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Unit 2 – The Atom (Chapter 3)

Unit 2 – The Atom (Chapter 3). nucleus – center of an atom containing __________ and _______________ Electrons surround the nucleus in a region known as the electron cloud .

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Unit 2 – The Atom (Chapter 3)

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  1. Unit 2 – The Atom (Chapter 3) nucleus – center of an atom containing __________ and _______________ Electrons surround the nucleus in a region known as the electron cloud. The nucleus is very small compared to the size of the atom. The diameter of the nucleus is 20,000 times smaller than the diameter of the atom (electron cloud). Most of the volume of an atom is ______________________.

  2. Subatomic Particles • protons – ____________ charge • Think • neutrons – ______ charge or ___________ charge • Think • electrons – ______________ charge

  3. Atoms and the Periodic Table (pg 77) atomic number – number of ______________ in an atom The atomic number identifies the element. No two elements have the same atomic number. mass number – number of ____________ plus number of _____________ ___________________ have such small mass that they are neglected in atomic mass calculations.

  4. Numbers vs. Mass (pg 80) Atomic number and mass number are numbers of particles NOT mass. For mass, the standard is the carbon-12 atom. atomic mass unit (amu) – exactly 1/12 the mass of a carbon-12 atom 1 amu = 1.661 x 10-27 kg proton mass = 1.007276 amu neutron mass = 1.008665 amu Protons and neutrons have different masses so two elements with the same mass number will have different masses.

  5. Average Atomic Mass weighted average = mass x percent abundance ex. average test grade mass number = protons + neutrons always a whole number average atomic mass is the weighted average of isotopes of an element rarely a whole number

  6. Average Atomic Mass Two isotopes of a certain element occur in nature. One is 19.80% abundant and has an atomic mass of 10.00 amu. The other is 80.20% abundant and has an atomic mass of 11.01 amu. Calculate the average atomic mass of this element and determine the identity of the element.

  7. Photoelectric Effect(pg 99) Initially, it was thought that light was waves only. In this theory, light waves at any frequency should be able to knock off an electron from metal if the light intensity is high enough. However, experiments showed that a minimum frequency is required to knock off an electron and nothing happens below that frequency regardless of the light’s intensity. Einstein proposed that light has a dual wave-particle nature – it acts as a wave and a particle at the same time. photon – particle of electromagnetic radiation (light) possessing a quantum of energy

  8. Photoelectric Effect(pg 99) Ephoton = hn where Ephoton = energy of a photon h = Planck’s constant n = frequency If the photon (light) has a frequency that is too low, then its energy is too low as well which means that it will not be able to knock off an electron. If the frequency is high enough, then the energy is high enough also so the electron will get knocked off. You do not have to be able to use this equation but be aware that increasing frequency means increasing energy.

  9. Electron Orbits Niels Bohr proposed that electrons circle the nucleus only in certain paths each of a specific level of energy. Farther out from the nucleus is higher energy. ground state – atom with electrons in lowest energy state excited state – atom with electrons gaining energy above ground state This is related to electron energy levels you may have done in physical science.

  10. Electron Orbits The ground state energy level is E1. The excited states are E2, E3, E4, and so on. Each element has its own unique ground state and many different excited states. emission – a photon is emitted when an electron falls to a lower energy level energy = Ephoton = E2 – E1 absorption – a photon is absorbed when an electron goes to a higher energy level energy = Ephoton = E1 – E2 The equation Ephoton = hn still holds so the electrons still exist only at certain frequencies and, therefore, certain energy levels. That’s why they jump directly from E1 to E2 – there is no middle ground.

  11. Electrons as Waves(pg 104) A French guy named Louis de Broglie finally put it together that electrons behave both as particles and as waves. Particles mean that they have mass while waves mean that they exist only at certain frequencies (and therefore energy levels) around the nucleus. This is called dual wave-particle theory and gets into some really weird quantum mechanics.

