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Understanding the Atom

Understanding the Atom. Learning Objectives. TSWBAT = The student will be able to: Explain the different models of the atom Explain how electronic transitions produce atomic spectra (light) Explain the wave-particle duality theory

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Understanding the Atom

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  1. Understanding the Atom

  2. Learning Objectives TSWBAT = The student will be able to: Explain the different models of the atom Explain how electronic transitions produce atomic spectra (light) Explain the wave-particle duality theory Explain the evidence that suggest that light exists as a wave Explain the composition of waves Calculate frequency & wavelength Explain the Michelson Experiment Describe how Rutherford’s experiment provides evidence that light consisted of particles

  3. ACOS Correlation ACOS #3c: • Utilizing benchmark discoveries to describe the historical development of atomic structure, including photoelectric effect, absorption, and emission spectra of elements

  4. Theory #1 p101 (Democritus) • He basically postulated that everything in life • was composed of atoms. • Democritus didn’t know much about • chemical behavior. • He defined the atom to be small indestructible • particles of matter. • No experimental evidence, simply theory

  5. Theory #2: John Dalton • Expanded on Democritus Believes that • everything is made of atoms • Atoms of the same element have the same • mass/different elements have different masses • Compounds contain atoms of more than one • element • In a particular compound, atoms of different • elements always combine in the same way ex H2O • Was created to explain why compounds always join in the • same fashion or had a fixed composition.

  6. Sizing Up The Atom • The diameter of an atom is approximately 1 × 10-8 cm • To help explain this concept, let’s use a • copper penny. • Now, let’s say I grind the penny to find • particles. • Let’s say, I pull a piece of the dust and grind it up to • the tiniest piece. This tiniest particle is known as an • atom.

  7. How do we visualize the atom? • By using a scanning electron microscope!

  8. Theory #3: Joseph John Thomson • Basically identified that the atom contained subatomic particles (charges). • Proved his hypothesis using the machine on p103 Figure B. • Basically he sent an electric current through a gas tube. • Once you sent an electric current through the gas tube, energy • was released in the form of light ex: fluorescent light bulbs. • He then placed negative/positive magnets outside the tubing • and noticed that the beam deflected • He discovered the electrons • Often called the chocolate chip ice cream model

  9. Robert Millikan • Millikan supported Thomson’s theories by defining the charge • and mass of the Electron. • The charge of the electron is negative • The mass of the electron is approximately 9.11 x 10-31 kg. • Because the mass is so small, it isn’t used in determining the • mass of the atom.

  10. Eugen Goldstein • Goldstein supported Thomson’s theories by identifying the presence • of a protons • The charge of the proton is positive • The mass of the proton is approximately (1840)x (9.11 x 10-31 kg)

  11. James Chadwick • Chadwick supported Thomson’s theories by identifying the • presence of a neutron. • The charge of the neutron is neutral • The mass of the neutron is approximately (1840)x (9.11 x 10-31 kg) • Notice: The mass of the neutron is identical to that of the proton.

  12. Theory #4 p107- p108 (Ernest Rutherford’s Theory) • Used alpha particles in uranium • An alpha particle is a fast moving particle that has a • positive charge • Contradicts Thompson’s model because we identified • the positively charged nucleus • Used the Gold Foil experiment that enabled him to • confirm his beliefs about the nucleus.

  13. Theory #4 p107- p108 (Ernest Rutherford’s Theory) • The basic premise of the Gold Foil experiment was that he shot a • beam of uranium light towards a gold piece of aluminum foil. • He observed that some particles were deflected off. • After further research he observed that the charges being deflected • were positively charged and they were only deflected when they came • in proximity with the positively charged nucleus. • Believed that the electrons remained in a fixed orbital in a • three-dimensional rotation.

  14. Theory #5 p128-129 (Bohr’s Theory) • Believed that the e- move with constant speed in an more orderly • system around the nucleus. • At atom’s energy can change if we gain or lose an electron • The possible energy that an electron in an atom can have is called • and energy level. • Energy levels increase as you go away from the nucleus. • This theory is the most important to learn in our class because it • helps us to understand the details of the periodic table.

  15. Notes We learned in the past section that electrons can gain or lose energy, ultimately causing them to change energy levels. We will extend upon this knowledge to explore how light is produced.

  16. Atomic Spectra • Atomic spectra – • atomic display of energy emission (see below) • In layman’s terms, it means it’s a visual image that occurs as a result of energy being released into the system. • As energy is released, it produces the colors below • Why is this important? • No two elements have the same emission spectra (patterns) therefore it can be used as an identifier for atomic spectra • Can definitely be used in forensic science

  17. Notes: • In order to see color, light must be present. • Evidence: Turn off your lights at home, what colors are you able to see? • Let’s now explore how light is produced. • Key Terms: • Quantum numbers (n)– these numbers are used to correspond to the energy levels • Photons – packets of energy that carry light

  18. How is light produced? Emission of light occurs as a single abrupt step called an electronictransition going from a high energy level to a lower energy level (Bohr’s Model relates to this) Light is produced because electrons absorb energy. When energy is absorbed, e- are bumped into a higher energy level. Once in the higher energy level, e- release heat and light as they return to the lowest energy level.

  19. Wave – Particle Duality Theories suggests that light behaves as both a wave and as a particle

  20. Wave Behavior • What evidence supports the idea that light behaves as a wave? • It is believed that light can behave as a wave because of it’s interference patterns

  21. Evidence that light has wave-like behavior • Thomas Young did an experiment using light and comparing the pattern formed by manipulating light. • Constructive Interference – light is amplified or is more intense because of the manipulation • Destructive Interference – light is reduced based on certain non-helpful manipulations ie blocking the light patterns.

