Unit 4. Electrons in Atoms

Transcription

1 Unit 4 Electrons in Atoms

2 When were most of the subatomic particles discovered? Who discovered densely packed nucleus surrounded by fast moving electrons?

3 Rutherford s Model Major development Lacked detail about how electrons occupy space surrounding the nucleus??????????? Scientists in the early twentieth century found Rutherford s nuclear atomic model to be fundamentally incomplete.

4 Why don t the electrons just fall into the positively charged nucleus?

5 Rutherford s nuclear model was lacking It did not begin to account for the differences in chemical behavior among the elements.

6 In the early 1900s, scientists began to unravel the puzzle of chemical behavior. They had observed that certain elements emitted visible light when heated in a flame. What caused these differences?

7 ELECTRONS! Analysis of the emitted light revealed that an element s chemical behavior is related to the arrangement of the electrons in its atoms. When excited electrons drop to lower energy levels they release light!

8 To better understand electrons in atoms. First, we need to understand the nature of light.

9 Light and Quantized Energy Wave Nature of light (Classical Physics) Visible light is a form of energy that exhibits wave-like properties known as electromagnetic radiation. Electromagnetic radiation A series of electromagnetic waves that travel in a vacuum at a speed of 3.0 x 10 8 m/s. Radio waves Microwaves Infrared waves Visible light Ultraviolet rays X-rays Gamma rays

10 Electromagnetic Spectrum

11 ROY G. BIV ROY G. BIV is an acronym that helps us remember the order of the visible light spectrum! Red, orange, yellow, green, blue, indigo, violet. As you approach violet, the frequency (and energy) of the wave increases

12 Light and Quantized Energy What are waves Mechanical waves require a medium to travel (air, water, or rope). Electromagnetic waves no medium Matter waves particles and electrons

13 Light and Quantized Energy Wave Properties Wavelength (λ or lambda) distance between equivalent points 10 9 nm = 1m Amplitude height from origin to crest, involves the intensity of the light

14 Light and Quantized Energy Wave Properties Frequency ( or nu) how many waves pass a given point per second. 1 hertz (Hz) = 1 wave per second 1MHz = Hz Hz 562 waves/second s s

15 Light and Quantized Energy Wave Nature of light All electromagnetic light moves at the speed of m/s and is represented by the symbol, c. The speed of light is the product of the wavelength (λ) and frequency ( ). c

16 Light and Quantized Energy Although the speed of electromagnetic waves are constant, the frequency and the wavelength may vary. As you can see from the equation, wavelength and frequency are inversely related; in other words, as one quantity increases, the other decreases. c

17 Light and Quantized Energy

18 Light and Quantized Energy What is the wavelength of a microwave having a frequency of 3.44 x 10 9 Hz? c s m/s -1 c = m

19 Light and Quantized Energy A helium-neon laser emits light with a wavelength of 633 nm. What is the frequency of this light? c 10 9 nm = 1 m c 1 m 633 nm 9 10 nm m s 7 m c = m/s λ = 633 nm 14 Hz

20 Light and Quantized Energy Particle Nature of Light (Quantum Physics) While considering light as a wave does explain much of its everyday behavior, it fails to adequately describe important aspects of light s interactions with matter. Glowing substances Photoelectric effect

21 Light and Quantized Energy Particle Nature of Light The wave model of light cannot explain why heated objects emit only certain frequencies of light at a given temperature, or why some metals emit electrons when colored light of a specific frequency shines on them.

22 Light and Quantized Energy Particle Nature of Light In 1900, the German physicist Max Planck began searching for an explanation as he studied the light emitted from heated objects. matter can gain or lose energy only in small, specific amounts called quanta. Quantum minimum amount of energy that can be gained or lost by an atom This is light acting like a particle!

