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Chapter 10 Modern Atomic Theory Models of the Atom • Rutherford’s model of an atom stated that an atom is mostly empty space with single dense positive core with negatively charged electrons moving around the nucleus in some way – Rutherford's model of an atom could not explain why metals or compounds of metals give off characteristic colors when heated in a flame, or why objects, when heated to higher and higher temperatures, first glow dull red, then yellow, then white – Rutherford’s atomic model could not explain the chemical properties of elements Electromagnetic Radiation • Our current model of atoms and how electrons are arranged around the atom came from the study of light Electromagnetic Radiation • Light is a form of electromagnetic radiation • Electromagnetic radiation – Radiant energy that exhibits wavelike behavior and travels through space at the speed of light in a vacuum Wave Length (λ) Crest Amplitude Trough 𝐹𝑟𝑒𝑞𝑢𝑒𝑛𝑐𝑦 ν = 𝐶𝑦𝑐𝑙𝑒𝑠 𝑆𝑒𝑐𝑜𝑛𝑑 1 Second 4 Cycles Electromagnetic Radiation – Electromagnetic radiation consists of waves • Waves cycle – Each complete wave cycle starts at zero, increases to highest value (the crest), passes through zero and down to its trough and back up to zero – The amplitude of the wave is its height at the crest – The wavelength, represented by λ (lambda), is the distance between the crests – The frequency, represented by ν (nu), is the number of waves that to pass a given point per unit of time » The units of frequency are usually cycles per second called a hertz (Hz) or S-1 – The product of frequency and wavelength always equals a constant (C), the speed of light 3.00x108 m/s c = λν – The wave length and frequency of light are inversely proportional to each other Electromagnetic Radiation • Light is a form of electromagnetic radiation that includes radio waves, microwaves, infrared waves, visible light, ultraviolet light, x-rays, and gamma rays – All electromagnetic rays travel in a vacuum at a speed of 3.00 x 108 m/s • Sunlight consists of light with a continuous range of wavelengths and frequencies • When the different frequencies of light pass through a prism they are separated into a spectrum Electromagnetic Radiation http://upload.wikimedia.org/wikipedia/commons/c/cf/EM_Spectrum_Properties_edit.svg Atomic Spectrum of Hydrogen • When H2 molecules absorb energy, some H-H bonds are broken and the resulting hydrogen atoms are excited – When excited, the hydrogen atoms contain excess energy that is released in the form of light at specific wavelength producing an special emission spectrum called a line spectrum – The spectrum indicated that only certain energies are allowed for the electron in the hydrogen atom • In other words, the energies were quantized Atomic Spectra • When atoms absorb energy, electrons move into higher energy levels, and these electrons lose energy by emitting light when they return to lower energy levels • Light emitted by an element separates into discrete lines to give an atomic emission spectrum of the element • Each discrete line in an emission spectra corresponds to one exact frequency of light emitted by the atom • The emission spectrum of each element is like a person’s finger print and can be used to identify each element Atomic Spectra and the Bohr Model • Niels Bohr proposed that an electron is found only in specific circular paths, or orbitals, around the nucleus – Each possible electron orbit in Bohr’s model has a fixed energy. • The fixed energies an electron can have are called energy levels • To move from one energy level to another, an electron must gain or lose just the right amount of energy called a quantum – Thus, the energy of electron is said to be quantized – The energy levels are not equally spread, the higher energy levels are closer together The Bohr Model • The electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits n=5 n=4 n=3 n=2 n=1 The orbits are known as principle energy levels aka shells Line Spectrum Wavelength When not excited, the electron in the hydrogen atom resides in the n = 1 energy level. A certain quantum of energy is required for an electron to move to a higher shell (energy level) An explanation of Atomic Spectra • The lowest possible energy of the electron is its ground state • When an electron absorbs energy and is excited, it jumps in quanta’s to higher energy levels • A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level • The emission occurs in a single abrupt step • The quanta of energy released is related to the frequency of the emitted light by the equation E = hν where h = 6.626x10-34 J-s • The light emitted by an electron moving from a higher to a lower energy level has frequency directly proportional to the energy change of the electron • Light has both wave and particle properties Light Frequency, Wavelength, and Energy Relationship 𝐶 = λν 𝐸 = ℎν • Increasing wave length result in decreasing frequency which result in light with less