1. Electron in Atoms Chapter 5

4. Rutherford’s Atomic Model Discovered dense positive piece at the center of the atom- “nucleus” Electrons would surround the nucleus. Atom is mostly empty space It did not explain the chemical properties of the elements.

5. Rutherford’s Atomic Model Rutherford’s model failed to explain why objects change color when heated. As the temperature of the metals is increased it changes from gray to red to yellow. The observed color change could only occur if the atoms in the iron gave off light in specific amounts of energy.

6. The Bohr Model of the Atom How does the understanding of light apply to the understanding of the electron configurations

7. Wave-particle Nature of Light Light is a form of energy that has a dual nature – behaves like wave and like particle- how do we know this

8. Nature of light – as wave The study of light led to the development of the quantum mechanical model. Light consists of waves and has the following properties : Amplitude: is the wave’s height from zero to the crest Wavelength: is the distance between the crests and is represented by Greek letter lambda ( ) Frequency : is the number of wave cycles to pass a given point per unit of time and is represented by (the Greek letter nu). SI unit - cycles per second or hertz (Hz).

9. Electromagnetic Spectrum C = Speed of light (2.998 X 10 8 m/s) Light consists of electromagnetic radiation. When sunlight passes through a prism, the different wavelengths separate into a spectrum of colors c = Low energy ( = 700 nm) High energy ( = 380 nm) Frequency (s -1 ) 3 x 10 6 3 x 10 12 3 x 10 22 10 2 10 -8 10 -14

10. Examples 1.Calculate the frequency of light with a wavelength of 585 nm. 2.Calculate the wavelength of light with a frequency of 1.89 x 10 18 Hz.

11. The Wave Description of Light Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space. Together, all the forms of electromagnetic radiation form the electromagnetic spectrum.

12. The Nature of Energy The wave nature of light does not explain how an object can glow when its temperature increases. Max Planck explained it by assuming that energy comes in packets called quanta.

13. Max Planck and Planck's Constant: Planck hypothesized that EM radiation was given off by vibrating atoms. He proposed that each atom vibrated at fundamental frequencies and therefore emitted radiation of only certain energies

14. Quantization of Energy An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA.

15. Albert Einstein's - The Photoelectric Effect The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal. A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy.

16. Examples 1.Calculate the energy of a photon of light with a frequency of 7.30 x 10 15 Hz. 4.84 x 10 -18 J 2.Calculate the energy of red light with a wavelength of 720 nm. ( 1nm = 10 -9 m) 2.76 x 10 -19 J 1.Calculate the wavelength of a photon with an energy value of 4.93 x 10 - 19 J. 403 nm (4.03 x 10 -7 m)

17. Electromagnetic Radiation Exhibits Both Wave Properties and Particle Properties A ray of electromagnetic radiation is a stream of many packets of energy, called photons. The more intense the light, the more photons are passing per second.

19. Atomic Emission Spectra White light is made up of all the colors of the visible spectrum and passing it through a prism separates it. By heating an element we can get it to give off colors and passing it through a prism does something different (discrete lines). The wavelengths of the spectral lines are characteristic of the element, and they make up the atomic emission spectrum of the element. No two elements have the same emission spectrum (fingerprints of an element).

20. Atomic Emission Spectra When atoms absorb energy, their electrons move to higher energy levels. These electrons lose energy by emitting light when they return to lower energy levels. 4000 A o 5000 6000 7000 Visible Light Spectrum Na H Ca Hg

21. Explanation of Atomic Spectra The energy level, and where the electron starts from, is called it’s ground state - the lowest energy level. Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be excited.

22. Explanation of Atomic Spectra As the electron falls back to the ground state, it gives the energy back as light (photons). They may fall down in specific steps. Each step has a different energy. The further they fall, more energy is released and the higher the frequency.

23. Flame emission spectrum Photographs of flame tests of burning wooden splints soaked in different salts. methane gaswooden splintstrontium ioncopper ionsodium ioncalcium ion

24. Bohr’s Atomic Model Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus. Each electron orbit has fixed energy called energy levels.

