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How did we arrive at this picture of the atom? Rutherford’s Atom Lord Kelvin The entire atom is like a positively charged pudding with negatively charged electrons dispersed throughout like raisins. This view lasted until 1910. Earnest Rutherford Used a Particle Accelerator to discover that the majority of the space in an atom is empty. › The nucleus is a concentrated mass in the center and is made up of positively charged protons. › The negative particles, named electrons by JJ Thomson, are located all around the nucleus. This view was introduced in 1919. James Chadwick Working with Rutherford, he discovered that there was excess mass in the nucleus that was not accounted for by the protons. He named these neutral, no-charge particles neutrons. BUT!! We still don’t know what the electrons are doing! Electromagnetic Radiation is several types of energy that exhibit wavelike behavior and travel at the speed of light in a vacuum. It was eventually discovered that electrons behave more like waves than like particles, perhaps because of their incredibly small mass… 9.10938188 × 10-31 kg Energy travels in the form of waves. Waves have special properties… 1. Wavelength 2. Frequency 3. Speed 1. Wavelength – the distance between two consecutive wave peaks or troughs. › The symbol for wavelength is λ, the Greek letter lambda. 2. Frequency – the number of waves (cycles) per second that pass a given point in space. › The symbol for frequency is ν, the Greek letter nu. 3. Speed – All types of electromagnetic radiation travel at the speed of light (c = 2.9979 x 108 m/s). › The speed of light is the product of the wavelength and frequency. c = λ ν › Wavelength and Frequency have an inverse relationship. Remember ROIGBIV? Light travels in waves of particles called photons which can be viewed as little packets of energy. › Photons can travel at different frequencies, and these frequencies correspond to light of different colors. Visible light includes light of all frequencies, and appears white to us. › When photons of certain frequencies are isolated, we see light in the various colors of the rainbow. Of the visible spectrum, violet light has the most energy, because it travels with the greatest frequency. Some light travels with higher energies than we can see. Examples are listed in order of increasing energy. 1. 2. 3. Ultraviolet Light X-rays Gamma Rays Some light travels with lower energies than we can see. Examples are listed in order of decreasing energy. 1. 2. 3. Infrared Microwaves Radio waves (FM, shortwave, AM) What happens to wavelength as frequency increases? What happens to the speed? What is the wavelength range for the visible light spectrum? What is the frequency of red and purple light? When atoms receive energy from some source, they become excited. Eventually, they release this energy by emitting photons. If the photons have wavelengths within the visible light spectrum, we can see them as light. As atoms take in energy from the heat of the flame, they become excited. › What’s actually happening is the atom absorbs energy, which allows its electrons to temporarily jump up levels, and occupy higher energy positions farther away from the nucleus. The atom can only hold this extra energy in for so long before it emits the energy in the form of energy packets, or photons of light. › As electrons return to their original positions, the atom returns to ground state, and it loses its extra energy by shooting off a photon. The light signature of each metal has been determined by the use of a special tool called a spectroscope. › The spectroscope splits the light of the visible spectrum into distinct lines of light that are separated in a way that our eyes cannot achieve unaided. Because of this research, we can identify the metal ions in an unknown solution by the color they emit when excited. Two Scientists are credited with the discovery of the true nature of light and energy. Max Planck & Albert Einstein discovered that energy is quantized…in other words, only certain values are allowed. Matter and Energy were considered to be separate entities. Matter had particles… energy did not. › The location of energy in space was considered indeterminable. It was thought that matter could absorb or admit any quantity of energy. Different wavelengths of light carry different amounts of energy per photon. › The energy contained in the photon corresponds to the change in energy that the atom experiences when an electron goes from an excited state to a lower state. (NOT necessarily the ground state.) Because the atom can exist only in specific energy states, it can contain only an amount of energy equal to the amount necessary to attain an energy level, not any amount in between levels. › We refer to this property as the quantized nature of energy. › Quantized – only certain values are allowed. › Think of stairs compared to a ramp… The photoelectric effect is the phenomenon in which electrons are emitted from the surface of a metal when light strikes it. Only light higher than a certain energy will cause electrons to be emitted. Einstein won the Nobel Prize in 1921 for his study of the photoelectric effect. 