Modern Atomic Theory

Transcription

1 Modern Atomic Theory In science, often times chemical or physical behavior can not be seen with the naked eye (nor with the use of some other device). Consequently, an understanding and explanation of the chemical or physical behavior must be determined indirectly. From these indirect means, several models of the atom have been presented. Atomic Models of the Atom Democritus Model Indivisible particle Dalton's Model Indivisible particles called atoms, etc. 1

2 J. J. Thomson s Model Atoms composed of negatively electrons and positively charged protons located randomly throughout the atom. Rutherford s Model He proposed that almost all of the mass (about 99.99%) and ALL of the positive charge of the atom is located at the center of the atom - NUCLEUS Since the protons and neutrons have a much greater mass than the electron, the mass is concentrated in a tiny nucleus. The nucleus occupies a very small part of the volume of the atom. The rest of the space is occupied by the electrons which are very small, do not take up much space. 2

3 Rutherford s model, although a major step forward in the understanding in the structure of the atom, left some important questions unanswered. 1. What were the electrons doing? 2. How are the electrons arranged in the atom? Randomly? In any specific order? 3. How did they electrons move? Like planets orbiting a sun? 4. Why aren t the electrons attracted to the nucleus of the atom? Why don t the atoms annihilate themselves? In order to answer these questions, additional work needed to be done in the areas of the nature of light and the transmission of energy. 3

4 Electromagnetic Radiation What do the following all have in common? 1. Heat coming from a hot oven. 2. Light emitted from a light bulb. 3. Light from the sun. They all involve energy being transmitted from one place to another. More correctly, they all involve ELECTROMAGNETIC RADIATION. Many forms of electromagnetic radiation exist. 4

5 How does electromagnetic radiation get transmitted? By Waves! What do waves look like? All waves can be described by 4 characteristics 1. Wavelength (distance from peak to peak) 2. Frequency (how many cycle the wave has for a specific period of time) 3. Speed (how fast the wave travels) 4. Amplitude (height or intensity of the wave) What makes any 2 forms of electromagnetic radiation different from each other is their wavelength and frequency. 5

6 The Components of a Wave Peak (above) & Valley (below) center line. Amplitude = height of a wave from the center line to a Peak. Wavelength ( - the Greek letter lambda) = The distance between peaks. ( with units nm or m) Frequency ( - the Greek letter nu) = The number of times that a wave passes through a given point at a given time. ( with units 1/s or s -1 or hertz) Speed (c AKA Speed of Light) = How fast the electromagnetic radiation travels it is the same for all forms. c = 3 x 10 8 m/s (approx=670,000,000 mile/hr) 6

7 A mathematical relationship exist between the wavelength and the frequency of a wave: c What this equation tells us is that: Large Small have Small have Large A mathematical relationship also exist between the Energy and the frequency of a wave: E (where h = x J s) What this equation tells us is that: h Large Small have Large have Small 7

8 Of all forms of electromagnetic radiation that exist, we only notice 2 of them: Visible Light (with our eyes) & Infrared (with our skin) However, we are either affected by or utilize all of them! It turns out that electromagnetic radiation not only can behave as a wave, but it can also behave as a particle. Albert Einstein described these little particles as tiny packets of energy call PHOTONS. Ultimately, it has been accepted that electromagnetic radiation consists of waves of oscillating photons traveling everywhere. Sunlight (and white light) consists of waves with a continuous range of wavelengths and frequencies in the visible light portion of the electromagnetic spectrum. When sunlight (and white light) passes through a prism, the light separates into a spectrum of colors (or colors of the rainbow). 8

9 Energy Colors of the Rainbow: ROYGBV R = Red O = Orange Y = Yellow G = Green B = Blue V = Violet Atom s Emitting Energy It had been known for some time that when atoms were given energy (by burning them for example) they would emit light. The diagram below helps to explain the process: Excited Atom Atom Given Energy Photon Emitted - Visible Light Ground State Atom Atom Returns to Ground State Color of light emitted dependant on the frequency & wavelength of the light. 9

