The_Atom class 9 Neutron Proton and Electrons
- 1. Models of the Atom • The impact of modern physics is most evident in the development of the atomic model of matter • We use term atomic model to indicate that we are trying to describe the key features of the atom • We do not know what an atom “looks like”, because we have no instruments for its direct observation.
- 2. Model of the Atom • Ancient Greeks • Matter is made up of particles, but not the elements as we know them today • Four elemental substances •Air, Fire, Earth, and Wind
- 3. A Change of Thought • Robert Boyle (1600’s) • Identified gold and silver as being elemental • In other words, they weren’t made up of air, earth, fire, or wind
- 4. Dalton’s Theory • In 1700’s he theorized that the basic unit of matter is a tiny particle called an atom.
- 5. Dalton’s Theory • All elements are composed of indivisible atoms • All atoms of a given element are identical • Atoms of different elements are different • Compounds are formed by the combination of atoms of different elements
- 6. But Wait…. • Experimental studies of the atom soon showed that it (the atom) was not indivisible… it has smaller parts!
- 7. Thomson’s Model • Just over 100 years ago, J. J. Thomson discovered that electrons are relatively low mass, negatively charged particles present in atoms
- 8. • Because he knew that atoms were electrically neutral, he concluded that part of the atom must posses positive charge equal to the total charge of the electrons
- 9. • He proposed a model in which the atom consists of a uniform distribution of positive charge, in which electrons are embedded (like raisins in plum pudding). • “The Plum Pudding” Model
- 11. Rutherford’s Model • Observations • Most of the alpha particles pass straight through the gold foil. • Some of the alpha particles get deflected by very small amounts. • A very few get deflected greatly. • Even fewer get bounced off the foil and back to the left.
- 12. Rutherford’s Model • Conclusions • The atom is 99.99% empty space. • The nucleus contains a positive charge and most of the mass of the atom. • The nucleus is approximately 100,000 times smaller than the atom.
- 14. The Bohr Model • While the Rutherford model focused on describing the nucleus, Niels Bohr turned his attention to describing the electron.
- 15. The Bohr Model • Neils Bohr proposed a model showing a dense nucleus with electrons in surrounding orbitals
- 17. The Bohr Model • For electrons to stay in orbit, they must have just the right amount of energy to keep it in place around the nucleus.
- 18. The Bohr Model • The maximum number of electrons in the first energy level is two. • The second level has a maximum of eight electrons.
- 19. The Wave Mechanical Model • The major difference between the wave-mechanical model and the Bohr model is found in the manner in which the electrons are pictured.
- 20. The Cloud Model
- 21. The Wave Mechanical Model • An orbital is described as a region in which an electron is most likely to be found.
- 22. The Structure of The Atom • All atoms are composed of a dense, positively charged nucleus, surrounded by a large space occupied by electrons.
- 23. The Nucleus • The nucleus contains two types of particles • Protons - with a positive charge • Neutrons - with no charge
- 24. Subatomic Particles • Protons have a mass of 1.67x10-24 g • Because the mass is so small, we sometimes use atomic mass units or amu • A proton is assigned 1 amu. • A neutron is approx. the same
- 25. Subatomic Particles • Each atom of a specific element must contain the same number of protons as each other atom of that element.
- 26. Subatomic Particles • The number of protons in the nucleus of an atom is the atomic number of that element.
- 27. Subatomic Particles • Electrons • are much less massive than either the proton or the neutron • Have a charge equal to, but opposite, a proton • Occupy space outside the nucleus • Tape Demo
- 28. Subatomic Particles • The sum of the numbers of protons and neutrons in the nucleus is called the mass number.
- 29. Sample Problem • Find the number of neutrons in an atom of Selenium whose mass number is 79.
- 30. Chemistry Humor • A neutron walked into a restaurant and asked how much for a drink. The waiter replied, "for you, no charge."
- 31. Isotopes • The atoms of a given element must contain the same number of protons, but the number of neutrons can vary.
- 32. Isotopes • For Example: Most atoms of hydrogen contain 1 proton and no neutrons 1 1H • But some contain 1 proton and 1 neutron 1 2H • Still others contain 1 proton and 2 neutrons 1 3H • All three are still atoms of hydrogen
- 33. Isotopes • Isotopes are atoms of the same element that have different numbers of neutrons • Thus, they have different mass numbers.
