals) 3d 3p 3s n=3 2p n=2 2s n=1 1s Describing Energy Sublevels • The number of sublevels per energy level corresponds to the energy level number. • There is much overlap of energy sublevels. Fill in the Following Chart Principle Energy Level n=1 n=2 n=3 Number of Sublevels Type of Sublevel • There is a set number of orbitals per energy sublevel that corresponds to the number of shapes each orbital possesses. •s=1 •p=3 •d=5 •f= 7 • Only 2 e- per orbital max Electron Configuration in Atoms (Section 5.2) • Electron Configurations • Exceptional Electron Configurations I.) Electron Configuration • The way in which the electrons of an atom are arranged in various orbitals around the nuclei of atoms. • 3 basic rules allow us to determine this: the aufbau principle, the Pauli exclusion principle, & Hund’s rule. • Use these rules to place all the electrons of an atom into the different orbitals. The 3 Rules for Electron Configurations The Aufbau Principle • Electrons occupy the orbitals of lowest energy first. Pauli Exclusion Principle • The maximum number of electrons per orbital is two. • When two electrons share an orbital they must have the opposite spins (i.e. electron spins must be paired). Hund’s Rule • 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 Aufbau Principle Electrons occupy the orbitals of lowest energy first. The filling of atomic orbitals is not simple beyond the 2nd energy level. Pauli Exclusion Principle • The maximum number of electrons per orbital is two. • When two electrons share an orbital they must have the opposite spins (i.e. electron spins must be paired). ↑↓ Hund’s Rule 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. Electrons repel and do not like being around each other, and so if they can they will separate. 2 p orbitals The Maximum # of Electrons That Can Occupy a Principle Energy Level # e = 2 2n – n = the principle energy level Let’s Start Write the electron configuration for phosphorus. (Atomic # = 15) How many total electrons? Where do the electrons go first? Where do the others go? Let’s Practice Write the electron configuration for carbon. (Atomic # = 6) Write the electron configuration for argon. How many total electrons? Write the electron configuration for nickel. How many total electrons? Shorthand Method for Electron Configuration This method involves writing the energy level and the symbol for every sublevel occupied by an electron. • We indicate the number of electrons occupying a particular sublevel with a superscript. • When writing shorthand, keep sublevels of a given energy level together, even though this distorts what is obtained by following the aufbau principle. Let’s try this with our previous examples. Write the electron configuration for the following atoms using the shorthand method. • Phosphorus • Carbon • Argon • Nickel Simple Table to Help Write Electron Configurations Shorthand Use the table to write the electron configuration for potassium. Let’s Practice 1. 2. 3. 4. 5. Write the electron configuration for the following elements using the shorthand method. magnesium nitrogen chlorine silicon calcium Identifying Elements from Electron Configurations 1. Count the total number of electrons. 2. This is the atomic number. 3. Use this number to ID element from the periodic table. 4. Example: 1s22s22p63s1 Let’s Practice Give the symbol and name of the elements that correspond to the following electron configuration. 1.) 1s22s22p63s1 2.) 1s22s22p3 Electrons in Atoms Chapter 5 The End ?