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The Structure of Atoms

The Structure of Atoms. Chapter 5. Why Study This Chapter First?. A basic understanding of the atom is required to comprehend many fundamental concepts that will be discussed in this chemistry course. The behavior of atoms is a result of its fundamental particles and how they are arranged.

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The Structure of Atoms

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  1. The Structure of Atoms Chapter 5

  2. Why Study This Chapter First? • A basic understanding of the atom is required to comprehend many fundamental concepts that will be discussed in this chemistry course. • The behavior of atoms is a result of its fundamental particles and how they are arranged. • You will be able to answer • Why atoms form compounds? • Why atoms/elements have similar chemical and physical properties? • Why atoms combine in certain ratios to form compounds?

  3. The Fundamental Particles in the Atom 1 amu = 1.660  10-24 g

  4. Discovery and Properties of Electrons • Humphry Davy and Michael Faraday performed experiments in the early 1800s illustrating the existence of charged particles. • Termed as ‘electrons’ • These charged particles were suggested to be responsible for holding substances together.

  5. Discovery and Properties of Electrons • J.J. Thomson (1897) used cathode-ray tubes to study these ‘electrons’ in more detail. He was able to determine the ratio of charge to mass for the electron. • e/m = 1.75882  108 coulomb/gram Original instrument used by J.J. Thomson

  6. Discovery and Properties of the Electron • Robert Millikandetermined the mass and charge of an electron using the ‘Millikan oil-drop experiment. • Measured charges were all integrals of the same number (homework question). • Charge (e-) = 1.60218  10-19 C • The mass of the electron was determined using the charge-to-mass ratio determined by Thomson. Mass (e-) = 9.10940  10-28 grams

  7. First Attempt at the Structure of the Atom (Incorrect) • In 1886, Eugen Goldstein discovered positively charged ions by using modified cathode-ray tubes. • Termed as ‘protons’ • The scientific community proposed that these positive and negative charges were mixed together. • Plum pudding model

  8. The Structure of the Atom Revisited (More correct) • In 1909, Ernest Rutherfordillustrated that the plum pudding mode for the atom was incorrect. Rutherford used alpha particles (He2+) to establish the arrangements of protons and electrons in the atom. • Scattering of alpha particles through a thin piece of gold foil was observed on a scintillation screen. If the plum pudding model were correct, all the alpha particles would be deflected by very small angles. What did he observe?

  9. The Structure of the Atom Revisited (More correct) Important observations from the Rutherford experiment. • Most of the particles passed through the foil with no or little deflection. • Surprisingly, a few were deflected at very high angles, and a few came directly back to the source (i.e. starting point).

  10. According to the Rutherford modeleach atom consists of a very dense region of positive charge surrounded by a diffuse region of negative charge. The positive region contains most of the mass. The diameter of the positive region is 1/10,000 of the atom. The density of positive region is ~1015 g/mL The atom is primarily empty space. The Structure of the Atom

  11. More Atom Information • H.G.J. Moseley used X-rays to propose that each element differs from the preceding element by one positive charge (i.e. proton) in its nucleus. • This led to the arrangement of the periodic table (according to proton number). The number of protons in the nucleus is equal to the atomic number. • Understand the periodic table and the atomic number. • James Chadwick discovered the presence of neutrons. • All elements except hydrogen contain neutrons.

  12. The Atom in Summary • Atoms contain small, dense nuclei surrounded by clouds of electrons. • The nuclei contains protons and neutrons. • Electrons are at relatively great distances from the nuclei. • Nuclear diameters are typically 10-5 nanometers while atomic diameters are 10-1 nanometers. • Basketball and six miles • All atoms contain protons. All atoms except hydrogen contain neutrons. Look at periodic table.

  13. Mass Number and Isotopes • Mass number is the sum of the ____ and ____ in an atom’s nucleus. The number of ____ determined the atomic number. This number also determines the identity of the atom. • Isotopes are atoms of the same element with different number of neutrons. • Isotopes of he same element contain the same number of ______. • Nearly all elements have more than one isotope. • hydrogen-1, hydrogen-2 (deuterium), and hydrogen-3 (tritium) • chlorine-35 and chlorine-37

  14. Using Nuclide Symbols to Indicate Isotopes • The isotopes is represented in the format, . • A is the mass number, Z is the atomic number, and E is the element. • Write the chlorine-35 and chlorine-37 with nuclide symbols. How many protons, neutrons and electrons are in each isotope? • Write plutonium-244 with nuclide symbols. What is the number of protons, neutrons, and electrons in this isotope?

