Chapter 4 arrangement of electrons in atoms

1. Chapter 4 arrangement of electrons in atoms Electromagnetic radiation – form of energy that exhibits wavelike behavior as it travels through space. Electromagnetic spectrum – all wavelengths of light Speed of light = 3.00 x 108 m/s = c Wavelength = λ (lambda) distance between corresponding points on adjacent waves. Frequencyυ (upsilon)= number of waves that pass a given point in a specific time. c = λυ

3. The Photoelectric Effect • Emission of electrons from a metal when light shines on the metal. • Involves the frequency of the light striking the metal

4. Particle Description of Light • Max Planck • Object emits energy in small, specific packets called quanta not in a continuous amount. • Quantum – minimum quantity of energy that can be lost or gained by an atom. E = hυh = 6.626 x 10-34 Jxs Planck’s constant Unit of Energy = Joules

5. Photon – particle of electromagnetic radiation having zero mass and carrying a quantum of energy. Ephoton= hυ Different metals require different minimum frequencies to exhibit the photoelectric effect.

6. The Hydrogen-atom Line-emission Spectrum When a current is passed through a gas at low pressure, the potential energy of the gas atoms increases (they become excited). Ground state – lowest energy state of an atom Excited state – state when atoms have a higher potential energy than it has in its ground state.

7. Hydrogen Gave off an emission-line spectrum – just a certain range of light. Scientists concluded that hydrogen’s electron contains a fixed amount of energy.

8. Bohr Model Energy level 1 – 7, electrons have more energy at higher energy levels. When e- move to a higher energy level it absorbs energy. When e-move to a lower energy level e- give off energy, as light. E = hv

9. Sec.2 The Quantum Model of the Atom De Broglie hypothesized that electrons have wave-like properties and made up of a particle (wave-particle nature). Just like light.

10. Heisenberg Uncertainty Principle Photons have about the same energy as electrons. States that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.

11. 4th shell – period 4 5th shell – period 5 6th shell – period 6

12. Schrodinger Wave Equation Quantumtheory – wave properties of electrons and other very small particles. Equation gives a probability of finding an electron at a given place around the nucleus. e- exist in certain regions called orbitals. Orbital – 3d region around the nucleus that indicates the probable location of an e-.

13. Atomic Orbitals and Quantum Numbers Bohrmodel – e- of increasing energy occupy orbits farther and farther from the nucleus. Schrodingerequation – e- in atomic orbitals also have quantized energies. Quantumnumbers – specify the properties of atomic orbitals and the properties of electrons in orbitals. Schrodinger equations shows us the energy level, shape, orientation(which orbital), then spin.

14. Principal Quantum Number Symbolized by n, indicates the main energy level occupied by the e-. N are positive integers As n increases, the e-’s energy and its average distance from the nucleus increases. Angular momentum quantum number –indicates the shape of the orbital.

15. Aufbau principle – an electron occupies the lowest-energy orbital that can receive it. Lowest = 1s Pauli exclusion principle – no two electrons in the same atom can have the same set of four quantum numbers. Ex. Angular momentum, energy, shape, orientation of an orbital Hund’s rule – orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, e- in singly occupied orbitals must have the same spin state.