1. Studies in which the frequency of the light is varied show that no electrons are emitted by a given metal below a specific threshold frequency, v0. 2. For light with frequency lower than the threshold frequency, no electrons are emitted regardless of the intensity of the light. 3. For light with frequency greater than the threshold frequency, the number of electrons emitted increases with the intensity of the light. 4. For light with frequency greater than the threshold frequency, the kinetic energy of the emitted electrons increases linearly with the frequency of the light. Energy is quantized. It can occur only in discrete units called quanta. Electromagnetic radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well. It was later found that matter exhibits a dual nature as well. Electrons, which are considered to be particles also exhibit wavelike properties. › Excited ions emit photons of light which have a frequency which corresponds to the energy change that occurs when an electron jumps from an excited state to the ground state. There is a size requirement for the dual nature to be easily observed. › Particles that are too big behave like particles. › Particles that are too small behave like waves. The Important Point is that matter and energy are not distinct. When white light is dispersed, it displays a continuous spectrum of visible light, like that seen below. Why are there 2 rainbows? When atoms are excited, they give off of light (flame test lab), This light can also be dispersed. The spectrum that is produced, while not continuous like a rainbow, is unique to that element. Continuous Spectrum › a spectrum that contains ALL the wavelengths of visible light. › All the lines are present. Line Spectrum › a spectrum that contains only certain wavelengths of visible light. › Only some of the lines of are present. Hydrogen is a good place to start because it only has one electron. › Therefore, all the light it emits comes from the same source. Hydrogen emits light with a four-line visible light spectrum. › The wavelengths of the light are 410 nm, 434 nm, 486 nm, & 656 nm. The significance of the existence of atomic spectrum of hydrogen is that it indicates that only certain energies are allowed for the electron in the hydrogen atom. Pages 287 –288 Niels Bohr (1885 – 1962) was a Danish physicist. He developed a quantum model for the Hydrogen atom. He proposed that the electron in the hydrogen atom moves around the nucleus only in certain allowed circular orbits. In order for an electron to move between levels, it must gain or lose a specific quantity of energy. The wavelengths predicted matched those that were measured. The model correctly fits the quantized energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electrons. Electrons require energy to exist in orbits farther away from the nucleus because of the attraction of opposite charges. As the electron is brought closer to the nucleus, (like when it jumps from an excited state to ground state) energy is released from the system in the form of a photon.. An electric current is sent through a tube containing H2(g). The energy causes the molecules to separate and fly around in a high energy state. The atoms collide and recombine, excited electrons returning to lower energy states & releasing light energy. The light is split by a diffraction grating into the line spectrum of hydrogen. The distance to the lines in the line spectrum is measured using a spectroscope. The wavelengths of the corresponding photons given off when the electrons release energy are determined. The wavelengths are used to determine which quantum levels the electrons jump from in order to give off each respective color line in the spectrum. Mid 1920s Werner Heisenberg (1901-1976) Louis de Broglie (1892-1987) Erwin Schrodinger (1887-1061) Compares the electron bound to the nucleus to a standing wave. Bohr’s model was known to be incorrect by the mid 1920’s. The motion of a large particle (like a billiard ball) is predictable. Electrons, however, do not move predictably. › De Broglie proposed that the smaller a particle is, the more it behaves like a wave (light), and the less it behaves like a particle (billiard ball). The unpredictable position of an electron is the main reason why scientists knew that Bohr’s model was incorrect. › According to his model, we should know exactly where the electrons are…but we can’t find them! But how then, can we determine the position of a particle, if we do not understand how it moves? De Broglie and Schrodinger came up with a VERY complex mathematical equation that was able to predict the location of an electron based on the wave-particle theory. The result is our modern concept of the atom containing not only orbits, but orbitals. › Orbital – The space around the nucleus where an electron is