Ch 10: Modern Atomic Theory
10.1 Rutherford’s Atom Our understanding of atoms has changed over 2000 years. Democritus thought of atoms as small pieces of matter than cannot be subdivided. Over 2000 years later, Dalton’s atomic theory described atoms as indivisible spheres of matter that could be combined in specific ratios to make compounds. Thompson discovered that atoms were composed of parts when he discovered the electron and concluded there must be a positive part of the atom. Rutherford confirmed the existence of a small, positively charged nucleus surrounded by electrons in orbits; atoms were mostly empty space.
Rutherford’s model of the atom is still the popular model of the atoms today, though it is very inaccurate. However, his model did not explain how the electrons could remain in stable orbits around the nucleus. They should fall towards the protons in the nucleus atom and merge with it, thereby collapsing the atom
10.2 Electromagnetic Radiation To understand the model of the atom further requires knowledge of light and how it transmits energy. EMR is energy transmitted from one place to another by light. The fundamental unit of light is the photon. There are many forms of EMR: gamma rays, x-rays, UV rays, visible light, infra-red rays, microwaves, radio waves. Human eyes can only see a small portion of this spectrum – the visible light part.
EMR (light) has both particle and wave properties.
The wave properties of EMR include wavelength, frequency, and speed. Wavelength is the distance between two consecutive wave peaks. (Think of water waves as analogous to EMR waves. The frequency of the wave refers to how many wave peaks pass a given point per given time period. The speed of a wave indicates how fast a given peak travels through its medium. For light, that speed is 186,000 miles per second (in a vacuum). The speed of light in a constant.
Different wavelengths of EMR (light) carry different amounts of energy (see EMR spectrum). For example, photons that correspond to red light have less energy than photons that correspond to blue light. That’s because the wavelength of red light is longer than that of blue light. This is why we are not easily damaged by visible light rays (ROY G BIV) but can be injured by light of shorter, and therefore more powerful, wavelengths.
UV light has enough energy to cause damage to our skin. X-Rays are even stronger – they can go through our bodies. Can damage cells and cause cancer. Gamma rays are even stronger and can seriously damage or kill organisms even with brief exposures. But remember, even visible light, infra-red, microwaves and radio waves can cause damage if intense enough. After all, microwaves can be used for cooking foods!
10.3 Emission of energy by Atoms Different atoms give off different colored light when heated. When atoms receive energy from a source they become “excited.” They then release this energy by emitting light of a particular wavelength. The energy of the light emitted is equal to the energy received.
10.4 The Energy levels of Hydrogen Atoms with excess energy are in excited state. An excited atom can release some or all of its energy by emitting a photon (a particle) and moving to a lower energy level.
A sample of hydrogen atoms absorb energy and become excited, then release energy in form of protons to return to ground (unexcited) state
When hydrogen atoms return to lower energy states they release photons of visible light of only certain energies that correspond to the colors below. Scientists observed these lines with a prism but did not know why they appeared
Because only photons of certain energies (colors) are emitted, scientists conclude that energy levels in atoms are quantized (i.e., they can only have certain values). But why were they quantized?
Scientists thought atoms could exist at any energy levels and could emit light of any wavelength (color).
10.5 The Bohr Model of the Atom How is the particular emission spectrum of hydrogen explained? In 1913 Niels Bohr constructed a model of the hydrogen atom that had its electron moving in circular orbits corresponding to allowed energy levels. The electron could jump to a different orbit by absorbing or emitting a photon of light of the correct energy content (fig 10.17). When an electron jumped down to a lower orbit after being energized it gave off a photon of light that corresponded to the observed emission spectrum for hydrogen.
It was the action of electrons moving from higher to lower quantized energy levels in an atom that produced the characteristic emission spectrum colors. Bohr’s model fit the hydrogen atom very well but did not explain the emission spectra of other atoms.
10.6 The Wave-Mechanical Model It replaced the Bohr model of the atom. The Bohr model simply put electrons with higher energy in an orbit with a greater radius (distance from the nucleus). This model did not adequately explain the behavior of atoms other than hydrogen. In the 1920s Louis Victor de Broglie and Erwin Schroedinger proposed that electrons have both wave and particle properties, just like photons. By mathematically treating electrons as waves, rather than just particles, they came up with a model that works for all atoms, not just hydrogen.
Electrons, behaving as waves, do not move in orbits but are located in orbitals. Orbitals simply represent the region of space, as mathematically defined, occupied by an electron around an atom. Electrons do not move like planets around the sun. Orbitals are a probability map of where electrons are likely to be found, not precise locations. As you’ll see in Ch 10.7, and below, orbitals have different shapes that correspond to the different energy states of electrons. Shown are the s, p, d, and f orbitals. The nucleus of the atom is located at the intersection of the axes.