23 Light and Quantized Energy Particle Nature of Light Planck found that the energy of a quantum of energy (photon) is directly proportional to the frequency. E h h = J s

24 Light and Quantized Energy What is the energy of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x Hz? E h E = ( J s)( s -1 ) E = J

25 Light and Quantized Energy Particle Nature of Light Photoelectric effect electrons, called photoelectrons, will be emitted from a metal when light above a certain frequency is shined on it.

26 Light and Quantized Energy Einstein's explanation treated light like a particle. Unless the incoming light has a high enough frequency (energy) it can t release the photoelectron.

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28 Light and Quantized Energy Particle Nature of Light Atomic emission spectrum non-continuous spectra emitted by glowing atoms.

29 Energy levels electrons orbit in circles around the nucleus at fixed energy amounts (quantized). Ground State an atoms electrons are at the lowest energy levels The higher the energy level the farther it is from the nucleus. Quantum Theory and the Atom Electrons move around the nucleus in circular orbits. Quantum

30 Quantum Theory and the Atom Building on Planck s and Einstein s concepts of quantized energy (quantized means that only certain values are allowed), Bohr proposed that the hydrogen atom has only certain allowable energy states. Impressively, Bohr s model also correctly predicted the frequencies of the lines in hydrogen s atomic emission spectrum.

31 Quantum Theory and the Atom When an atom gains energy, it is said to be in an excited state. Although a hydrogen atom contains only a single electron, it is capable of having many different excited states.

32 Quantum Theory and the Atom Bohr s model worked well for hydrogen, but Fell apart for every other atom on the periodic table!!! It did, however, point in the right direction!

33 Quantum Theory and the Atom In 1924, French Louis de Broglie proposed an idea that accounted for the fixed energy levels of Bohr s model. If waves could be treated like a particle, could particles be treated like waves?

34 De Broglie knew that if an electron has wavelike motion and is restricted to circular orbits of fixed radius, the electron is allowed only certain possible wavelengths, frequencies, and energies. In other words, it would be quantized just like observed. = =

35 Quantum Theory and the Atom Developing his idea, de Broglie derived an equation for the wavelength (λ) of a particle of mass (m) moving at velocity (ν). h mv Does it work? Experiments show that the smaller the particle, the more important it s wave properties!

36 Quantum Theory and the Atom Step by step, scientists such as Rutherford, Bohr, and de Broglie had been unraveling the mysteries of the atom. However, a conclusion reached by the German theoretical physicist Werner Heisenberg a contemporary of de Broglie, proved to have profound implications for atomic models

37 When Heisenberg was pulled over for speeding Quantum Theory and the Atom Heisenberg Uncertainty Principle You can t precisely know both the position and velocity of a particle at the same time. No, but I Do you know how fast you were going? know where I m at!

38 Quantum Theory and the Atom In 1926, Austrian physicist Erwin Schrödinger furthered the waveparticle theory proposed by de Broglie Schrödinger derived an equation that treated the hydrogen atom s electron as a wave

39 Quantum Theory and the Atom Remarkably, unlike Bohr s model, Schrödinger s new model for the hydrogen atom seemed to apply equally well to atoms elements! With this equation, the modern Quantum Mechanical Model was born.

40 Quantum Theory and the Atom Quantum Mechanical Model nucleus electron cloud 90% probability of finding the electron within this space Electron position and energy are described using energy levels, energy sublevels, orbital shapes, and spin.

41 Quantum Mechanical Model

42 Quantum Theory and the Atom Principal Energy Level (n) Describes distance from the nucleus and general energy. n = 1, 2, 3, 4,. The higher the energy level the greater the average distance from the nucleus. Each energy level contains sublevels The number of sublevels on a level is equal to the energy level (n). 1 st energy level has 1 sublevel 2 nd energy level has 2 sublevels

43 Quantum Theory and the Atom Each sublevel contains orbitals. orbital: a three-dimensional region around the nucleus in which an electron moves and is found 90% of the time. Each orbital can hold up to two electrons. The total number of orbitals on a level = n 2. Each sublevel has a different shape of orbital on the level. These shapes are represented by the symbols s, p, d, or f.