MOST LIKELY TO BE. There are four types of orbitals and each has its own specific energy, shape, & # of electrons it can hold. › s, p, d, & f It is important to note that the concept of an orbital is not a method for finding an exact location of an electron. Orbitals are just a prediction based on statistics and probability. (See next slide) Electrons do not follow a direct and predictable path. (the firefly experiment) Atoms have 2 subdivisions for housing electrons. › Principal energy levels (like orbits) › Sublevels (the orbitals within each orbit) It is somewhere in its orbital, following an indeterminable path, in a position which can only be determined by probability! The definition most often used by chemists to describe the size of the hydrogen 1s orbital is the radius of the sphere that encloses 90% of the total electron probability. Principle Energy Levels are represented by the numbers 1-7. The total number of principle energy levels an atom has corresponds directly to the row in which the atom is found on the periodic table. › For example, Silver, which is found in row 5 has five principle energy levels (1,2,3,4 & 5). The higher in number a principle energy level is, the farther away it is from the nucleus. › Which means an electron must have higher energy in order to be there. There are four sublevels. › s, p, d, f Each principle energy level can contain a number of sublevels (orbitals) equal to or less than its principle energy level number up to a maximum of 4. › PEL 1 can contain only an s orbital. › PEL 3 contains s, p & d. › PEL 6 contains s, p, d & f. Only one of each sublevel is contained in each principal level, and once these are full, electrons must move up to the next energy level. The s Orbital as a spherical shape. The larger the value of n, the larger the sphere. In the picture, 1s and 2s are shown as white rings. › The blue center represents the nucleus. › The pink area between the 1s and 2s orbitals is an area of zero probability where no electrons exist. The p Orbital has two lobes separated by a node at the nucleus. The p Orbital is only possible for atoms with principle energy levels ≥ 2. The larger the value of n, the larger the lobes. There are three possible orientations in space. › 2px › 2py › 2pz Which p orbital is pictured? The d Orbital has an even more complex shape than p. The d Orbital is only possible for atoms with principle energy levels ≥ 3. The larger the value of n, the larger the shapes. There are five possible orientations in space. › 3dxz › 3dyz › 3dxy › 3dx2-y2 › 3dz2 The f Orbital has the most complex shape. The f Orbital is only possible for atoms with principle energy levels ≥ 4. The larger the value of n, the larger the shapes. There are seven possible orientations in space. › 4fz3 – (3/5) zr2 › 4fx3 – (3/5) xr2 › 4fy3 – (3/5) yr2 › 4fxyz › 4fy(x2 – z2) › 4fx(z2-y2) › 4fz(x2 –y2) Pauli Exclusion Principle states that an atomic orbital can hold a maximum of two electrons, and those two electrons must have opposite spins. Electrons spin much like tops. › They can spin in one of two directions, clockwise and counterclockwise. › We represent the fact that they spin in different directions by using arrows: ↑ & ↓ It is not important to know which way an electron is spinning, just that if there are two in one orbital, they MUST have opposite spin. n Orbitals Orientations Total # of Electrons 1 s 1 2 2 s, p 1, 3 2+6=8 3 s, p, d 1, 3, 5 2 + 6 + 10 = 18 4 s, p, d, f 1, 3, 5, 7 2 + 6 + 10 + 14 = 32 1s2 2s2, 2p6 3s2, 3p6, 3d10 4s2, 4p6, 4d10, 4f14 Electron Configuration – electron arrangement. There are two ways to represent electron configuration: › Electron configuration notation – Uses the principal energy level, the orbital shape, and the number of electrons present. Hydrogen: 1s1 › Orbital Diagram – Uses the principal energy level and the orbital shape plus a series of boxes labeled with arrows to represent electrons present as well as their spins. Hydrogen: 1s ↑ As atoms increase in # of protons and electrons, it can become confusing as to where the additional electrons go. › The electrons fill up atomic orbitals in a specific order. The periodic table of elements can be used to determine this order. › Lets consider the first 18 elements on the periodic table… › http://www.dayah.com/periodic/ Orbitals fill up in the following order: › 1s › 2s, 2p › 3s, 3p › 4s, 3d, 4p › 5s, 4d, 5p › 6s, 4f, 5d, 6p › 7s, 5f, 6d, 7p Do you notice any patterns? Energy Order of increasing energy is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s, Pages 303 - 307 Electrons in the outer most principle quantum level of an atom. These are the most important electrons to chemists because these are the electrons that are involved in bonding. The elements in the same group (vertical column of the periodic table) have the same valence electron configuration. › This is why Mendeleev was able to align the elements by their properties; because elements with = # of valence electrons behave in chemically similar ways. The quantum