Physical Science

Physical Science Atomic Physics Slides subject to change

Modern Physics Focus on the atomic, microscopic scale. Properties of electrons, and the electron cloud around each nucleus. Properties of the nucleus, subatomic particles. Nature of light.

Also Quantization • At the level of atoms, many quantities are found only in integer multiples of elementary amounts (for example, electric charge e). • These microscopic quantities are said to be quantized − it’s as tiny as you can get, cannot be subdivided.

Photons • In 1905 Einstein proposed that light exists as tiny packets, or particles, which he called photons. Light is quantized. • “Can be absorbed or generated only as a whole." • We live in a quantum universe, one built out of tiny, discrete chunks of energy and matter.

Light Energy • A photon has no mass, yet it has energy. • Energy of one photon equals Planck’s constant h times the frequency f of the light wave. E = hf • Planck’s constant” h = 6.626x10–34 J-s

Green Laser Energy of one photon from the green laser: f = 5.64x1014 Hz Ephoton= hf = (6.626x10–34)(5.64x1014) = 3.74x10–19 J

Photons Interact with Atoms • When light is absorbed by an object, the absorption occurs in the atoms of the object. Emission of light comes from the atoms. Emission Absorption

Things That Are Invisible Hebrews 11:3 By faith we understand that the universe was formed at God's command, so that what is seen was not made out of what was visible.

Dalton Model 1807 An atom is ... tiny indivisible uniformly dense solid sphere Each chemical element is composed of tiny, indivisible particles called atoms, which are identical for that element. Billiard Ball Model

J.J. Thompson Model Electron discovered in 1897 by English physicist J.J. Thompson. But atoms are electrically neutral, so he proposed the negative charges are spread throughout the atom (1903). Perhaps electrons are like raisins in Christmas “plum pudding?” Thompson’s “Plum Pudding” Model

Ernest Rutherford Model Ernest Rutherford (New Zealand chemist) discovers atoms have small charged nucleus (1911). His students Ernest Marden and Hans Geiger aim alpha particles (helium nuclei) at thin gold foil. 1 out of every 8,000 particles were reflected.

Rutherford’s Model All the positive charge is crammed inside a tiny, massive nucleus. Nucleus is equivalent to a marble in the middle of a football stadium. The much lighter electrons are well outside the nucleus. Atoms are almostly entirely empty space. Rutherford’s “Planetary” Model

Nuclear Dimensions

Hydrogen Emission • When light is emitted from a pure gas (say, in a gas discharge tube) there are a limited number of frequencies in the line spectrum. Why?

Bohr Model • Niels Bohr model (1913) looks like a planetary model. • Orbit radius, and total energy of each electron orbit has a quantum value. • Result: electrons only orbit at certain radiuses. • It has lowest energy when it is closest to the nucleus, higher energy farther from the nucleus.

Bohr Model for Hydrogen • An electron’s energy in its orbit is characterized by a whole number value, principal quantum numbern = 1, 2, 3 ... • When hydrogen electron is in the lowest state (closest to the nucleus), n = 1, it’s in the ground state.

Bohr Model • Hydrogen emits a photon when the electron drops from a higher energy level to a lower energy level. Emission

Bohr Model Hydrogen absorbs a photon when E=hf is exactly the right energy to enable a jump from a lower energy level to a higher energy level. Absorption

Bohr Characteristics • Bohr hydrogen orbit radii • r1 = “Ground State” radius = 0.053x10−9 m = 0.053 nanometer. • r2 = 4 r1 • r3 = 9 r1 • rn = n2 r1

Energy in Orbit Bohr hydrogen orbit energy E1 = “Ground State” energy = − 13.6 electron volts (approximately 2.18x10−18 joules. E2 = E1/4 = −3.4 eV E3 = E1/9 = −1.5 eV En = E1/n2

Hydrogen • Sometimes a hydrogen electron’s energy is drawn like this: When atom absorbs energy, electron goes to a higher energy level. –1.5 E3 Electron volts (eV) –3.4 E2 E1 –13.6 Ground State at this energy level 1 electron volt (eV) is 1.6x10–19 J

Hydrogen • Electron needs energy to jump to higher state. May come from heat (bumped by another atom), light (photons), or electricity (bumped by electrons). • Enables it to jump from E1 to say, E2 or higher. • Energy must be exact amount.