  12. Heisenberg Uncertainty Principle Electrons are detected using photons (light) but electrons are so small that the photons used to detect them actually knock the electrons off course. The more photons you hit the electrons with allows you to detect them more easily but the presence of the photons themselves makes the electrons act in an unnatural way. If you back off with the photons, the electrons act more normally but then it is more difficult to detect the electrons with the lower photon level. Bottom line is that you can’t know everything about an object at the same time. Red light, green light…

  13. Energy Levels Electrons exist as waves which means that they can only exist at certain energy levels in the electron cloud. This is why there are energy levels in an atom. Think of the rungs of a ladder or a stool… Remember how you find the energy level of an element?

  14. Orbitals(pg 111) INCREASING ENERGY orbital The 1st energy level has 1 sublevel: 1s (with 1 orbital). The 2nd energy level has 2 sublevels: 2s (with 1 orbital) and 2p (with 3 orbitals). The 3rd energy level has 3 sublevels: 3s (with 1 orbital), 3p (with 3 orbitals), and 3d (with 5 orbitals). and so on… energy level sublevel

  15. Filling in Electrons aufbau principle (pg 111) – electrons occupy the ______________ energy orbital available From the previous page, electrons will occupy the 1s orbital before the 2s orbital, the 2s orbital before the 2p orbital, and so on... Note, however, that the 4s orbital is lower energy than the 3d orbitals. Therefore, electrons will actually fill the 4s orbital first. To help with this, there’s yet another diagram…

  16. Figure 19 on pg 116 • Write the s sublevels down the first column, the p sublevels down the second column keeping the energy levels in the same row. Continue with the d and f sublevels in the same fashion. • Go from upper right to lower left. Therefore, the order to fill the orbitals with electrons is as follows: • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, etc.

  17. Filling in Electrons Pauli exclusion principle (pg 112) – max of two electrons can occupy one orbital and they must have opposite spins means spin in one direction means spin in the opposite direction So to fill one orbital, the electrons must look like this:

  18. Filling in Electrons Hund’s rule (pg 112) – an orbital cannot hold two electrons until each orbital in that sublevel has at least one electron For example, fill in the following p orbitals with 4 electrons:

  19. Orbitals and Quantum Numbers(pg 107) principal quantum number – indicates major energy levels of an atom: symbol is n aka “shell” This is EXACTLY the same as the period on the periodic table. What are the values of n for bromine? What about for iridium?

  20. Orbitals and Quantum Numbers(pg 107) angular momentum quantum number – indicates sublevel (aka subshell) within major energy level: symbol is l The values of l are any whole number between ____ and ____ Energy level 1 (n = 1), Energy level 2 (n = 2), Energy level 3 (n = 3), and so on… l = 0 is the ____ subshell l = 1 is the ____ subshell l = 2 is the ____ subshell l = 3 is the ____ subshell

  21. Notation so far… n = 1 has 1 sublevel l = 0 (s) n = 2 has 2 sublevels l = 0, 1 (s, p) n = 3 has 3 sublevels l = 0, 1, 2 (s, p, d) n = 4 has 4 sublevels l = 0, 1, 2, 3 (s, p, d, f) The p sublevel for the 2nd energy level is labeled as 2p. The d sublevel for the 4th energy level is labeled as 4d. All others work the same way.

  22. Orbitals and Quantum Numbers Each sublevel has orbitals which are specifically where the electrons are located. magnetic quantum number – indicates number of orbitals: symbol is ml ml can be any whole number between -l andl s sublevel (l = 0) p sublevel (l = 1) d sublevel (l = 2) f sublevel (l = 3) Table 2 on page 110 is a good summary.

  23. Orbitals and Quantum Numbers Electrons in an orbital spin around on an axis like the planets do. spin quantum number – indicates spin state (spin direction) of an electron This number can be either -1/2 or +1/2. Each orbital contains a max of two electrons and each must spin in opposite directions, i.e. one is -1/2 and the other is +1/2.

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