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Chapter 2

Chapter 2. The Structure of the Atom and the Periodic Table. 2.1 Composition of the Atom. Atom - the basic structural unit of an element The smallest unit of an element that retains the chemical properties of that element. Electrons, Protons and Neutrons.

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Chapter 2

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  1. Chapter 2 The Structure of the Atom and the Periodic Table

  2. 2.1 Composition of the Atom • Atom - the basic structural unit of an element • The smallest unit of an element that retains the chemical properties of that element

  3. Electrons, Protons and Neutrons • Atoms consist of three primary particles • electrons • protons • neutrons • Nucleus - small, dense, positively charged region in the center of the atom • protons - positively charged particles • neutrons - uncharged particles

  4. Characteristics of Atomic Particles • Electrons are negatively charged particles located outside of the nucleus of an atom • Protons and electrons have charges that are equal in magnitude but opposite in sign • A neutral atom that has no electrical charge has the same number of protons and electrons • Electrons move very rapidly in a relatively large volume of space while the nucleus is small a dense

  5. Symbolic Representation of an Element Charge of particle Mass Number Symbol of the atom Atomic Number • Atomic number (Z) - the number of protons in the atom • Mass number (A) - sum of the number of protons and neutrons

  6. Atomic Calculations number of protons + number of neutrons = mass number number of neutrons = mass number - number of protons In a neutral atom, the number of protons = number of electrons Very Important: the number of protons does not change!

  7. Selected Properties of the Three Basic Subatomic Particles Name Charge Mass(amu) Mass (grams) Electrons (e) -1 5.4 x 10-4 9.1095 x 10-28 Protons (p) +1 1.00 1.6725 X 10-24 Neutrons (n) 0 1.00 1.6750 x 10-24

  8. Determining the Composition of an Atom Calculate the number of protons, neutrons and electrons in each of the following:

  9. Isotopes • Isotopes - atoms of the same element having different masses • contain same number of protons • contain different numbers of neutrons 4 Isotopes of Hydrogen Hydrogen (Hydrogen - 1) Tritium (Hydrogen - 3) Deuterium (Hydrogen - 2)

  10. Isotopic Calculations • Isotopes of the same element have identicalchemical properties • Some isotopes are radioactive • Find chlorine on the periodic table • What is the atomic number of chlorine? • What is the mass given? • This is not the mass number of an isotope

  11. Atomic Mass • Find chlorine (Cl) on the table. What is this number, 35.34? • The atomic mass - the weighted average of the masses of all the isotopes that make up chlorine • Chlorine consists of chlorine-35 and chlorine-37 in a 3:1 ratio • Weighted average is an average corrected by the relative amounts of each isotope present in nature

  12. Determining Atomic Mass Calculate the atomic mass of naturally occurring chlorine if 75.77% of chlorine atoms are chlorine-35 and 24.23% of chlorine atoms are chlorine-37 Step 1: convert the percentage to a decimal fraction 0.7577 chlorine-35 0.2423 chlorine-37

  13. Step 2: Multiply the decimal fraction by the mass of that isotope to obtain the isotope contribution to the atomic mass. For chlorine-35: 0.7577 x 35.00 amu =26.52 amu For chlorine-37 0.2423 x 37.00 amu =8.965 amu Step 3: sum these partial weights to get the weighted average atomic mass of chlorine: 26.52 amu + 8.965 amu =35.49 amu

  14. It’s just like calculating your grade! • 20% of the quiz average • 60% of the exam average • 20% of the final exam Score = 0.2*quiz ave. + 0.6*exam ave. + 0.2*final

  15. Atomic Mass Determination • Nitrogen consists of two naturally occurring isotopes 99.63% nitrogen-14 with a mass of 14.003 amu 0.37% nitrogen-15 with a mass of 15.000 amu • What is the atomic mass of nitrogen?

  16. Ions • Ions - electrically charged particles that result from a gain or loss of one or more electrons by the parent atom • Cation - positively charged result from the loss of electrons 23Na  23Na+ + 1e- • Anion - negatively charged results from the gain of electrons 19F + 1 e- 19F-

  17. 2.2 Development of Atomic Theory • Dalton’s Atomic Theory - the first experimentally based theory of atomic structure of the atom.

  18. Postulates of Dalton’s Atomic Theory • All matter consists of tiny particles called atoms • An atom cannot be created, divided, destroyed, or converted to any other type of atom • Atoms of a particular element have identical properties

  19. Atoms of different elements have different properties • Atoms of different elements combine in simple whole-number ratios to produce compounds (stable aggregates of atoms) • Chemical change involves joining, separating, or rearranging atoms • Postulates 1, 4, 5 and 6 are still regarded as true.

  20. Evidence for Subatomic Particles: Electrons, Protons and Neutrons • Electrons were the first subatomic particles to be discovered using the cathode ray tube Indicated that the particles were negatively charged.

  21. Evidence for Protons and Neutrons • Protons were the next particle to be discovered, by Rutherford • Protons have the same size charge but opposite in sign • Proton is 1837 times as heavy as electron • Neutrons • Postulated to exist in 1920’s but not demonstrated to exist until 1932 by Chadwick. • Almost the same mass as the proton (slightly heavier).

  22. Thomson’s model of the atom http://nobelprize.org/educational_games/physics /quantised_world/structure-images/fig2b.gif The “plum pudding” model.