Hydrogen • Electrons give up energy to jump to a lower state. • May jump from E2 down to E1. • Will give up exact amount with a photon of E = hf. When atom emits energy, electron goes to lower level. –1.5 E3 Sometimes a hydrogen electron’s energy is drawn like this –3.4 E2 Electron volts (eV) E1 –13.6 Ground State at this energy level 1 electron volt (eV) is 1.6x10–19 J

Energy Transitions • If an H atom is in excited state E2, it will normally spontaneously emit a photon to go from E2→ E1. • ΔE = E1 – E2 = (− 13.6) − (−3.4) = –10.2 eV ·1.6x10−19 J/1 eV = –1.63x10–18 J • The minus sign indicates the atom lost energy (the energy goes into a photon).

Energy Transitions • From Einstein and Bohr, loss of energy goes into a emission of a photon of specific frequency. • E = hf = 1.63x10–18 J • f = 1.63x10–18 / 6.634x10–34 • f2→1= 2.46x1015 Hz • This is a photon in the ultraviolet range.

Energy States • The characteristic red line in the hydrogen spectrum represents an electron transition from the n = 3 state to n = 2 state.

More on Light • In 1905 Einstein proved that energy and matter are linked in the most famous relationship in physics: E = mc2. • The energy content of a body is equal to the mass of the body times the speed of light squared. • He suggested that the heat produced by radium could mark the conversion of tiny amounts of the mass of the radium salts into energy.

Even More on Light • Mass can be converted into energy and energy into mass. • Example, there are two particles with the same mass, an electron and a positron (an “anti-particle”). • Their mass is each 9.11x10−31 kg. • Total equivalent energy is 2mec2 = 1.64x10−13 J. • Equal but opposite charge −e, +e.

Even More on Light • What is a photon of that energy? • E = hf • 1.64x10−13 = (6.626x10−34) f • f = 2.47x1020 Hz, a gamma ray. • An electron-positron collision conserves total charge, and annilates both particles and creates a gamma ray.

Historical Perspectives Up to 1900, “Classical Physics” Laws of motion Galileo, Newton Gravitation Newton Electricity Coulomb Magnetism, Faraday, Orsted Electromagnetic waves Maxwell Optics Post-1900, “Modern Physics”

Lasers • Light Amplification by Stimulated Emission of Radiation (LASER). • Some atoms can stay in certain “metastable” energy states for a long period of time. • Sending a photon through the material will “stimulate” transition from metastable state to a lower state • Creates an avalanche of photons that bounce off mirrors at each end. • Keep pumping energy into lasing medium.

Heisenberg Uncertainty • At the microscopic level, it is impossible to know a particle’s exact position and velocity simultaneously. • There is a limit on measurement precision at microscopic levels. The measurement process itself disturbs what you are measuring in this environment.

Absorption and Emission • When visible light of all wavelengths pass through a cool sample of the gas, we get an absorption spectrum. • High voltage across a sample, get an emission spectrum. Absorption Emission

Electron Cloud • Erwin Schrödinger model (1926). • The electron’s position is is only a probability. Hydrogen Ground state, most probable radius ~0.053 nm Exactly the same as Bohr’s model.

Matter Waves If light can be a wave and act like a particle, can particles act like waves? Louis deBroglie proposed: λparticle = h/mv Example: Wavelength of small ball, 0.5 kg at 26 m/s λparticle = h/mv = (6.634x10−34)/(0.5)(26) = 5.1x10−35 m Too short to see any effect.

Electron Diffraction Electron beam through a crystal does show diffraction effects.

Light, As Bullets • You expect light to hit the target as shown. • However, … Double Slit

Split Personality • The particles form a diffraction pattern. CCD camera detects single photons Double Slit Diffraction pattern

Physical Science