Electrons in Atoms Chemistry

Electrons in Atoms Objectives 1. Compare the wave and particle models of light. 2. Define a quantum of energy and explain how it is related to an energy change in matter. 3. Contrast continuous electromagnetic spectra and atomic emission spectra. 4. Compare the Bohr and quantum mechanical models of the atom. 5. Explain the impact of de Broglie’s wave particle duality and the Heisenburg uncertainty principle on the modern view of electrons in atoms.

Light and Quantized Energy Even though Rutherford’s nuclear model was a major discovery, it did not explain how electrons actually behaved. -did not: a. state how electrons occupy the space around the nucleus b. account for the differences in chemical behavior of the various elements This chapter describes how electrons are arranged and how it plays a role in chemical behavior.

Light and Quantized Energy By the early 1900’s, the puzzle of chemical behavior was beginning to be understood. -observed certain elements emit a visible light when heated To understand this relationship, you must understand the way light behaves. electromagnetic radiation: form of energy that exhibits wavelike behavior as it travels through space -visible light, microwaves, x rays, radio waves, are a few examples

Light and Quantized Energy Waves are characterized by: -wavelength (l): shortest distance between equivalent points on a continuous wave, usually crest to crest or trough to trough ♦measured in m, cm, or nm -frequency (n): number of waves that pass a given point per second ♦measured in hertz (Hz), 1/s or s-1 -amplitude: waves height from the origin to a crest or from origin to the trough

-speed of light (c): all electromagnetic waves travel at a speed of 3.0 x 108 m/s in a vacuum

Light and Quantized Energy Speed is represented by: - c = l n ♦ wavelength & frequency are inversely proportional

Light is actually a continuous range of wavelengths and frequencies. -can be separated into colors by a prism, which is called the visible spectrum, though the visible spectrum makes up a tiny portion of the complete electromagnetic spectrum Pictures came from http://www2.chemistry.msu.edu/faculty/reusch/VirtTxtJml/Spectrpy/UV-Vis/spectrum.htm

electromagnetic spectrum: includes all forms of electromagnetic radiation, with the only differences in the types of radiation being their wavelengths and frequencies. -p 120, Fig 5-5 Pictures came from http://www.lbl.gov/MicroWorlds/ALSTool/EMSpec/EMSpec2.html

Particle Nature of Light Though light behaves as a wave, it does not adequately describe important aspects of light’s interactions with matter. -why heated objects emit only certain frequencies at given temperatures

Max Planck began searching for an explanation for these phenomena. -he concluded that matter can gain or lose energy only in small, specific amounts called quanta, the minimum amount of energy that can be lost or gained by an atom -proposed the emitted light was quantized and showed a mathematical relationship between quanta and frequency Equantum = hn where h = Planck’s constant = 6.626 x 10-34 J and Equantum and n are proportional

Photoelectric Effect the wave model of light could not explain a phenomenon called the photoelectric effect. In the photoelectric effect, electrons (called photoelectrons) are emitted from a metal’s surface when light of a certain frequency shines on the surface.

While a beam of light has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or bundles of energy, called photons -photon: a particle of electromagnetic radiation with no mass that carries a quantum of energy. Extending Planck’s idea of quantized energy, Einstein calculated that a photon’s energy depends on its frequency. For the photoelectric effect to occur, a photon must possess, at a minimum, the energy required to free an electron from an atom of the metal.

Atomic Emission Spectra atomic emission spectrum: set of frequencies of the electromagnetic waves emitted by atoms of that element -not a continuous range, but several distinct lines of color -unique to each element (fingerprint) -can be used to identify that element

Quantum Theory of the Atom Neils Bohr proposed a quantum model to explain why hydrogen’s atomic emission spectrum was discontinuous. -hydrogen has only certain allowable energy states ♦ground state: lowest allowable energy level ♦excited state: when an atom gains energy and goes into a higher energy state than the ground state -the larger the orbital, the higher the energy state ♦he assigned a quantum number, n, to each orbit (n = 1 is the ground state, or first orbit) ♦there are 7 orbits, or energy levels

