Ch 4

1. Ch 4 Electron Energies

2. Electromagnetic Spectrum • Electromagnetic radiation is a form of energy that exhibits wave-like behavior as it travels though space. • EM radiation is organized into a spectrum according to wavelength ( ) and frequency (v) of the waves. • The spectrum includes the areas; radio waves, infrared, visible light, ultraviolet, x-rays, and gamma rays.

4. Mathematical Relationships • Wavelength and frequency are inversely proportional. As one increases, the other decreases. • The speed of EM waves is related to wavelength and frequency in the following way. c = v • Because the speed of light is constant, it is possible to conclude that wavelength and frequency are inversely proportional.

5. Photons • All areas on the electromagnetic spectrum carry particles of electromagnetic radiation called photons. • A photon has zero mass but carries a quantum of energy.

6. Quanta • This minimum energy contained by the photon is called quantum energy. • This quantum energy is determined by the frequency of the radiation carried by the photon. E = hv

7. Planck • It was Planck who discovered that all energy comes in these packets of quanta. • He was able to prove photons of quantum energy existed by observing the photoelectric effect.

8. Photoelectric Effect • When light shines on metal, photons in the light can knock electron’s off of the atoms in the metal. • This only occurs if the photon that hits the metal has at least the minimum energy required to knock the electron loose. • Therefore, matter absorbs only whole numbers of photons of EM energy.

10. Jumping Electrons • Photons not only knock electrons lose from an atom, they can also be absorbed or released from an atom. • Photons of energy can be absorbed by electrons in an atom causing the atom to be at an excited state. • Photons can also be released by electrons in an atom causing the atom to be at ground state.

11. Orbital Energy • When current is passes through a gas at low pressure, the atoms within the gas become excited. • As an electron falls from the excited state to the ground state, energy is given off in the form of a photon of radiation. • The energy of the photon is equal to the difference in energy between the two orbits. • Ephoton = E2 - E1

12. Emission/absorption Spectra • The released photon of radiation can be sent through a prism where it becomes separated into its specific frequencies, forming a line-emission spectrum. • The color and position of the light on the emission spectrum relate to the wavelength and frequency of the photon and therefore its quantum energy.

13. Continuous or Line Spectra • Because they originally thought atoms would become excited by any amount of energy added, it was thought this spectrum would be continuous. • Instead the spectrum that was produced had only lines of distinct frequencies. • This indicated that only fixed amounts of energy, quanta, were being released or absorbed as electrons moved between orbits.

15. Quantized Energy Levels • This fixed line spectrum suggested that energy differences between the atom’s energy states were also fixed. • These set energy levels were named orbits. • The energy of the orbits increases with increasing distance from the nucleus.

16. Bohr Atomic Model • Bohr proposed a hydrogen atom model that links the atom’s single electron with its photon line emission spectrum. • Bohr found the wavelength from the radiation’s frequency on the line emission spectrum. • Using the wavelength, he calculated the energies that an electron must have to have at each energy level. • This technique allowed him to model the hydrogen atom correctly, but doesn’t work very well with atoms containing more than one electron.

18. Particle/Wave Duality • It was already known that electrons exhibited particle like qualities. • However, the fact that electrons confined to orbits produce only certain frequencies, they were exhibiting wave like properties as well.

19. De Broglie • De Broglie found more evidence for the wave like properties of electrons • Electrons interact with one another just as waves do. • They diffract/bend as they pass by the edge of an object. • They can interfere with one another, producing areas of constructive and destructive interference.

21. Quantum Theory • Quantum theory mathematically describes the wave properties of very small objects such as electrons. • It has become the leading branch of physics that deals with atomic and subatomic systems

22. Orbitals • Based on Heisenberg’s principle, only the probability of the location of an electron can be determined. • Therefore, Bohr’s theory of neat orbits was thrown out. • Instead, it is now thought that electrons orbit the nucleus in three dimensional regions called orbitals. • The orbital give the probable location of an electron.

23. Heisenberg Uncertainty Principle • Electrons are detected by their interaction with photons. But any attempt to locate an electron with a photon knocks the electron off its course. • As a result there is a basic uncertainty in trying to locate an electron. • Heisenberg’s principle states that it is impossible to determine the position and speed of an electron at the same time. • Although difficult for scientists to accept, it has become one of the fundamental principles of our present understanding of light and matter.

25. Quantum Numbers • Quantum numbers describe the properties of orbits and the electrons within the orbits. • Using these, it is possible to figure out why each orbit contains its specified amount of electrons.

26. Principle Quantum Number • The principle quantum number, n, gives the main energy level occupied by the electron. • Electrons that share the same main energy level are said to be in the same shell. • n=1 1st shell n=2 2nd shell n=3 3rd shell

27. Angular Momentum Quantum Number • Angular momentum quantum number, L, indicates the sublevels in the main shell. • L values are zero and all numbers less than n. • The L values correspond to certain shapes of orbits. • 0 = s-shaped and spherical 1 = p-shaped and dumbbell shaped 2 = d-shaped and cross shaped

28. Magnetic Quantum Number • Magnetic quantum number, m, gives the orientation of an orbital around a nucleus. • s = 1 orientation p = 3 orientations d = 5 orientations • The total number of orbitals within each shell is n2. • Each orientation of an orbital can hold two electrons. • Therefore the total number of electrons per shell is 2n2.

30. Spin Quantum Number • Spin quantum number indicates the spin of the electrons in each orbit • The spin of electrons in the same orbit must be opposites. • The two values of these spins are + ½ and – ½.

31. Electron Configuration Notation • Gives the main energy levels and sublevels of the element. • The number of electrons in each sublevel is also shown in superscript. • You start at 1s and continue filling up until the correct number of electrons are used. • 1s22s22p6……

32. Noble Gas Notation • This shortened version of electron configuration allows noble gas symbols to represent part of the configuration. • The noble gas that occurs before the element on the periodic table is the one used. • Only the notation after that noble gas has to be written.

33. Orbital Notation • In this notation, an orbital is represented by a line with the main level and sublevel written underneath it. • Arrows showing electrons and their spin is written above the line. Each line can only hold two electrons. • It is necessary to write the notation for the level as may times as there is orientations for that level.

34. Three Principles • There are three principles that must be followed when writing electron configurations and orbital notations. • Aufbau principle- an electrons occupies the lowest-energy orbital it can. • Pauli exclusion principle- no two electrons in the same orbit can have the same spin quantum number. • Hund’s rule- orbitals of equal energy must each have one electron before any is allowed to have a second.