  15. Atomic Weight • Atomic weight scale is based on the carbon-12 isotope. • One amu (atomic mass unit) is exactly equal to 1/12 of a carbon-12 atom. • Atomic weight is a weighted average of all _____ occurring isotopes. This is the reason that most numbers in the periodic table are significantly less or more than a whole number. • To calculate the atomic weight in amu multiply each isotope by the natural abundance and add the terms.

  16. Atomic Weight Calculations • The two naturally occurring isotopes for chlorine are chlorine-35 and chlorine-37. Isotopic masses for the two isotopes were measured at 34.9689 and 36.9659 amu. The natural abundances are 75.77% and 24.23%, respectively. Calculate the atomic weight. • The naturally occurring isotopes for strontium are Sr-84, Sr-86, Sr-87, and Sr-88. The natural abundances of these isotopes are 0.56%, 9.86%, 7.00%, and 82.58%. What is the atomic weight? • The atomic weight of gallium is 69.72 amu. The two naturally occurring isotopes are Ga-69 and Ga-71. What is the natural abundance of each isotope? Look at webelements.com

  17. Mass Spectrometers • Mass spectrometers are instruments that measure the mass-to-charge ratio of charged particles. The extent of deflection of ions in a magnetic field depends on the mass-to-charge ratio.

  18. Mass Spectrometers • Mass spectrometers are able to separate and determine the isotopic abundances of a particular element. • Ions are detected in order of mass-to-charge ratio (m/q). The ions with the smallest m/q are detected first, and the ions with greatest m/q are detected last. The intensity or heights of the peaks in a spectrum reflect the natural abundance. • The naturally occurring isotopes of B are boron-10 and boron-11. The natural abundances of the isotopes are 20% and 80%. How would the spectrum for boron appear? (mass spectra plotter) • How would the mass spectrum for Ge appear? http://www.webelements.com/webelements/elements/text/Ge/isot.html

  19. The Nature of Light • Much of the information concerning the arrangement of electrons in the atom has been based on the light emitted (given off) or absorbed by atom. • Activate a light stick; Why are only certain color emitted? This pertains to most objects that exhibit a specific color (neon signs) • Light has both ______ properties and ______ properties.

  20. Wave-like Properties of Light • Light is electromagnetic radiation that exhibits oscillating wave-like behavior. All waves are in motion. • Wavelength () is the distance between two identical points of the wave. • Frequency () is the number of wave crests passing a point per second. • The product of the wavelength and frequency is equal to the speed of light, c. =c c=3.00  108 meters/second

  21. Wave-like Properties of Light • The visible region of the electromagnetic spectrum ranges from about 4.0  10-7 m to 7.5  10-7 m. • A very tiny slice of the spectrum. • Demo: Light through a prism. White light contains many wavelengths (all visible). • The spectrum through a prism is called a _____ _____ because it contains a continuous range of wavelengths.

  22. Wave-like Properties of Light • Some problems • What is the wavelength, in angstroms, of electromagnetic radiation having a frequency of 3.00  1020 Hz (cycles per second)? • What is the corresponding frequency of electromagnetic radiation that has a wavelength of 15 nanometers? 1 nanometer = 1.0  10-9 meters

  23. Particle-like Properties of Light • Light can also exhibit particle-like properties under certain conditions. The particles that compose light are referred to as _______. Each _____ has its own frequency, and, therefore, a unique energy. • Energy of a photon can be expressed as E = h or hc/. • h (Planck’s constant) = 6.6260755  10-34 Joulessecond Energy is _____ proportional to the frequency and _____ proportional to the wavelength.

  24. Particle-like Properties of Light • Some problems • My cordless phone operates at 900 MHz. Calculate the energy of a photon at this frequency. • Red light has a wavelength near 7500 Å. Calculate the energy of a photon of red light. Increasing the intensity of light only increases the number (or flux) of photons. For a specific color to be observed the photons emitted have to posses a specific energy. How did this relate to the atom?

  25. Quantization of the Electron in the Atom • The Bohr Atom • When electric current is passed through hydrogen gas only several emission lines are observed in a spectrum (only certain colors). • DEMO: Emission spectrum of hydrogen • For the hydrogen atom, Balmer and Rydberg illustrated that the observed wavelengths are related to the following expression: Where R has a value of 1.097  107 m-1 (Rydberg constant). The n’s are positive integers, with n2 greater than n1.

