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Chapters 4 & 5 Atomic Theory. Section 4.3 Atomic Number The key distinguishing trait between different elements is the number of protons in the nucleus.

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Presentation on theme: "Chapters 4 & 5 Atomic Theory. Section 4.3 Atomic Number The key distinguishing trait between different elements is the number of protons in the nucleus."— Presentation transcript:

1 Chapters 4 & 5 Atomic Theory

2 Section 4.3

3 Atomic Number The key distinguishing trait between different elements is the number of protons in the nucleus. The “atomic number” (Z) is the number of protons in an element Consider tin (atomic symbol “Sn”). If a neutral tin atom contains 50 protons, how many electrons does it have?

4 Mass Number The total number of protons and neutrons in an atom is called the “mass number”. The atomic number of gold (Au) is 79. If the mass number for gold is 197, how many neutrons does a typical gold atom contain? – a. 197 b. 79 c. 118 d. 122

5 Chemical Symbols The complete symbol for an element on the Periodic Table includes the atomic symbol, the mass number, and the atomic number. Note: you can also refer to atoms by using the mass number and the name of the element, i.e. gold-197. Superscript → Subscript → Mass number Atomic number X

6 Comprehension Check Identify the atomic number and mass number in the symbol for bromine below. Then, find the number of protons, neutrons and electrons in a typical bromine atom. Draw the complete chemical symbol for the element that has the atomic number 19. What is the name of this element? A scientist has a mystery atom with 47 protons and 62 neutrons. What element is this? How is this atom unique? Use your Periodic Table to answer these questions. Br 80 35

7 Isotopes

8 The Three Known Hydrogen Isotopes

9 Atomic Mass Recall, the number of protons plus the number of neutrons in an atom is the mass number. The atomic mass of an element is the weighted average mass of all naturally occurring isotopes of that element. » Math Check: What is the difference between a weighted average and an arithmetic mean? – The weighted average reflects both the mass and the percent abundance of the isotopes.

10 Measuring Atomic Mass Scientists do not measure the mass of single atoms in grams. – The numbers would always be tiny and complicated Mass of 1 proton = 1.673x10 -24 grams Instead of grams, we use atomic mass units based on the mass of a carbon-12 atom. – By definition 1 amu = 1/12 mass of carbon-12 atom. Why carbon-12? – High isotope purity

11 Isotopes of Carbon Carbon = 12.011

12 Calculating a Weighted Average Naturally ocurring iron (Fe) has 4 isotopes: iron- 54 (5.845%), iron-56 (91.754%), iron-57 (2.119%), and iron-58 (0.282%). Find iron’s atomic mass by computing the weighted average of these four isotopes. 54 x 5.845% = 3.156 56 x 91.754% = 51.382 57 x 2.119% = 1.208 58 x 0.282% = 0.164 Total = 55.91 amu

13 Preview of the Periodic Table A periodic table is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties. The periodic table allows you to easily compare the properties of one element to another. Each horizontal row is called a period. Each vertical column is called a group, or family. – Elements in a group typically have similar physical and chemical properties

14 Section 5.1

15 Atomic Orbitals Principal Quantum Number (n) = the energy level of the electron: 1, 2, 3, etc. These are called atomic orbitals (coined by scientists in 1932) - regions where there is a high probability of finding an electron. Sublevels- like theater seats arranged in sections: letters s, p, d, and f

16 Principal Quantum Number Generally symbolized by “n”, it denotes the shell (energy level) in which the electron is located. Maximum number of electrons that can fit in an energy level is: 2n 2

17 Summary s p d f # of shapes (orbitals) Maximum electrons Starts at energy level 1 2 1 3 6 2 5 10 3 7 14 4

18 By Energy Level First Energy Level Has only s orbital only 2 electrons 1s 2 Second Energy Level Has s and p orbitals available 2 in s, 6 in p 2s 2 2p 6 8 total electrons

19 By Energy Level Third energy level Has s, p, and d orbitals 2 in s, 6 in p, and 10 in d 3s 2 3p 6 3d 10 18 total electrons Fourth energy level Has s, p, d, and f orbitals 2 in s, 6 in p, 10 in d, and 14 in f 4s 2 4p 6 4d 10 4f 14 32 total electrons

20 By Energy Level  Any more than the fourth and not all the orbitals will fill up. You simply run out of electrons The orbitals do not fill up in a neat order. The energy levels overlap Lowest energy fill first.

21 Section 5.2 Electron Arrangement in Atoms OBJECTIVES: Describe how to write the electron configuration for an atom.

22 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f aufbau diagram - page 133 Aufbau is German for “building up”

23 Electron Configurations… …are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how: 1) Aufbau principle - electrons enter the lowest energy first. This causes difficulties because of the overlap of orbitals of different energies – follow the diagram! 2) Pauli Exclusion Principle - at most 2 electrons per orbital - different spins

24 Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli To show the different direction of spin, a pair in the same orbital is written as:

25 Electron Configurations 3) Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until they have to. Let’s write the electron configuration for Phosphorus  We need to account for all 15 electrons in phosphorus

26 The first two electrons go into the 1s orbital Notice the opposite direction of the spins only 13 more to go... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

27 The next electrons go into the 2s orbital only 11 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

28 The next electrons go into the 2p orbital only 5 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

29 The next electrons go into the 3s orbital only 3 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

30 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into separate shapes (Hund’s) 3 unpaired electrons = 1s 2 2s 2 2p 6 3s 2 3p 3 Orbital notation

31 Orbitals fill in an order Lowest energy to higher energy. Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. However, half filled orbitals have a lower energy, and are next best Makes them more stable. Changes the filling order

32 Section 5.3 Physics and the Quantum Mechanical Model OBJECTIVES: Describe the relationship between the wavelength and frequency of light.

