Chapter 5

1. Chapter 5 Electrons in Atoms

2. Models of the Atom Section 1

3. The Development of Atomic Models • Elements • Rutherford’s atomic model couldn’t explain the chemical properties of elements.

4. The Bohr Model • Niels Bohr (Danish 1885-1962) a student of Rutherford saw that his model needed improvement. • Bohr proposed that an electron is found only in specific circular paths, orbits, around the nucleus. • Energy levels – the fixed energies an electron can have.

5. Energy levels are like steps or rungs on a ladder. • Quantum – amount of energy required to move an electron from one energy level to another energy level. • Energies between levels are not all the same

6. The Quantum Mechanical Model • Erwin Schrödinger (Austrian 1887-1961) used math to describe the behavior of the electrons. • Quantum mechanical model – modern description of electrons in atoms based on mathematical solutions to the Schrödinger equation.

7. The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus. • Think of as a fuzzy cloud of chance.

8. Atomic orbitals • Atomic orbitals – a region of space in which there is a high probability of finding an electron. • Distinguished by n (principle quantum number or energy level) and a number (n = 1, 2, 3…)

9. In each energy level there are orbitals (shapes) called sublevels. • Each energy sublevel corresponds to a different shape, which describes where the electron is likely to be found.

14. Each energy level has as many sublevels as the level number (ex: level 1 has 1 sublevel, level 2 has 2 sublevel (shapes).

15. To find the maximum number of electrons in an energy level use 2n2.

16. 1s orbital = 1 total orbital 2s orbital 2p orbitals 4 total orbitals 3s orbital 3p orbitals 3d orbitals 9 total orbitals

17. Electron Arrangement in Atoms Section 2

18. Electron Configurations • Electron configuration – way in which electrons are arranged in atoms. • Three rules – the aufbau principle, the Pauli exclusion principle, and Hund’s rule – tell you how to find the electron configuration.

19. Aufbau principle • Aufbau principle – states that electrons occupy the orbitals of lowest energy first. • Orbitals on any sublevel are always the same energy.

21. Pauli exclusion principle • Pauli exclusion principle – an atomic orbital may describe at most two electrons. • When electrons pair they must have opposite “spins” so they don’t repel as much.

22. Hund’s rule • Hund’s rule – electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. • Basically singles in a sublevel until they have to double up.

23. Exceptional electron configurations • Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.

24. Physics and the Quantum Mechanical Model Section 3

25. Light • The quantum mechanical model (QMM) came out of the study of light. • Parts of a wave • Amplitude – the wave height from zero to crest. • Wavelength (λ) – distance between two crests. • Frequency (ν) – number of wave cycles to pass a given point per unit of time. • Hertz (Hz) – SI unit of frequency (can also be expressed as s-1).

26. The product of the frequency and wavelength always equals a constant (c), the speed of light.

27. The wavelength and frequency of light are inversely proportional to each other. • Electromagnetic radiation – light in these forms: radio waves, microwaves, infrared waves, visible light, ultraviolet waves, x-rays, and gamma rays. • Spectrum – light separated by a prism into frequencies (or colors).

29. Atomic Spectra • When atoms absorb energy, electrons move into higher energy levels. • The electrons then lose energy by emitting light when they return to lower energy levels. • Atomic emission spectrum – frequencies of light emitted by an element separate into discrete lines. • These lines are unique to each element. Mercury Nitrogen

30. White light through a prism Helium light through a spectrum

31. An Explanation of atomic spectra • Remember Bohr said electrons can have only specific energies. • Ground state – an electron has its lowest possible energy. • Exciting electrons can move them up to a higher energy level but only if the energies match up (the right quanta).

32. The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron.

33. Quantum mechanics • Photons – light quanta (particles of light) • 1924 Louis DeBroglie (French 1892-1987) proposed a thought: • If light behaves like a particle, can other things behave like waves? • He and others found that yes they can, and we’ve begun to think of the electron differently.

34. Classical mechanics adequately describes the motions of bodies much larger than atoms. • Quantum mechanics describes the motions of subatomic particles and atoms as waves. • Heisenberg uncertainty principle – states that it is impossible to know exactly both the velocity and position of a particle at the same time.