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Chapter 5 Electrons in Atoms

Chapter 5 Electrons in Atoms. Ms. Wang Lawndale High School. Section 5.1 – Models of the Atom. In 1897 J. J. Thomson discovered the electron. Observed that a magnet deflected the straight paths of the cathode rays.

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Chapter 5 Electrons in Atoms

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  1. Chapter 5Electrons in Atoms Ms. Wang Lawndale High School

  2. Section 5.1 – Models of the Atom In 1897 J. J. Thomson discovered the electron Observed that a magnet deflected the straight paths of the cathode rays

  3. Atoms were known to be electrically neutral which meant that there had to be some positively charged matter to balance the negative charges

  4. Ernest Rutherford’s experiment disproved the plum pudding model of the atom and suggested that there was a positively charged nucleus because many of the alpha particles hit the thin gold foil and bounced back BUT, Rutherford’s atomic model could not explain the chemical properties of elements

  5. The Bohr Model In 1913, Niels Bohr came up with a new model (Bohr was a student of Rutherford) He noticed that light given out when atoms were heated always had specific amounts of energy, so he proposed that electrons in an atom must be orbiting the nucleus and can reside only in fixed energy levels.

  6. Energy Levels • Energy levels – fixed energy that an electron can have • This is similar to steps of a ladder • Quantum – amount of energy required to move an electron from one energy level to another energy level (to be quantized)

  7. The Quantum Mechanics View of the Atom The quantum mechanical model that scientist use today does not describe the exact path an electron takes around the nucleus but more concerned with the probability of finding an electron in a certain place.

  8. Atomic Orbitals • Atomic Orbitals – a region of space in which there is a high probability of finding an electron • Each energy sublevel corresponds to an orbital of different shape describing where the electron is likely to be found

  9. Labeling Electrons in Atoms • Each electron in an atom is assigned a set of four quantum numbers. These help to determine the highest probability of finding the electrons. • Three of these numbers (n, l, m) give the location of the electron • The fourth (s) describes the orientation of an electron in an orbital.

  10. Quantum letters can be thought of like the numbers and letters on a concert ticket

  11. Labeling Electrons in Atoms

  12. n= principal quantum number • Used to describe the energy of the electron. The farther away from nucleus, the higher the energy • The n quantum number can have values = 1, 2, 3, …. n n = 1 can hold 2 electrons n = 2 can hold 8 electrons n = 3 can hold 18 electrons n = 4 can hold 32 electrons

  13. Nucleus Draw the electron shell diagram for Beryllium. Be has 4 electrons Be Electrons Draw the electron shell diagram for Nitrogen. N has 7 electrons N

  14. Draw the electron shell diagrams for these elements • Nickel • Aluminum • Argon • Carbon • Calcium • What does n represent? • How many electrons can each n hold?

  15. l = sublevel • Provides a code for the shape of orbitals • They are designated by letters • l =0, 1, 2, (n-1)

  16. Answer these questions • If n = 1 what does l =? Which letter does that correspond to? • If n = 2 what does l = Which letter does that correspond to? • If n = 3 what does l =? Which letter does that correspond to? • If n= 4 what does l =? Which letter does the correspond to?

  17. For principal energy level 3, there are 3 sublevels s < p< d <f in energy

  18. m=magnetic quantum number • Used to describe each orbital within a sublevel

  19. Section 5.2 – Electron Configurations • Each orbital holds 2 electrons • Filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p • Example He = 2 electrons 1s2 • Example Li = 3 electrons 1s22s1 • Example B = 5 electrons 1s22s22p1

  20. Practice Problems Write electron configurations for the following atoms Li 5. P N 6. Si Be 7. Mg C 8. Al

  21. Electron Configurations can be written in terms of noble gases To save space, configurations can be written in terms of noble gases • Example 1: Ne = 1s22s22p6 S = 1s22s22p63s23p4 Or S = [Ne] 3s23p4 • Example 2: Ar = 1s22s22p63s23p6 Mn = 1s22s22p63s23p64s23d5 Mn = [Ar] 4s23d5

  22. Reading the Periodic Table

  23. Locating Electrons in Atoms So far we have discussed 3 quantum numbers • n= principal quantum level (principal energy level) • l= Sublevel • m = magnetic quantum number (shape of orbitals) 1s2 n Number of electrons in sublevel l

  24. s = spin • When an electron moves, it generates a magnetic field. • s describes the direction of electron spin around its axis. • They must spin in opposite directions • Spin= up down • There are two values of s: +1/2 and -1/2

  25. Orbital Diagrams • The electron configuration gives the number of electrons in each sublevel but does not show how the orbitals of a sublevel are occupied by the electrons.

  26. Orbital Diagrams • They are used to show how electrons are distributed within sublevels. • Each orbital is represented by a box and each electron is represented by an arrow. • The direction of the spin is represented by the direction of the arrow Example: Boron 1s22s22p1 2p 2s 1s

  27. Orbital Diagrams Steps to writing orbital diagrams:ex F (Z=9) • Write the electron configuration 1s22s22p5 2. Construct an orbital filling diagram using boxes for each orbital 3. Use arrows to represent the electrons in each orbital. 2p 2s 1s 2p 2s 1s

  28. Aufbau Principle • Electrons must occupy the orbital with the lowest energy first • Example: Oxygen 1s22s22p4 2p 2p 2s 2s 1s 1s

  29. Pauli Exclusion Principle • An atomic orbital may describe at most two electrons • The 2 electrons must have opposite spins • Example: Oxygen 1s22s22p4 2p 2p 2s 2s 1s 1s

  30. Hund’s Rule • Orbitals of equal energy are each occupied by one electron before any pairing occurs • Example: Oxygen 1s22s22p4 2p 2p 2s 2s 1s 1s

  31. Draw orbital diagrams for these elements • Li 5. P • N 6. Si • Be 7. Mg • C 8. Al

  32. Section 5.3 - Atomic Spectra • When atoms absorb energy, electrons move into higher energy levels • These electrons lose energy by emitting light when they return to lower energy levels • Atomic Emission Spectrum – the discrete lines representing the frequencies of light emitted by an element

  33. Atomic Spectra • Each discrete line in an emission spectrum corresponds to one exact frequency of light emitted by the atom • Ground State – lowest possible energy of the electron in the Bohr model • The light emitted by an electron moving from higher to a lower energy level has a frequency directly proportional to the energy change of the electron

  34. Homework Chapter 5 Assessment Page 148 #’s 22-24, 27, 29, 30-39, 50-53, 57, 60, 68, 70-72

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