- 34. Isotope Symbols • Isotopes can be identified by using a symbol that indicates both the element and its mass number • Examples C-14 14C Carbon-14 6 14C
- 35. Mass Number • The mass number must be an integer. mass number = atomic number + neutrons
- 36. Atomic Masses • Why are the atomic masses on the periodic table fractional values (i.e. not integers)? • Because… the atomic masses are the average mass of all the naturally occurring isotopes in a sample of the element
- 37. Sample Problem • Atomic mass of carbon • on overhead • Lab - “Modeling Isotopes” • Page • Homework: Review Questions 13-20
- 38. Location of Electrons - Energy Levels • Each electron has its own distinct amount of energy that corresponds with the energy level it occupies. • Electrons can gain or lose energy and move to a different energy level, but they do so in a unique way….
- 39. • Electrons can only absorb a “correct” amount of energy that allows it to move to a higher energy level • These “packets of energy” are photons of light. • Different colors of light carry different amounts of energy
- 41. • By convention there is color, by convention sweetness, by convention bitterness, but in reality there are atoms and space. - Democritus (400 BC)
- 42. Name this compound: •Ba(Na)2 •answer: Banana.
- 43. •When electrons occupy the lowest available orbital, the atom is said to be in the ground state.
- 44. • When electrons are subject to heat, light, or electricity, an electron may absorb energy and (temporarily) move to a higher energy level. This unstable condition is called an excited state.
- 45. • When the electron returns to a lower level it emits energy in the form of infrared, ultraviolet, or visible light. • While the light appears as one color to our eyes, it is actually composed of many different wavelengths (or colors of light)
- 46. • Because each atom has its own distinct orbital energy levels, each atom has its own distinct pattern of emission lines (also known as bright line spectrum), that can be used to identify elements. • Astronomy
- 47. The Bohr Model • Bohr built upon spectroscopic observations of atoms. Spectroscopists noticed that an atom can only absorb certain energies (colors) of light (the absorption spectrum) and once excited can only release certain energies (the emission spectrum) and these energies happen to be the same. Bohr used these observations to argue that the energy of a bound electron is "quantized." Absorption Spectrum Emission Spectrum
- 48. The Bohr Model • In the animation, you will see a model of a Hydrogen atom and to the right of it, a Bohr energy level diagram. • In the animation you will notice that if the energy of the photon of light is just right, it will cause the electron to jump to a higher level. • When the electron jumps back down, a photon is created for each jump down. • A photon without the right amount of energy (the pink one) passes through the atom with no effect. • Photons with too much energy will cause the electron to be ejected which ionizes the atom. An ionized electron is said to be in the n=infinity energy level. • Keep in mind that these rings are not actually orbits, but are levels that represent the location of an electron wave. The number n corresponds to the number of complete waves in the electron. Get out your reference tables and find the hydrogen and mercury energy level diagrams.
- 49. The Bohr Model • Ephoton = Einitial - Efinal (reference tables) • This formula can be used to determine the energy of the photon emitted (+) or absorbed(-).
- 50. Sample Problem • Calculate the energy of the photon that is emitted when a hydrogen atom changes from energy state n=3 to n=2. What color corresponds with the photon emitted? • Solution • From reference tables… E3 = Einitial = -1.51 eV E2 = Efinal = -3.40 eV Ephoton = Einitial – Efinal = (-3.40 eV) – (-1.51 eV) = -1.89 eV Ephoton = hf ==== f = Ephoton / h But we need Ephoton in Joules, because Planck’s constant is in Joules Ephoton = (-1.89 eV) (1.60x10-19J / eV) = 3.02x10-19 J f = (3.02x10-19 J) / (6.63x10-34 J∙s) = 4.56x1014 Hz From reference tables, this frequency corresponds with red light.
- 51. The Bohr Model • Summary of the Bohr Model • All forms of energy are quantized. An electron can gain or lose kinetic energy only in fixed amounts, or quanta. • The electron in the hydrogen atom can occupy only certain specific orbits of fixed radius and no others. • The electron can jump from one orbit to a higher one by absorbing a quantum of energy in the form of a photon. • Each allowed orbit in the atom corresponds to a specific amount of energy. The orbit nearest the nucleus represents the smallest amount of energy that the electron can have.
- 52. Bohr Model Vocabulary • When the electron is in the lowest energy level (n=1), it is said to be in the ground state. • An electron in any level above the ground state is said to be in an excited state. • A spectral line is a particular frequency of absorbed or emitted energy characteristic of an atom
- 53. The Cloud Model • The cloud model represents a sort of history of where the electron has probably been and where it is likely to be going. • The red dot in the middle represents the nucleus while the red dot around the outside represents an instance of the electron. • Imagine, as the electron moves it leaves a trace of where it was. This collection of traces quickly begins to resemble a cloud. • The probable locations of the electron predicted by Schrödinger's equation happen to coincide with the locations specified in Bohr's model.