44 Quantum Theory and the Atom s Orbitals Each level has one s shaped (spherical) sublevel Only 1 orientation on sublevel An s sublevel can hold 2 electrons 1s 2s 3s

45 Quantum Theory and the Atom p Orbitals 2 nd energy level and above have a p sublevel 3 orientations on each sublevel p Sublevel can hold up to 6 electrons 2p x 2p y 2p z 3p

46 Quantum Theory and the Atom d orbitals 3 rd energy level and above have a d sublevel 5 orientations on each sublevel d sublevel can hold up to 10 electrons d d 2 2 ( x y ) xy d xz d yz d z 2

47 Quantum Theory and the Atom f Orbitals 4 th energy level and above have a f sublevel 7 orientations on each sublevel f sublevel can hold up to 14 electrons f 2 2 f x( x 3y ) xyz 2 f 2 2 xz y(3x y ) f f 2 3 yz f 2 2 z z( x y ) f

48 Electron Configurations Electron Configuration Rules 1. Aufbau principle Electrons enter orbitals of lowest energy first.

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50 Actual Energy Levels HOW CAN YOU REMEMBER THIS?!?!

51 Aufbau Chart 1s 1s 2 2s 2p 2s 2 2p 6 3s 3p 3d 3s 2 3p 6 3d 10 4s 4p 4d 4f 4s 2 4p 6 4d 10 4f 14 5s 5p 5d 5f 5s 2 5p 6 5d 10 5f 14 6s 6p 6d 6f 6s 2 6p 6 6d 10 6f 14 7s 7p 7d 7f 7s 2 7p 6 7d 10 7f 14

52 Electron Configurations 2. Pauli exclusion principle no 2 electrons in an atom can have the same four quantum numbers (level, sublevel shape, orientation, and spin). Carbon 1s 2s 2p 2p 2p Same level, sublevel, and orientation, but different spin Same level, sublevel, and spin, but different orientation

53 Electron Configurations 3. Hund s rule - When filing a sublevel with multiple orbitals (p, d, or f), each orbital must have one electron before any orbital has a second electron. Carbon 1s 2s 2p 2p 2p The 6 th electron has to go in the 2 nd 2p orbital NOT fill the first 2p!

54 Electron Configurations Orbital Notation shows every occupied orbital in every sublevel with electrons. Arrows are used to show the electrons and the direction of spin ( or ) Carbon 1s 2s 2p 2p 2p

55 Electron Configurations Draw the orbital notation of Cobalt. 1. Cobalt has 27 electrons. 2. Draw about 27/2 = 13.5 or 14 dashes (orbitals). 3. Label each orbital. 4. Add electrons following Aufbau and Hund s rules. Co 1s 2s 2p 2p 2p 3s 3p 3p 3p 4s 3d 3d 3d 3d 3d

56 Electron Configurations Electron Configuration Notation each sublevel with electrons is described with the number of electrons in the sublevel as a superscript. Carbon 1s 2s 2p 2p 2p 1s 2 2s 2 2p 2

57 1 s 2 He p d f

58 Electron Configurations Noble Gas or Shorthand Notation like electron configuration except the inner level electrons are described by writing the last noble gas in brackets and then describing the other electrons. lead 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 2 Xe [Xe]6s 2 4f 14 5d 10 6p 2

59 Electron Configurations Electron Dot Notation The element symbol represents the inner level electrons and dots are used to show the valence(outside) electrons. Space out electrons with no more than 2 to a side X C 1s 2 2s 2 2p 2 C

60 Exceptions to the predicted electron configurations Two elements of the first 40 elements have electron configurations different from what would be normally predicted. Predicted: Cr: [Ar] 4s 2 3d 4 Actual: Cr: [Ar] 4s 1 3d 5 Predicted: Cu: [Ar] 4s 2 3d 9 Actual: Cu: [Ar] 4s 1 3d 10