energy • Increasing frequency results in decreasing wavelength which result in light with more energy The Quantum Mechanical Model Schrodinger • Like the Bohr model, electrons are restricted to certain energy levels • Unlike the Bohr model, the quantum mechanical model does not involve an exact path the electron takes around the nucleus – Electron paths are not circular • The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus Quantum Mechanics • Photons have wave an particle properties • Everything travels in wave like motion but the mass of the particle must be small enough to observe the wave like nature • Heisenberg uncertainty principle states that is impossible to know exactly both the velocity and the position of a particle at the same time Quantum Mechanical Model • Heisenberg uncertainty principle – There is a fundamental limitation to how precisely both the position and momentum of a particle can be known at a given time Atomic Orbitals • Atomic orbitals are often thought of as a region of space in which there is a high probability of finding an electron – Each orbital is characterized by a series of numbers called quantum numbers, which describe various properties of the orbital: • Energy levels of electrons in the quantum mechanical model are called principle quantum numbers (n) and are assigned the numbers 1,2,3,4, and so forth Atomic Orbitals • An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron • For each principle energy level, there may be several orbitals with different shapes and different energy levels that constitute energy sublevels (subshells) • Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found • Different atomic orbitals are denoted by letters: s, p, d, and f – – – – • The number and kinds of atomic orbitals depend on the energy sublevel – – – – • There is only one type of s orbital There are three types of p orbitals There are five types of d orbitals There are seven types of f orbitals The lowest energy level n = 1 can only hold one orbital s The second energy level n = 2 can hold four orbitals s, p1, p2, p3 The third energy level n = 3 can hold 9 orbitals s, p1, p2, p3, d1, d2, d3, d4, d5 The fourth energy level n = 4 can hold 16 orbitals ... Two electrons with opposite spin can potentially occupy each orbital – The maximum number of electrons that can occupy a principle energy level is given by the formula 2n2, where n is the principle quantum number Atomic Orbitals Electron Arrangements in Atoms • In atoms, electrons and the nucleus interact to make the most stable arrangement possible • Three rules you need to know about electron arrangements are: the Afbau principle, the Pauli exclusion principle, and Hund’s rule that tell how to determine the electron arrangement of atoms Electron Arrangements in Atoms • Aufbau principle – electrons occupy the orbitals of lowest energy first – The range of some energy levels within a principle energy level can overlap the energy levels of another principle level • Pauli exclusion principle – an atomic orbital may describe at most two electrons – To occupy the same orbital, the two electrons must have opposite spin represented with an up or down arrow ↑↓ • 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 – One electron enters each orbital until all the orbitals contain one electron with the same spin direction • Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations • Find the electron configuration of Zinc 1s2 1. Find Zinc’s atomic number • Zinc has an atomic number of 30. 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 6f14 2. Follow the arrows adding the superscripts until you reach 30 Answer 1s2 2s2 2p6 3s2 3p6 4s2 3d10 Notice 2 + 2 + 6 + 2 + 6 + 2 + 10 = 30 3. Rearrange electron configuration in order of increasing energy level Answer 1s2 2s2 2p6 3s2 3p6 3d10 4s2 • Find the electron configuration of Ruthenium (Ru) 1s2 1. Find Ruthenium’s atomic number • 2s2 2p6 3s2 3p6 Ruthenium has an atomic number of 44. 2. Follow the arrows adding the super scripts until you reach 44 3d10 Answer 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d6 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 6f14 Notice 2 + 2 + 6 + 2 + 6 + 2 + 10 + 6 + 2 + 6 = 44 Notice Only 6 electrons were needed for the 4d sublevel 3. Rearrange electron configuration in order of increasing energy level Answer 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d6 5s2 • Writing Orbital Diagrams 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 6f14 1s2 2s2 2p6 1. Write orbital boxes for each sublevel that contains electrons in the same order as you would an initial electron configuration 2. Add one electron (represented by an up arrow) to each box (orbital) in a sublevel until each orbital in that sublevel contains one electron 3. Once all orbitals in the sublevel contain one electron start paring them up with a down arrow until all boxes in the sublevel are filled 4. Once one sublevel is filled, move to the next sublevel repeating steps 3 and 