10 Emission Spectra of Hydrogen Let s see what happens when electricity is passed through a gas tube that contains hydrogen gas. What color do you see? Let s use a spectroscope (a prism connected to a scale that can measure wavelengths) to separate the components of the light. What colors do you see? Color Red Blue-Green Blue Violet Wavelength (meters) 6.56 x 10-7 m or 656 nm 4.86 x 10-7 m or 486 nm 4.34 x 10-7 m or 434 nm 4.10 x 10-7 m or 410 nm Recall: 1 nm = 1 x 10-9 m Which line has the highest energy? Which line has the lowest energy? Why ONLY these lines??? 10

11 Why not every color of the rainbow? Specific colors must correspond to specific wavelengths. Specific wavelengths must correspond to specific energies. These specific colors suggest that hydrogen must only have certain allowable, discrete energy levels allowed! They never emit light at wavelengths between those emitted. We say that the hydrogen atom must be QUANTIZED. Only certain energy values or levels are allowed. It was later found that all atoms are quantized. 11

12 Try to identify which colors of light correspond to the energy level changes depicted in the diagram: B Excited State 4 Excited State 3 Excited State 2 Energy C Excited State 1 A D Ground State 12

13 Bohr Model of the Atom In 1913, Niels Bohr ( ) a Danish physicist, and former student of Rutherford, came up with a more advanced model of the atom. Niels Bohr agreed with many parts of Ernest Rutherford s model of the atom. For example: 1. The idea of the nucleus 2. The electrons not being in the nucleus He also needed to incorporate the idea that the hydrogen atom was quantized. The big problem with Rutherford's model of the atom was what would happen if 2 oppositely charged particles (protons & electrons) are close to each other? They attract to each other!!! 13

14 Therefore, Bohr suggested that: If protons & electrons were simply sitting randomly within the atom, they would attract and smash into each other. The atom would annihilate itself. Bohr came up with a new picture of the atom. Electrons were arranged in the atom in concentric, circular paths or orbits, around the nucleus. Nucleus + Possible Orbits These orbits were specific distances for the nucleus of the atom. He suggested that when a hydrogen atom was excited, an electron would move to an orbit farther from the nucleus. 14

15 This transition would take energy a specific amount of energy. When the atom return to its ground state, photons would be emitted. This was consistent with the observations made for hydrogen when it was excited. 15

16 Modern Model of the Atom Bohr's model of the atom, which contained orbits in which electrons moved about, explained the Hydrogen atom very well. The emission could be predicted (i.e. the frequency, the wavelength and the energy). The Bohr model DID NOT work very well for atoms that had more than 1 electron. The more electrons in an atom, the worse the model got in explaining the atom! The emission could NOT be predicted (i.e. the frequency, the wavelength and the energy). As a result, a NEW theory needs to be developed. Two ideas in physics paved the way. 1. In 1924, Louis de Broglie (French) developed a mathematical relationship that supported the fact that the photons of light have wavelike properties. 16

17 2. In 1926, Erwin Schrödinger ( ), an Austrian physicist developed equations for the energy of electrons in an atom which describes the properties of an electron as wave-like. Not only do Schrödinger's equations describe the electrons in the atom, but they also describe the behavior of waves in the electromagnetic spectrum. The work of de Broglie and Schrödinger was so important that a branch of physics developed that studies the quantized energy changes of electrons in atoms - Quantum Mechanics. The modern picture of the atom comes from the answers to Schrödinger equations (very complicated). 17

18 The modern picture of the atom is called the Quantum Mechanical Model of the Atom. A representation of the hydrogen atom is shown here: 1. Like the Bohr model, the quantum mechanical model restricts the energy of the electrons to certain values. 2. The path of the electron around the atom is NOT an exact path or orbit. Instead, it estimates the probability of finding an electron in a specific region. These regions are called ORBITALS. 3. The mathematics that comes from this model does not tell us of the exact motion of the electron or the exact time when it is in a particular orbital. 18