  22. Wave Composition Amplitude – height of the wave Wavelength – the distance between two identical points on a wave Crest– the peak Trough – the valley

  23. Frequency • Frequency is defined as the number of cycles (revolutions) divided by time. • Written as v (pronounced mu) • SI Unit = the hertz (s-1 or 1/s)

  24. Frequency • Frequency and time are inversely proportionate • Frequency and energy are directly proportionate • Frequency and wavelength are inversely proportionate • Large frequency/high energy can be very dangerous to • human health i.e. tanning salons.

  25. Period The time it takes for one complete cycle to occur

  26. Why does this matter? Color is a reflection of a wave frequency Light consists of electromagnetic waves Visible wavelengths are 10-2 through 10-11 (m) approximately Knowing this information can help us design devices

  27. Calculating wavelength and frequency • c= speed of light & λ =wavelength • Speed of light is approx. equal to 3 x 108 m/s • Solving for frequency (v= c/λ): • A microwave (from an antenna) is part of the electro-magnetic spectrum therefore it’s speed is the speed of light.  What is it’s wavelength if its frequency is 3.44 x 10 9 hertz? • Solving for wavelength (λ = c/v) • Calculate the wavelength of an FM radio with frequency waves of 102.5 x 102 m.

  28. How did we calculate the speed of light? • Michelson Experiment • It helped us to place a value on the speed of light. He used mirrors to time how fast the light would travel based on known data for the mirrors. He basically used the formula for speed.

  29. Evidence that Light behaves as particles Let’s revisit Rutherford’s experiment He shot a beam of Uranium (U) light into a piece of (Gold)Au and noticed particle deflection. If light didn’t exist in particles, we wouldn’t have observed this deflection

  30. The Photoelectric Effect The photoelectric effect is the name given to the observation that when light is shone onto a piece of metal, a small current flows through the metal. The light is giving its energy to electrons in the atoms of the metal and allowing them to move around, producing the current.

  31. So where do we stand now?

  32. Heisenberg Uncertainty Principle • States that it is impossible to know exactly both the velocity and the position of a particle at the same time. • When the e- collides with a photon of light, it changes the velocity/direction of the e- in ways that scientists are unsure of . • Based on what we learned about motion in Physical Science, we must know the speed and direction of an object in order to clearly describe it’s motion.

  33. Louis de Broglie & Schrodinger’s • Both individuals did tremendous work on explaining how electrons move in this wavy like pattern thus explaining why light behaves as both as wave and a particle. (p144-145)

  34. Why does any of this matter? • Understanding properties of light have helped scientist to enhance technology. • For example: • light telescopes manipulate light to help us view the stars • electron microscopes manipulate light, mirrors and electrons to view images • lasers manipulate light during surgery to cut things

  35. Electron Cloud Model of the Atom • Contradicts Bohr’s model because we know that the electrons move • around in a less predictable way. • The density of the cloud is greater in the region where the cloud is • the thickest • The thicker the region, the higher probability of finding an electron • An orbital is a region of space where an electron is most likely to • be found • Ground state refers to the lowest energy level. • There are four known orbitals: s,p,d,f

  36. Counting Atoms 2H20 3CO2 Al2O3 C2H6 FeSO4 Pb(SO4)3 Na2SO4 3K3Fe(SCN)6

  37. Most elements can be found on earth (with the exception of • those elements that too unstable and thus must be • synthesized in the laboratory). • Since all elements have isotopes then we must consider how • much of one isotope of an element exists versus another • isotope of the same element. • These are called the "natural" abundances on earth.

  38. Next, we can inquire what the mass of element X is? • Since each isotope has a different mass (because each isotope has • a different number of neutrons) the simplest answer is to give the • "average" mass of element X - the atomic weight. • After more analysis the mass of each isotope is determined to be • the following:

  39. Then the average mass (atomic weight) is given by:

  40. Rubidium has two common isotopes, 85Rb and 87Rb. If the abundance of85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average • atomic mass of rubidium?

  41. 2) Uranium has three common isotopes. If the abundance of 234U is 0.01%, the abundance of 235U is 0.71%, and the abundance of 238U is 99.28%. What is the average atomic mass of uranium?

  42. 3) Titanium has five common isotopes: 46Ti (8.0%), 47Ti (7.8%), 48Ti (73.4%), 49Ti (5.5%), 50Ti (5.3%). What is the average atomic mass of titanium?

  43. Here are three isotopes of an element: 612C 613C 614C • The element is: __________________ • The number 6 refers to the _________________________ • The numbers 12, 13, and 14 refer to the ________________________ • How many protons and neutrons are in the first isotope? ________________ • How many protons and neutrons are in the second isotope? _________________ • How many protons and neutrons are in the third isotope? _______________

  44. Complete the following chart:

  45. Naturally occurring europium (Eu) consists of two isotopes was a mass of 151 and 153. Europium-151 has an abundance of 48.03% and Europium-153 has an abundance of 51.97%. What is the atomic mass of europium?

  46. Strontium consists of four isotopes with masses of 84 (abundance 0.50%), 86 (abundance of 9.9%), 87 (abundance of 7.0%), and 88 (abundance of 82.6%). Calculate the atomic mass of strontium.

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