  26. Quantization of the Electron in the Atom • The Bohr atom • Niels Bohr proposed that the electron energy is quantized; meaning that only certain values of energy are allowed/possible for the electron. Energy will be absorbed or emitted in specific amounts as the electron moves from one energy level to another. • Even though the Bohr atom model has flaws, it illustrates the quantization of the electron. • Look at Fig. 5-16 in book and discuss.

  27. Quantization of the Electron in the Atom • What does this mean in summary? • An atom possesses a number of energy levels in which electrons can exist. • The lowest possible energy level is called the ground state (n=1). The higher energy levels are called excited excited states. • In order for an electron to move from a lower energy level to a higher energy level, it must ____ an amount of energy equal to the difference between the two levels.

  28. Quantization of the Electron in the Atom • What does this mean in summary? • The electron may stay located in an energy level without emitting or absorbing energy. • Movement of an electron to a lower energy level will emit energy equal to the difference between the two levels. Conservation of energy is observed in this process.

  29. Atomic Emission • Emission spectrum – certain wavelengths of light are emitted from a substance • Light can be separated by a prism or grating. This is called a bright-line spectrum. • The lines are caused by electron transitions from higher energy levels to lower levels. DEMO: Gratings on a few gases.

  30. Atomic Absorption • Absorption spectrum - certain wavelengths of light are absorbed by a substance. • These wavelengths correspond to electron being promoted to higher levels. The wavelengths are removed. • The wavelengths that are absorbed are usually given off in the emission spectrum.

  31. Electrons Can Behave Like Waves • Louis de Broglie illustrated that small particles, such as electrons, can exhibit wave-like properties. The wavelength can be expressed as =h/m h(Planck’s constant) = 6.626  10-34 Joulesec In units, Therefore, Planck’s constant can also be written as

  32. Electrons Can Behave Like Waves • Some problems • What is the wavelength of a proton moving at 2.50  107 m/s? • What would be the wavelength of a car (mass 1500 kg) moving at a velocity of 15.2 m/s? • For electrons in atoms, classical interpretations do not work. The electron must be treated as having wave-like properties. This is the basis for quantum mechanics. DEMO: Diffraction limit related to the optical microscope versus the electron microscope (calculation).

  33. Basic Ideas of Quantum Mechanics • Atoms or molecules can only exist in certain energy states. • Atoms or molecules emit or absorb energy when they change their energy state. • Equal to the difference between the two energy states. • The allowed energy states of atoms or molecules can be described by a set of quantum numbers.

  34. Quantum Mechanics and the Atom • The electron in the atom can be treated as a standing wave. Only certain “waves” are allowed/permitted for an electron around the nucleus. Each wave corresponds to an energy state. • Schrödinger developed an equation that utilizes standing waves to calculate the energies of the electron in the hydrogen atom.

  35. Quantum Mechanics and the Atom • The equation can be solved for the hydrogen atom to produce the values that are allowed for the electron in the hydrogen atom. Each solution (or set of values) can be described by a set of quantum numbers. More complex equations are required for atoms with many electrons. • Four quantum are necessary to describe the energy state of an electron in the atom. • Principal quantum number, angular momentum quantum number, magnetic quantum number, and spin quantum number

  36. Quantum Numbers (#1) • Principal quantum number, n • The shell number or main energy level. It may be any positive integer. • Primary indicator of energy • Distance from the nucleus increases with n • The further the electron from the nucleus the greater the energy. Therefore, energy increases with n.

  37. Quantum Numbers (#2) • Angular momentum quantum number, l • Corresponds to the subshell label • Designates the shape of the region of space that the electron occupies (orbital) • The number of subshells (or values of l) is equal to n, the principal quantum number • l = 0, 1, 2, … (n-1) What are the possible values of l if n is equal to 3? • s=0, p=1, d=2, f=3

  38. Quantum Numbers (#3) • Magnetic quantum number, ml • Specifies the orbital in a subshell that the electron is assigned • This number specifies the orientation of that orbital in space • The number of orbitals in a subshell increases with l • ml = (-l), …, 0, …., (+l) • Number of individual orbitals = 2l+1 What are the possible values of ml if l = 2 and 3?