19 Orbitals and Energy Levels The plot of the probability map is represented as an orbital. Chemists have defined the 3D space around the nucleus of the hydrogen in which the electron occupies 90% of the time as the orbital. Always keep in mind that the electron occupies the space within the surface NOT on the surface itself. When a hydrogen atom is in its lowest energy state, it is called the ground state. The probability of finding the electron in the lowest energy orbital is very high. If the hydrogen atom absorbs energy, the electron moves to a higher energy level (perhaps the 2nd, 3rd or higher energy level) with contain different orbitals. The atom is then said to be in an excited state. 19

20 We say that when an electron is in an orbital where the electron is in its the 1st energy level, it has a principal energy number, n = 1. Principle Energy Number n 1st energy level n = 1 2nd energy level n = 2 3rd energy level n = 3 4th energy level n = 4...and so on. Unlike the Bohr model of the atom, each energy level may have more than one orbital type. The number of orbitals for each energy level can be determined by a simple equation: # of Orbitals = n 2 where n = energy level 20

21 Orbital of the Hydrogen Atom Energy Level # of Orbitals n n Types of Orbitals (Sublevels) There are several types of orbitals that exist in an atom. All atoms have all types of orbitals, however, they may not have any electrons in them. The types of orbitals depend on the energy level. Orbital Type Number of Orbitals s 1 p 3 d 5 f 7 21

22 The orbital types are s, p, d, and f s - 1st orbital of an energy level p - 2nd orbital of an energy level, if possible 22

23 d - 3rd orbital of an energy level, if possible f - 4th orbital of an energy level, if possible 23

24 Increasing Energy When put all of these pieces together, the atom is series of layers of energy levels. Each energy level can have several orbitals. Together it looks like the chart below (shown are 4 of the 7 energy levels known.) The Energy Levels and Orbitals of the Hydrogen Atom n = n = 4 4s 4p 4p 4p 4d 4d 4d 4d 4d 4f 4f 4f 4f 4f 4f 4f n = 3 3s 3p 3p 3p 3d 3d 3d 3d 3d n = 2 2s 2p 2p 2p n = 1 1s Do you recall the emission for Neon? Why was it so complicated? If hydrogen only has 1 electron that can be excited, how many does neon have??? Which one of them got excited? 24

25 We need a better way of figuring out how the electrons of an atom are arranged. The next section will do this for us. Electron Configuration What do we know about how the electrons are arranged in the atom? 1. Atoms have ENERGY LEVELS. 2. These energy levels correspond to regions of space called ORBITALS. 3. We know that their 4 types of orbitals (sublevels - s, p, d, f). LADIES and GENTLEMEN, we are now ready to produce a realistic picture of what the atom looks like. How are the electrons arranged within the atom? 25

26 The Order of Filling of Electrons in the Atom So, the order of filling is: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p Recall the emitted light for hydrogen gas in the experiment that you performed in the laboratory. The electron in the ground state energy level was excited to a higher, excited energy level. 26

27 Well, as with everything in nature, electrons do not like to work if we do not need to. Aufbau Principle Electrons enter orbitals of lowest energy first. As a result, electrons will occupy orbitals which require the least amount of energy first. When those orbitals are full, they will occupy orbitals of higher energy next. Pauli Exclusion Principle Electrons rather be by themselves than share an orbital with another electron. But if they must share an orbital with another electron, the 2 electrons must have opposite spin. Hund's Rule When electrons can occupy multiple orbitals of equal energy, one electron enters each orbital until all the orbital contain one electron with parallel (same) spins. If more electrons need to be added, these electrons will pair with the opposite spin. 27

28 Electron Configuration and Orbital Diagrams Orbital Diagram A shorthand way of showing the distribution of electrons in every orbital. Electron Configuration A shorthand way of describing the location and number of electrons that an atom possesses. Orbital Notation Electron Configuration H He Li Be B C N O 1s 2s 2p 2p 2p 1s 2s 2p 2p 2p 1s 2s 2p 2p 2p 1s 2s 2p 2p 2p 1s 2s 2p 2p 2p 1s 2s 2p 2p 2p 1s 2s 2p 2p 2p 1s 2s 2p 2p 2p 1s 1 1s 2 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p 2 1s 2 2s 2 2p 3 1s 2 2s 2 2p 4 F Ne 1s 2s 2p 2p 2p 1s 2 2s 2 2p 6 1s 2s 2p 2p 2p 28 1s 2 2s 2 2p 5