33 Section 5.3 Physics and the Quantum Mechanical Model OBJECTIVES: Identify the source of atomic emission spectra.

34 Section 5.3 Physics and the Quantum Mechanical Model OBJECTIVES: Explain how the frequencies of emitted light are related to changes in electron energies.

35 Light The study of light led to the development of the quantum mechanical model. Light is a kind of electromagnetic radiation. Electromagnetic radiation includes many types: gamma rays, x-rays, radio waves… Speed of light = 2.998 x 10 8 m/s, and is abbreviated “c” All electromagnetic radiation travels at this same rate when measured in a vacuum

36 - Page 139 “R O Y G B I V” Frequency Increases Wavelength Longer

37 Parts of a wave Wavelength Amplitude Origin Crest Trough

38 Equation: c = c = speed of light, a constant (2.998 x 10 8 m/s) (nu) = frequency, in units of hertz (hz or sec -1 ) (lambda) = wavelength, in meters Electromagnetic radiation propagates through space as a wave moving at the speed of light.

39 Wavelength and Frequency Are inversely related As one goes up the other goes down. Different frequencies of light are different colors of light. There is a wide variety of frequencies The whole range is called a spectrum

40 - Page 140 Use Equation: c =

41 Radiowave s Microwave s Infrared. Ultra- violet X-Rays GammaRays Low Frequency High Frequency Long Wavelength Short Wavelength Visible Light Low EnergyHigh Energy

42 Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table

43 Atomic Spectra White light is made up of all the colors of the visible spectrum. Passing it through a prism separates it.

44 If the light is not white By heating a gas with electricity we can get it to give off colors. Passing this light through a prism does something different.

45 Atomic Spectrum Each element gives off its own characteristic colors. Can be used to identify the atom. This is how we know what stars are made of.

46 These are called the atomic emission spectrum Unique to each element, like fingerprints! Very useful for identifying elements

47 Light is a Particle? Energy is quantized. Light is a form of energy. Therefore, light must be quantized These smallest pieces of light are called photons. Energy & frequency: directly related.

48 Equation: E = h E = Energy, in units of Joules (kg·m 2 /s 2 ) (Joule is the metric unit of energy) (Joule is the metric unit of energy) h = Planck’s constant (6.626 x 10 -34 J·s) = frequency, in units of hertz (hz, sec -1 ) = frequency, in units of hertz (hz, sec -1 ) The energy (E ) of electromagnetic radiation is directly proportional to the frequency ( ) of the radiation.

49 The Math in Chapter 5 There are 2 equations: 1) c = 2) E = h Know these!

50 Explanation of atomic spectra When we write electron configurations, we are writing the lowest energy. The energy level, and where the electron starts from, is called it’s ground state - the lowest energy level.

51 Changing the energy Let’s look at a hydrogen atom, with only one electron, and in the first energy level.

52 Changing the energy Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be “ excited ”

53 Changing the energy As the electron falls back to the ground state, it gives the energy back as light

54 They may fall down in specific steps Each step has a different energy Changing the energy

55 What is light? Light is a particle – it contains discrete photons (it comes in chunks) Light is a wave - we can measure its wavelength and it behaves as a wave If we combine E=mc 2, c=, E = 1/2 mv 2 and E = h, then we can get: = h/mv (from Louis de Broglie) called de Broglie’s equation Calculates the wavelength of a particle.

56 The physics of the very small Quantum mechanics explains how very small particles behave Quantum mechanics is an explanation for subatomic particles and atoms as waves Classical mechanics describes the motions of bodies much larger than atoms

57 Wave-Particle Duality J.J. Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave!

58 First paradox! Light exists and acts as both a particle and a wave Each nature of light can be measured and observed Light is both at the same time!!!

59 Heisenberg Uncertainty Principle It is impossible to know exactly the location and velocity of a particle simultaneously. The better we know one, the less we know the other. Measuring changes the properties. True in quantum mechanics, but not classical mechanics

60 Heisenberg Uncertainty Principle You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! “One cannot simultaneously determine both the position and momentum of an electron.” Werner Heisenberg

61 It is more obvious with the very small objects To measure where a electron is, we use light. But the light energy moves the electron And hitting the electron changes the frequency of the light.

62 Moving Electron Photon Before Electron velocity changes Photon wavelength changes After Fig. 5.16, p. 145

63 Second Paradox!!! It is impossible to measure both the velocity and position of an electron at the same time. It is only possible to measure or observe one or the other Heisenberg Uncertainty Principle Science is cray!!!

64 Erwin Schrödinger and his cat Austrian Physicist (1887-1961) Challenged the paradox of the particle/wave duality Created famous thought experiment about a cat in a box

65 Schrödinger’s Cat

66 Your assignment!! Work within your table groups Develop a model that could be used to explain the Heisenberg Uncertainty Principle to a student who is unfamiliar with quantum physics. Assume the student is familiar with atomic structure You can include objects that we don’t actually have (like ping pong balls or string or… BE CREATIVE!!) You will describe the activity, including the materials needed You will explain how the activity demonstrates the Heisenberg Uncertainty Principle, as well as a summary of the principle itself.

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