Electrons can only circle the nucleus only in allowed paths, called orbits, though they can jump from one orbit to another, if 1) space allows and 2) there is enough energy for the electron to make the jump

Bohr stated that in the ground state, an atom does not radiate energy, though it will in a higher energy state. -as an excited atom falls back to its ground state (lower energy state), it emits a photon of radiation - Ephoton = E2 - E1 = hn where E2 = higher energy state and E1 = lower energy state -these photons emitted occur only at certain frequencies, giving off distinct colors

This shows the relative energies of the electron transitions responsible for hydrogen’s four spectral lines. -called the Balmer series (visible series)

Though Bohr’s model explained hydrogen’s spectral lines well, it did not -explain the spectrum of any other element -fully account for the chemical behavior of atoms Why was his theory so important? It laid the groundwork for later atomic models

Quantum Mechanical Model Scientists (mid 1920’s) were convinced that Bohr’s model was incorrect. -Louis de Broglie proposed an idea to explain the fixed energy levels ♦Bohr’s quantized orbits have wave-like characteristics ♦light has both wave and particle characteristics: ‘If waves have particle-like behavior, can particles of matter (including electrons) behave like waves?’

de Broglie knew that if an electron has wave-like motion and is restricted to circular orbits, the electron is allowed only certain wavelengths, frequencies and energies -developed a wave equation, the de Broglie equation, which predicts all moving particles have wave characteristics l = h where v = velocity (speed) mv and m = mass We now have begun making great strides in understanding the atom.

Heisenburg Uncertainty Principle Werner Heisenburg used the work of Rutherford, Bohr and de Broglie to make his own conclusion, which had profound implications for the atomic model. -Heisenburg Uncertainty Principle: states it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time. -even though scientists at the time had trouble understanding this concept, it was able to describe the fundamental limitations on what can be observed

Schrodinger Wave Equation Erwin Schrodinger used de Broglie’s ideas and the Heisenburg Uncertainty Princple to further the wave-particle theory -derived an equation treating hydrogen’s electron as a wave -worked with all atoms, not just hydrogen’s -electrons do not travel around the nucleus in neat orbits, but exist in regions called atomic orbitals, a three dimensional region around the nucleus that indicates the probable location of an electron We now have a working quantum mechanical model of the atom.

Quantum Mechanical Model Did: limits an electron’s energy to certain values Did Not: makes no attempt to describe the electron’s path around the nucleus.

Intro to Spectroscopy Lab A white light source can be separated into the individual wavelengths in the visible region of the spectrum (400 nm to 650 nm.) This separation of the different wavelengths of visible light can be done with a diffraction grating (which works like a prism.) The diffraction grating can be adjusted so that only the energy of a specific wavelength of light is passed through a sample, and a detector can then determine how much of that energy was absorbed by the sample.

Intro to Spectroscopy Lab This is how the Spectronic 20 works. Refer to the diagram below to follow the path of the light through the Spectronic 20. The diffraction grating used to separate the wavelengths is turned when you turn the knob on top of the instrument. In doing so, you determine the specific wavelength of light that passes through your sample, and the Spectronic 20 gives you a readout of the light that was absorbed by your sample. diffraction grating light source detector cuvette

Intro to Spectroscopy Lab Purpose This lab will familiarize you with the operation of the Spectronic 20. At the conclusion of the lab, you should understand the function and know the names of each of the controls on the spectrophotometer. You should also be able to set the spectrophotometer to a given wavelength and zero it using a blank. Hypothesis I will demonstrate the proper way to use a spectrophotometer. Materials/Equipment spectrophotometer chalk distilled water 2 cuvettes Kimwipes

Terminology: 1. A cuvette looks like a test tube. It fits inside the spectrophotometer to hold your sample. It is different from a test tube because the glass it is made of is more uniform in width than a test tube. This is important because the light must pass through both walls of the cuvette before it reaches the detector, and the thickness of the glass should not affect the readings you get from the spectrophotometer. 2. Kimwipes are small tissues like Kleenex. They are used to wipe fingerprints and dirt off the cuvette before inserting it into the spectrophotometer. Safety Considerations: 1. Always wear goggles and an apron in the lab. 2. Please be careful with the cuvettes; they are expensive.