29 Let s write orbital notations and electron configuration for the following: Orbital Not. Electron Config. Mg Si S Ar Sc Mn Cu Electron Configuration Shortcut When writing electron configurations for atoms that have many electrons, a long row orbitals result. A shorter way of writing the electron configuration exist. 29

30 For the element magnesium, electron configuration = 1s 2 2s 2 2p 6 3s 2 The 1s 2 2s 2 2p 6 part of the electron configuration is the same as that of Ne. So, 1s 2 2s 2 2p 6 = the electron configuration Ne. THE SHORTCUT Replace 1s 2 2s 2 2p 6 with [Ne] This must the last completely filled energy level.* The new way to write the electron configuration for Mg is [Ne]3s 2 Try the shortcut method for Tin (Sn) 30

31 Electron Configuration of Ions The electron configuration of fluorine is 1s 2 2s 2 2p 5 Does fluorine have a charge? NO. What would be the electron config. of F 1-? F 1s 2s 2p 2p 2p 1s 2 2s 2 2p 5 F 1-1s 2 2s 2 2p 6 1s 2s 2p 2p 2p The electron configurations that we have been writing have been for NEUTRAL atoms. 31

32 Let's try to write the electron configurations ions. 1. O 2-2. C Li Ti Fe 2+ 32

33 Atomic Structure and Periodic Trends The whole of point of learning chemistry is to try and understand why the physical world that we live in behaves the way that it does. I hope that it is obvious to you by now that matter behaves the way that it does because of the atoms that make it up. Understanding the internal structure of atoms helps us to explain why atoms do what they do. Also, developing an understanding of this structure allows use to predict how other atoms (and the materials that they make) act. Based on similarities in electron configuration, one can predict the physical and chemical behavior of an element not even ever work with it before. That s powerful!!! For example, let s consider the transition metals in group 11 (Cu, Ag, and Au). All 3 atoms have the same ending to their electron configuration s 2 d 9. 33

34 Cu oxidizes (rusts) very easily. Ag oxidizes but not as easily as Cu. Based on periodic trends, one could conclude that Au is least oxidized. You would be correct! Two very important periodic trends that we will discuss are Ionization Energy and Atomic Size. Ionization Energy The amount of energy need to remove the outermost electron from an individual atom in the gas phase. A(g) A 1+ (g) + e - In general, the stronger the nucleus of an atom attracts (holds) its outermost electron, the more difficult it will be to remove that electron. The more electrons that exist between the nucleus and the outermost electron, the less the attractive force there will be. This effect is called Shielding. 34

35 This effect is really significant when comparing elements with a group. As an example, here are the trends in ionization energy of two groups on the periodic table: Be > Mg > Ca > Sr > Ba O > S > Se > Te Why would any atom ever lose an electron? What would be the advantage? The main reason STABILITY. All atoms want it! How do atoms obtain stability? By having a completely fill outermost energy level. Now for metals, they generally only need to lose one or two electrons to develop a completely fill outermost energy level. That is way metals tend to have a very low ionization energy compared to nonmetals. 35

36 Ionization Energy Decreases Nonmetals would rather gain electrons than lose them. It is very hard to remove electrons from metals. In general, the following trend exist for ionization energy: Ionization Energy Increases 36

37 Atomic Size What general statements could you make about the size (or radius) of atoms across a period and down a group on the periodic table? 37

38 1. Atomic size increases down a group. Why? Each successive period has electrons in energy levels that farther away from the nucleus of the atom. 2. Atomic size decreases across a period. Why? As one goes across the period, the energy level does not change. However, the number of protons in the nucleus increases. This successive increase in the number of protons makes the attractive ability of the nucleus stronger. This stronger attraction allows the nucleus to pull in all of its electrons closer to it. The Result: The atoms are progressively smaller across the period. 38