  39. Quantum Numbers (#4) • Spin quantum number, ms • Refers to the spin of an electron • For every value of ml(or orbital), ms can have a value of +1/2 or –1/2 • Each orbital can accommodate only ___ electrons? What are the possible values of ms if ml equals 2? How about if ml equal 55? Look at Table 5-4 for allowed values of quantum numbers. If n=3, what are the allowed quantum numbers? These values are the possible energy states for an electron in the 3rd shell.

  40. Atomic Orbitals • An atomic orbital is a region of space in which the probability of finding an electron around the nucleus is high (usually > 75%). • The electrons must occupy an orbital. Electrons in orbitals do not orbit or circle around the nucleus. • Orbitals contain diffuse clouds of electrons that are at significant distances away from the nucleus.

  41. Atomic Orbitals • Three quantum numbers are used to specify the orbital in which the electron is located • n, l, and ml • The fourth quantum number, ms, is used to specify the spin of the electron in the orbital. • The shell number, n, of an atomic orbital has a value of 1, 2, 3, ….. • Indicates the energy of the electron in that orbital • Each shell is further away from the nucleus (higher energy) • Each shell has a capacity for 2n2 electron

  42. Atomic Orbitals and Subshells • s subshell • For the s subshell, l = 0. • Every shell has an s subshell that contains ___ orbital (ml =0). How many electrons can each subshell hold? What are the values? • These subshells are designated as 1s, 2s, 3s, etc. How many subshells for n=1? n=3? • The s orbital is spherically symmetric (illustrate) • The diameter of s orbital increases with shell number. • The energy of an electron in the 4s orbital is higher than the energy of an electron in the 2s orbital. Why?

  43. Atomic Orbitals and Subshells • s subshell • There are regions of space where the electrons are not allowed to be for s orbitals having a shell number greater than 1. These regions are called _____. • DEMO: Show with orbital software • The number of nodes increases with shell number. Which has more nodes, 3s or 1s?

  44. Atomic Orbitals and Subshells • p subshell • All shells equal to 2 or higher possess a p subshell • For any p subshell, l = 1 • The three p orbitals are mutually perpendicular and shaped like dumbbells • The subscript indicates the Cartesian scale (px, py, and pz) • ml = __ __ __ • With increasing n, the p subshells increase in energy and the number of nodes. The distance from the nucleus also increases. DEMO: Illustrate the 2p and 4p

  45. Atomic Orbitals and Subshells • d subshells • All shells equal to 3 or higher possess a d subshell • For any d subshell, l = 2 • Each subshell consists of five d orbitals • ml= __ __ __ __ __ • There are four clover-leaf-shaped orbitals and 1 orbital shaped like a peanut with a donut around the center DEMO: Illustrate orbitals with software

  46. Atomic Orbitals and Subshells • f subshell • All shells equal to 4 or higher possess and f subshell • For any f subshell, l = 3 • Each subshell consists of seven orbitals • ml = __ __ __ __ __ __ __ • Shapes are complicated DEMO: Illustrate a 4f orbital with software

  47. Atomic Orbitals and Subshells • Recall that a particular orbital is described by a unique set of n, l, and ml values. • Each orbital can hold two electrons. These two electrons are described by the ___ quantum number • Values equal __ and __ • The electrons are spin-paired (opposite spin). One electron has an up spin, , and one electron has a down spin, . Spin pairing produces lower energy.

  48. Atomic Orbitals and Subshells • Summarizing • Number of subshells in a shell is equal to __ • Number of orbitals in a shell is equal to __ • Number of electrons in a shell is equal to __ Relation of atomic orbitals is illustrated nicely in Table 5-4 (also on demo. CD, screen 7-12) Handouts Problems

  49. Electron Configurations – Electron Buildup • For atoms with more than one electron, the Schröndinger equation becomes much more complicated. There are not exact solutions. For these atoms, the electrons are approximated to occupy atomic orbitals similar to those developed for the hydrogen atom. This is called the orbital approximation. With appropriate modifications, this can be a very good approximation.

  50. Electron Configurations – Electron Buildup • How electrons are placed into individual orbitals is best understood by examining the electron configurations as a function of increasing atomic number. • Aufbau Principle – electrons will be added to orbitals in such a way to acquire the lowest energy for the atom. • Orbitals increase in energy with increasing, n, or shell number. • Orbitals increase in energy with increasing value of l or subshell. • In the 2nd shell, the 2s would fill before the 2p

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