mechanical model can be used to explain the arrangement of the elements in the periodic table. Quantum mechanics allows us to understand that similar chemistry exhibited by the members of a given group arises from the fact that they all have the same valence electron configuration. Only the principle quantum number changes going down a group on the periodic table. The Wave Mechanical Model is further supported by several properties, which are directly related to electron configuration, that demonstrate a periodic repetition on the Table of Elements. We will talk about four: › Metals, Nonmetals, & Metalloids › Chemical Reactivity › Ionization Energies › Atomic Size Metals › Metals are ductile and malleable, can conduct heat and electricity, and have a shiny luster. › Metals give up electrons to form positive ions. › The majority of the periodic table, from the staircase to the left. Nonmetals › Do not have the properties of metals. › Nonmetals gain electrons to form negative ions. › The upper right corner of the periodic table to the right of the staircase. Metalloids › Exhibit the properties of both metals and nonmetals. › Include only 7 elements: Si, Ge, As, Sb, Te, Po, & At Electronegativity is the property of an atomic nucleus to attract electrons. › Electronegativity increases going up a group. › Electronegativity increases going across a row. The chemical reactivity of an element depends on its ability to gain or lose electrons. › We will look at chemical reactivity through two properties. The property of metals to lose electrons. The property of nonmetals to gain electrons. In the Alkali Metal group, Cesium is much more likely to lose an electron than Lithium. › Each has an s1 electron in its valence shell. › For Li, that electron in principle energy level 2, and very close to the + charge of the nucleus. › For Cs, that electron is in principle energy level 7, and is much further away from the pull of the positive nucleus. The same trend is seen in the Alkaline Earth Metals. The most chemically active metals reside in the lower left hand corner of the table. In the Halogen group, Fluorine has a much higher electronegativity than Astatine. › At has 76 more electrons than F › The more massive the electron cloud, the bigger the force of negative repulsion a new electron must face. › In F, the nucleus does not have so many electrons to pull through, and it can attract that extra electron more easily. The same trend is seen in Oxygen’s group. The most chemically active nonmetals are found in the upper right hand corner. Ionization Energy – the energy required to remove an electron from an individual atom in the gas phase. › M(g) M+(g) + eionization energy Metals have relatively low ionization energies. › Metals increase in reactivity going down the group. Nonmetals have relatively high ionization energies. › Nonmetals would rather gain than lose electrons. Ionization Energy increases going up a group and going across a row. Atomic Size in a group › As we go down a group, we add a principle energy level of electrons. › It is easy to see why this would cause an increase in atomic radius. As we go across a row, however, atomic radius decreases, even though we are adding electrons. › This is because we are also adding protons. › As the # of protons increases, so does the nucleus’s ability to pull the electron cloud in. Atomic radius increases going down a group, and decreases going across a collumn. 1. 2. 3. It is the number and type of valence electrons that primarily determine an atom’s chemistry. The organization of the periodic table can be used to approximate the electron configuration of quantum mechanical model atoms. Certain groups have special names that should be known. • • • • • • • Alkali Metals Alkaline Earth Metals Halogens Noble Gases Lanthanides Actinides Transition Metals. 4. Elements can also be divided into Metals and Nonmetals. • • • Metals give up electrons to form positive ions, and have low ionization energies. They are found on the left side of the table. The most chemically reactive metals are in the bottom left corner of the table. Metals are usually reducing agents. Nonmetals gain electrons to form negative ions when reacting with a metal. Nonmetals have high ionization energies. They are on the right side of the table. The most reactive nonmetals are in the upper right corner (ignoring the noble gases because they are inert). Nonmetals are usually oxidizing agents. The division between metals and nonmetals is defined by the stair step line. But the division is only approximate. Many of the elements touching the division line have the properties of both metals and nonmetals. These elements are called metalloids. Where are the nonmetals found on the periodic table? Where are the metalloids found? Where are the metals found? Where are the most chemically active metals located? Where are the most chemically active nonmetals located? How does the ionization energy of an atom change periodically? How does atomic size change periodically?