Bell-Ringer. Define the term isotope. [2 marks] Answer per IB Mark scheme:

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2 Bell-Ringer Define the term isotope. [2 marks] Answer per IB Mark scheme: atoms of the same element/same number of protons/same atomic number that have different numbers of neutrons/different mass numbers Award only [1] max if reference made to elements but not atoms.

3 TEST TOPIC 2 TEST IS NEXT WEEK Wednesday Oct 19 Topic 2 Answers PDF with answers to all Topic 2 exercises from the book has been uploaded. RESOURCE FOR STUDYING.

4 Tricky IB Question for RAM #1 Iridium (Ir) has a relative atomic mass of and consists of Ir-191 and Ir-193 isotopes. Calculate the percentage composition of a naturally occurring sample of iridium.

5 Iridium (Ir) has a relative atomic mass of and consists of Ir-191 and Ir-193 isotopes. Calculate the percentage composition of a naturally occurring sample of iridium. Answer: Only use the following formula when you are dealing with only two isotopes, we can deduce the percent abundance of the heavier isotope, and then subtract that value from 100 to get the percent abundance of the lighter isotope. Ar - mass number of lighter isotope x 100% = % of heavier isotope Difference in mass number of two isotopes x 100% = 61% Therefore, there exists 61% of Ir-193 and 39% of Ir-191.

6 Tricky IB Question for RAM #2 Magnesium has three stable isotopes, ²⁴Mg, ²⁵Mg, and ²⁶Mg. The lightest isotope has an abundance of 78.90%. Calculate the percentage abundance of the other two isotopes. Answer: Magnesium s RAM according to the periodic table is = (24 x 78.90) + (25 x) + [26 ( x)] 100 2,431 = x x 2,431 = x x = 11.2% abundance for Mg-25, therefore 9.90% abundance for Mg-26

7 Mass Spectrometer The proportion of each isotope present in a sample of an element can be measured using an instrument called a mass spectrometer. The readout from a mass spectrometer is called a mass spectrum. In a mass spectrum of an element, we get one peak for each individual isotope. The height of each peak is proportional to the number of atoms of this isotope in the sample tested. Mass spectrum for Rubidium (Rb)

8 Test Yourself Deduce the relative atomic mass of the element rubidium from the data given in the figure. Mass spectrum for Rubidium (Rb)

9 Test Yourself Deduce the relative atomic mass of the element rubidium from the data given in the figure. Answer: Pretend there is a sample of 100 atoms, meaning that 77 out of the 100 are Rb-85 atoms and the remaining 23 are Rb-87 atoms. (85 x 77) + (87 x 23) Mass spectrum for Rubidium (Rb)

10 Niels Bohr s Experiment Niels Bohr was a physicist interested in further understanding the structure of the atom. Bohr subjected a discharge tube filled with hydrogen gas to a very high voltage, and then observed that the gas began to emit a pinkish light.

11 Bohr broke the light up using a spectroscope, which contains a diffraction grating and separates the various wavelengths of light, such as with a prism. Glass Prism breaking white light into continuous spectrum

12 These discrete lines in hydrogen s line emission spectra data represent the different frequencies/wavelengths of light emitted by hydrogen. Feeding energy to an atom can cause its electrons to be promoted to higher energy levels in the atom. The electron, however, is unstable in this higher level and will fall to a lower energy level. As it returns from a level at energy E₂ to E, the extra energy (E₂ - E₁) is given out in the form of a photon of light. This contributes to a line in the emission spectrum of an atom. The fact that a line spectrum is produced provides evidence for electrons being in energy levels (shells); i.e. electrons in an atom are allowed to have only certain amounts of energy.

13 Types of Emission Spectras Continuous Spectrum: unbroken sequence of frequencies, such as the spectrum of visible light when it is passed through a prism. Line Emission Spectrum: only shows the frequencies of light that are emitted by an atom; i.e. the top figure shows the frequencies of visible light emitted by a hydrogen atom. Line Absorption Spectrum: does NOT show the frequencies of light that were emitted by an atom, but instead show the frequencies of the light that were absorbed by the atom; i.e. the absorption figure shows all frequencies of visible light that were NOT absorbed by hydrogen.

14 Before we discuss the Bohr model further, let s consider the types of light emitted by atoms. Electrons can indeed emit light. All atoms emit their own spectra of light, and that spectra. The electromagnetic spectrum describes all the types of LIGHT in our universe. These are called electromagnetic radiation or electromagnetic waves.

15 These are the emission spectra of each of the respective elements on this list. Emission spectra shows you the frequencies/wavelengths of electromagnetic radiation that are emitted from the atoms of an element. These are specifically showing the frequencies/wavelengths of visible light that are emitted by atoms of an element.

16 Wavelength ( ): is the distance between the crests, or the distance between troughs. Frequency ( ): is the number of wave cycles to pass during a given unit of time; i.e. the amount of waves passing in a 1 second period. Amplitude: is the wave s height from zero to the crest Crest: the top of the wave Trough: the bottom of the wave

17 High frequency Short wavelength High energy Low frequency Long wavelength Low energy -Wavelength is inversely proportional to frequency. -Frequency is directly proportional to energy.