25. Bohr’s Atomic Model An atom with its electron in the lowest energy level is in its ground state. The energy of the electron is an orbit is proportional to its distance from the nucleus. Quantum = the amount of energy needed to move an e - from one energy level to the next. The energy of an e - is therefore said to be quantized.

26. Quantum Mechanical Model Bohr’s model - Not quite right..... The model was only able to explain hydrogen atom. Why??? Electrons do not orbit the nucleus in circular paths.

27. Quantum Mechanical Model In 1926, Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom. e - are not found in specific orbitals. Electrons are in a ‘cloud’. Equation for the probability of a single electron being found along a single axis (x-axis)

28. Quantum Mechanical Model 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 of an atom. Electron cloud

29. Atomic Orbitals and Quantum Numbers Atomic orbitals - regions where there is a high probability of finding an electron The quantum mechanical model uses quantum numbers (n, l, m l ) to describe an orbital.

30. Quantum Numbers 1. Principal Quantum Number (n): the energy level where the e - is found.  n = 1, 2, 3, 4 Maximum number of electrons that can fit in an energy level is: How many e - in level 2? 3? 2n 2

31. 2. Azimuthal Quantum Number ( l ): defines the shape of the orbital. Also known as angular momentum quantum number Energy sublevel – correspond to orbitals of different shapes and describes where e - may be found  Letter symbols: s, p, f, d  s orbital: e - are found in spherical shapes. No sharp edges  p orbital: Dumbbell shape and there are 3 different shapes.  d orbital: 5 different shapes  f orbital: 7 different shapes

32. Quantum Numbers 3. Magnetic Quantum Number ( m l ): describes the orientation of the orbital in space. 4. Spin Magnetic Quantum Numbers (m S ): describes the electron spin.  m S = +1/2 or - 1/2

33. Energ y Level Sublevels# Orbitals Total Orbitals Total Electron s 1 2 3 4

34. Energ y Level Sublevels# Orbitals Total Orbitals Total Electron s 1s112 2s148 p3 3s1918 p3 d5 4s11632 p3 d5 f7

35. Electron Configuration Electron configuration is the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how: 1. Aufbau Principle According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram, each box represents an atomic orbital. Increasing energy 6s6s 5s5s 4s4s 3s3s 2s2s 1s1s 6p6p 5p5p 5d5d 4p4p 4d4d 4f4f 3p3p 3d3d 2p2p

36. Electron Configuration 2. Hund’s Rule: every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

37. 3. Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.  To show the different direction of spin, a pair in the same orbital is written as: Each electron in an atom has a unique set of 4 quantum numbers which describe it. i.Principal quantum number ii.Angular momentum quantum number iii.Magnetic quantum number iv.Spin quantum number

38. Orbital Diagrams Each box represents one orbital. Half-arrows represent the electrons. The direction of the arrow represents the spin of the electron.

39. Electron Configuration A convenient shorthand method for showing the electron configuration of an atom involves writing the energy level and the symbol for every sublevel occupied by an electron. Indicate the number of electrons occupying that sublevel with a superscript  Hydrogen (H) : 1s 1  Oxygen (O) : 1s 2 2s 2 2p 4  Phosphorus (P) : 1s 2 2s 2 2p 6 3s 2 3p 3

41. Condensed Electron Configurations (Noble gas notation) Neon completes the 2p subshell. Sodium marks the beginning of a new row. So, we write the condensed electron configuration for sodium as Na: [Ne] 3s 1 [Ne] represents the electron configuration of neon. Core electrons: electrons in [Noble Gas]. Valence electrons: electrons outside of [Noble Gas].

42. The periodic table can be used as a guide for electron configurations. The period number is the value of n. Groups 1 and 2 have the s-orbital filled. Groups 13 - 18 have the p-orbitals filled. Groups 3 - 12 have the d-orbitals filled. The lanthanides and actinides have the f-orbital filled. Electron Configurations and the Periodic Table