  23. Evidence for the Nucleus • Earnest Rutherford’s “Gold Foil Experiment” lead to the understanding of the nucleus • Most alpha particles pass through the foil without being deflected • Some particles were deflected, a few even directly back to the source

  24. Rutherford’s Gold Foil Experiment • Most of the atom is empty space • The majority of the mass is located in a small, dense region

  25. Rutherford’s model of the atom http://www2.kutl.kyushu-u.ac.jp/seminar/ MicroWorld1_E/Part2_E/P25_E/atom.gif

  26. 2.3 Light, Atomic Structure, and the Bohr Atom • Rutherford’s atom – tiny, dense, positively charged nucleus of protons surrounded by electrons • How do we describe the relationship of the electrons to each other and the nucleus? • The problem... our classical understanding of physics didn’t work for the atom! This will take some explaining.

  27. Light and Atomic Structure • Spectroscopy - absorption or emission of light by atoms. • Used to understand the electronic structure. • To understand the electronic structure, we must first understand light, Electromagnetic Radiation • travels in waves from a source • speed of 3.0 x 108 m/s

  28. Radio • Knew from radio that if we accelerate an electron back and forth in a wire it will radiate a radio wave (electromagnetic radiation).

  29. Wavelengths • Light is propagated (moves) as a collection of sine waves • Wavelength is the distance between identical points on successive waves • Each wavelength travels at the same velocity, but has its own characteristic energy

  30. Electromagnetic Spectrum low energy long wavelength high energy short wavelength

  31. Bohr Theory • Atoms can absorb and emit energy via promotion of electrons to higher energy levels and relaxation to lower levels • Energy that is emitted upon relaxation is observed as a single wavelength of light • Spectral lines are a result of electron transitions between allowed levels in the atoms (in other words, allowed energies are “quantized” in that only certain quantities of energy are allowed.

  32. The Bohr Atom Electrons exist in fixed energy levels surrounding the nucleus Quantizationof energy Energy is released as the electron travels back to lower levels Relaxation Promotion of electron occurs as it absorbs energy Excited State

  33. Electronic Transitions • Amount of energy absorbed in jumping from one energy level to a higher energy level is a precise quantity • Energy of that jump is the energy difference between the orbits involved • Orbit- what Bohr called the fixed energy levels • Ground state - the lowest possible energy state

  34. Modern Atomic Theory • Bohr’s model of the atom when applied to atoms with more than one electron failed to explain their line spectra • One major change from Bohr’s model is that electrons do not move in orbits • Atomic orbitals - regions in space with a high probability of finding an electron • Electrons move rapidly within the orbital giving a high electron density

  35. The Quantum Mechanical Atom • Bohr’s model of the hydrogen atom didn’t clearly explain the electron structure of other atoms • Electrons in very specific locations, principal energy levels • Wave properties of electrons conflict with specific location • Schröedinger developed equations that took into account the particle nature and the wave nature of the electrons

  36. Schröedinger’s equations • Equations that determine the probability of finding an electron in specific region in space, quantum mechanics • Principal energy levels (n = 1,2,3…). n is also known as the principal quantum number. But there is more to the structure of the arrangement of electrons found around the nucleus of an atom, as we will see.

  37. Energy Levels and Sublevels PRINCIPAL ENERGY LEVELS • n = 1, 2, 3, … • The larger the value of n, the higher the energy level and the farther away from the nucleus the electrons are • The number of sublevels in a principal energy level is equal to n • in n=1, there is one sublevel • in n = 2, there are two sublevels

  38. The angular momentum quantum number. • An object, such as an electron, that moves around another object (the nucleus in our case) will have angular momentum. • Given by l=0,1,...n-1 • l=0 is s subshell, l=1 is p subshell, etc. • Subshells increase in energy: s<p<d<f • e.g., electrons in 3d subshell have more energy than electrons in the 3p subshell

  39. Sublevels in Each Energy Level

  40. Orbitals Orbital - a specific region of a sublevel containing a maximum of two electrons • The number of orbitals in a subshell is given by the magnetic quantum number. • m = -l,...,0,...,+l • Orbitals are named by their sublevel and principal energy level • 1s, 2s, 3s, 2p, etc. • Each type of orbital has a characteristic shape • s is spherically symmetrical • p has a shape much like a dumbbell

  41. s is spherically symmetrical Each p has a shape much like a dumbbell, differing in the direction extending into space Orbital Shapes

  42. m = -l, ...0,...+l • An orbital can only hold two electrons!! Once it is filled it cannot accept more.

  43. Electron Configuration • Electron Configuration - the arrangement of electrons in atomic orbitals • Aufbau Principle - or building upprinciple helps determine the electron configuration • Electrons fill the lowest-energy orbital that is available first • Remember s<p<d<f in energy • When the orbital contains two electrons, the electrons are said to be paired and the orbital is full

  44. Rules for Writing Electron Configurations • Obtain the total number of electrons in the atom from the atomic number • Electrons in atoms occupy the lowest energy orbitals that are available – 1s first • No more than 2 electrons in any orbital • Maximum number of electrons in any principal energy level is 2(n)2 • Follow the periodic table!

  45. H Hydrogen has only 1 electron It is in the lowest energy level & lowest orbital Indicate number of electrons with a superscript 1s1 Li Lithium has 3 electrons First two have configuration of Helium – 1s2 3rd is in the orbital of lowest energy in n=2 1s2 2s1 Writing Electron Configurations

  46. Classification of Elements According to the Type of Subshells Being Filled

  47. Electron Configuration Examples • Give the complete electron configuration of each element Be N Na Cl Ag

  48. Shorthand Electron Configurations • Uses noble gas symbols to represent the inner shell and the outer shell or valance shell is written after • Aluminum- full electron configuration is: 1s22s22p63s23p1 What noble gas configuration is this? • Neon • Configuration is written: [Ne]3s23p1

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