18 You should be able to identify: -The relative wavelengths and frequencies of all electromagnetic waves, including those of visible light. -How frequency, wavelength, and the speed of light are related: c= All of these waves travel at the same speed, the speed of light (c), despite having different wavelengths/frequencies.

19 Practice 1. Arrange the following in order of increasing wavelength: red light, yellow light, infrared radiation, ultraviolet radiation. 2. Arrange the following in order of increasing energy: UV radiation, red light, infrared radiation, green light

20 Answers: 1. ultraviolet radiation; yellow light; red light; infrared radiation 2. infrared radiation; red light; green light; ultraviolet radiation

21 Bohr Model of the Atom This lead Bohr to conclude that electrons travel along well defined orbits, and that electrons emit light when falling back down to their ground state orbitals, producing a specific emission spectrum. Bohr was correct about this; however, when asked to explain why each atom had a different emission spectrum he concluded it must be because each atom has its own special energy orbitals. We know today this is not true, all atoms have the same layout, just different configurations. In order to develop the model of the atom further, we needed to reconsider the nature of light and matter.

22 Heisenberg s Uncertainty Principle: we cannot know where an electron is at any given moment in time - the best we can hope for is a probability picture of where the electron is likely to be. Wave-Particle Duality: electrons actually behave as both particles and waves. Waves are in continuous motion, and therefore are not localized to a specific position in space. The possible positions of an electron are spread out in space in the same way that a wave is spread across a water surface.

23 We cannot find an electron even if it is behaving as a particle. Recall that the mere act of locate an object in space (whether via naked eye or instrument) requires light to reflect off of that object. The subatomic particles are so small that light reflecting off of them will cause their trajectory to change.

24 Why can t we observe an electron directly? Why do we believe electrons have particle-wave duality? VIDEO: QUANTUM WORLD Figure: beam of electrons going through a slit produces a diffraction pattern, a wave pattern, instead of a straight line. Evidence for particle-wave duality of subatomic particles.

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26 Can we see atoms today? Well, kind of... Video: How Can You See an Atom?

27 Schrödinger and the Electron Cloud -Based on Heisenberg s uncertainty principle, Erwin Schrödinger in 1926 developed a mathematical equation to figure out where an electron might be located within the electron cloud based on its level of energy. -He knew that an electron s location within the electron cloud could not be known for certain- but a probability of where an electron might be could be calculated! -We call areas of high probability of where an electron might be an atomic orbital.

28 Atomic orbitals describe the three-dimensional areas where there is a high probability that the electron will be located. They are regions around an atom s nucleus in which there is a 90% probability of finding the electron. Shapes of orbitals will depend on the energy of the electron and its motion in that area. When an electron is in an orbital of higher energy, it means it is most likely far from the nucleus. s orbitals are spherical p orbitals are dumbbell-shaped. There are three of them, located on an x, y, and z axis.

29 Current Model: Quantum Mechanical Model Schrodinger is a huge contributor to our current model of the atom, the Quantum Mechanical Model. Protons and neutrons are located in the nucleus, or center, of the atom. Protons and neutrons are roughly the same size in mass. Together, protons and neutrons are referred to as nucleons. Electrons are much smaller than protons and neutrons, and are located at a very far distance from the nucleus in areas of high probability called atomic orbitals, which are within the electron cloud.

30 Current Atomic Model: Quantum Mechanical Model Incorrect: Bohr thought the energy levels around a nucleus were predictable orbits that electrons traveled on; he figured that electrons had fixed positions in an atom. He believed these orbits were equidistant from each other. Correct: Electrons occupy areas of space called atomic orbitals within the area we call the electron cloud. These orbitals have different levels of energy depending, in part, on their proximity to the nucleus. These orbitals are not equidistant from each other.

31 The energy levels in an atom and all the orbitals that pertain to those energy levels are not equidistant from each other. Energy levels closer to the nucleus are further away from each other. Energy levels further away from the nucleus are closer to each other. WRONG!

32 Energy levels increase in energy the farther away they are from the nucleus. Each energy level is sometimes referred to as a principal quantum number, and is given the variable, n. For instance, the first energy level (the level nearest to the nucleus) would be referred to as n=1, or principal quantum number 1.

33 Due to the unequal distances between energy levels in an atom, different frequencies of light can be emitted from an atom. This figure illustrates a hydrogen atom.

34 Hydrogen produces visible light when its excited electron falls to the second energy level (n = 2). It doesn t matter where it jumped to when excited; if it falls and takes a pit stop at energy level 2, then it will release visible light. It releases Infrared light if it lands on n=3. If its ground state is n=1 then it will release UV radiation once it returns there. Notice that the electron travels a longer distance to get back to n=1 versus n=2. This explains why it releases UV radiation, which is much stronger than visible light.

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37 Test Yourself a) b) c) d) A B C D

38 Test Yourself Answer: b) B