4 until you have as many arrows total as you do electrons in the atom 5. The orbital diagram for 3s2 3p6 4s2 26𝐹𝑒 is shown below 3d10 • Writing Orbital Diagrams 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 6f14 1s2 2s2 2p6 1. Write orbital boxes for each sublevel that contains electrons in the same order as you would an initial electron configuration 2. Add one electron (represented by an up arrow) to each box (orbital) in a sublevel until each orbital in that sublevel contains one electron 3. Once all orbitals in the sublevel contain one electron start paring them up with a down arrow until all boxes in the sublevel are filled 4. Once one sublevel is filled, move to the next sublevel repeating steps 3 and 4 until you have as many arrows total as you do electrons in the atom 5. The orbital diagram for 3s2 3p6 4s2 26𝐹𝑒 is shown below 3d10 Aufbau Principle and the Periodic Table Valence Electrons • Valence electrons are the electrons in the outermost principal quantum level of an atom – The inner electrons are known as core electrons • Valence electrons are the most important in determining the chemical properties of an element • The elements in the same group have the same valence electron configuration Periodic Trends • Atomic Size • Ionization Energy • Electronegativity Periodic Trends: Atomic Size Periodic Trends – The atomic radius is one half of the distance between the nuclei of two atoms of same element when atoms are joined – In general, atomic size increases from top to bottom within a group and decreases from left to right across a period – Why • Periodic Trends – The increases of positive charge in the nucleus of the atom draws in it’s electrons in its highest energy level as you move left to right across a period and thus, the atomic size decreases • Group Trends – As you move down a group, new principle energy levels are filled with electrons. These filled inner energy levels shield the electrons in the outer energy levels from the positive charge of the nucleus and thus atoms are larger down a group in the periodic table Periodic Trends • Ions – Cations are smaller than the atoms they were formed from – Anions are larger than the atoms they were formed from Periodic Trends • Trends in Ionization Energy – Electrons can move to higher energy levels when atoms absorb energy – If there is enough energy, electrons overcome the attraction of protons in the nucleus and escape – The energy required to remove an electron from an atom is called ionization energy • Measured with when an element is in its gaseous state • The energy required to remove the first electron from an atom is called the first ionization energy and produces a cation with a 1+ charge • The first ionization energy tends to decrease from top to bottom within a group and increases from left to right across a period – Atoms with low ionization energy levels tend to lose electrons easily in chemical reaction becoming cations – Atoms with high ionization energy tend to gain electrons in chemical reaction becoming anions Periodic Trends: Ionization Energy Periodic Trends • Trends in Ionization Energy – Periodic Trends in ionization energy • Increases from left to right due to increasing nuclear charge, and constant shielding of the outer most electrons – Results in increased attraction of the outer electrons to the nucleus and greater energy required to remove outermost electron – Group Trends in ionization energy • Decreases from top to bottom in a group due to increased shielding of outer electron by completely filled inner energy levels and the outer electrons further distance from the nucleus – Results in decreased attraction of the outer electrons to the nucleus and less energy required to remove outermost electron Periodic Trends: Electronegativity Trends in Electronegativity • Electronegativity – When atoms react to form a compound, two kinds of bonds can result between the atoms called ionic bonds and covalent bonds – Whether an ionic or covalent bond forms can be predicted by the electronegativity of the atoms involved in making the compound • Electronegativity is the ability of an atom of an element to attract electrons when the atom is in a compound • Scientist use factors such as ionization energy to calculate values of electronegativity – In general, electronegativity values decrease from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period • Metals at the far left of the periodic table have low electronegativity values • Nonmetals at the far right of periodic table (excluding noble gases) have high electronegativity values • Electronegativity among transition metals are not as regular Bond Type and Polarity • The difference in electronegativity between the two atoms involved in the bond will determine the most probable type of bond that will form – If the difference is: • • • • 0.0—0.4 the bond is nonpolar covalent 0.4—1.0 the bond is moderately polar covalent 1.0—2.0 the bond is very polar covalent >2.0 the bond is ionic