•Atomic Structure

• This electron microscope high-resolution image shows magnification of the thin edges of a piece of mica. The white dots are "empty tunnels" between layers of silicon-oxygen tetrahedrons, and the black dots are potassium atoms that bond the tetrahedrons together. Note the 10 Angstrom width, which is 0.000001 mm.

• First Definition of the Atom

• The Early Greeks thought of everything as being made up of four basic elements.– Earth.– Air.– Fire.– Water.

• Democritus and Leucippus thought that matter was discontinuous or made up of individual particles.– Democritus called these fundamental particles atoms.

• Atomic Structure Discovered

• Introduction– In 1661 Robert Boyle defined an element as a simple

substance which could not be broken down into simpler substances.

– We now define an element as a pure substance that can not be broken down into simpler things by either chemical or physical methods.

– Since elements always combine in fixed ratios, this lends support to the idea of elements being made of discrete particles.

• (A) Oxygen and lead combine to form yellow lead oxide in a ratio of 1:13. (B) If 1 atom of oxygen combines with 1 atom of lead, the fixed ratio in which oxygen and lead combine must mean that 1 atom of lead is 13 times more massive than 1 atom of oxygen.

• Reasoning the existence of atoms from the way elements combine in fixed- weight rations. (A) If matter were a continuous, infinitely divisible material, there would be no reason for one amount to go with another amount. (B) If matter is made up of discontinuous, discrete units (atoms), then the units would combine in a fixed-weight ratio. Since discrete units combine in a fixed-piece ratio, they must also combine in a fixed-weight-based ratio.

• Discovery of the Electron– A cathode ray is a beam of electrons that moves between

metal plates in an evacuated tube from a negative to a positive terminal.

• The electron beam is seen as a green beam.

– These rays can be deflected by a magnet.

• A vacuum tube with metal plates attached to a high voltage source produces a greenish beam called cathode rays. These rays move from the cathode (negative charge) to the anode (positive charge).

– In 1897 JJ Thompson place a positively charges plate on one side of the tube and a negatively charged plate on the other side of the tube.

• The beam was deflected away from the negative plate toward the positive plate.

• Thompson realized that the particles that made up the beam must be negatively charged, since like charges repel and opposite charges attract.

• By balancing the deflections made by the magnet with that made by the electrical field, Thompson was able to calculate the ratio of the charge to mass of an electron as 1.7584 X 1011 coulomb/kilogram

• These particles were later named electrons.

• What appears to be visible light coming through the slit in this vacuum tube is produced by cathode ray particles striking a detecting screen. You know it is not light, however, since the beam can be pulled or pushed away by a magnet and since it is attracted to a positively charged metal plate. These are not the properties of light, so cathode rays must be something other than light.

• A cathode ray passed between two charged plates is deflected toward the positively charged plate.

– The ray is also deflected by a magnetic field. – By measuring the deflection by both, J.J. Thomson was

able to calculate the ratio of charge to mass. – He was able to measure the deflection because the detecting

screen was coated with zinc sulfide, a substance that produces a visible light when struck by a charged particle.

– In 1906 Robert Millikan passed mineral oil through a vaporized into an apparatus where he could observe the drops with a magnifier and make measurement on them as they drifted downward.

• he found that the least charge on any of the droplets was 1.60 X 10-19 coulombs and that larger droplets always had a charge that was some multiple of this value.

• Knowing Thompson’s work of charge to mass ratio and the charge on an individual electron, it was possible to calculate the mass of the electron as 9.11 X 10-31 kg.

• Thompson proposed that an atom was a blob of positively charged matter in which electrons were stuck like raisins in plum pudding.

• Millikan measured the charge of an electron by balancing the pull of gravity on oil droplets with an upward electrical force. – Knowing the charge-to-mass ratio that Thomson had

calculated, Millikan was able to calculate the charge on each droplet.

– He found that all droplets had a charge of 1.60 x 10-19 coulombs or multiples of that charge.

– The conclusion was that this had to be the charge of an electron

• The Nucleus– Ernst Rutherford determined that there was a positively

charge nucleus associated with the atom, that was surrounded by electrons.

• Rutherford calculated that the radius of the nucleus to be about 10-13 cm and the radius of the atom to be about 10-8 cm.

• Electrons therefore took up about 100,000 times the radius of the nucleus.

• Rutherford and his co-workers studied alpha particle scattering from a thin metal foil. – The alpha particles struck the detecting screen, producing

a flash of visible light. – Measurements of the angles between the flashes, the

metal foil, and the source of the alpha particles showed that the particles were scattered in all directions, including straight back toward the source

– In 1917 Rutherford broke up the nucleus of the nitrogen atom by bombarding it with alpha particles and was able to identify a particle with a positive charge called a proton.

• He also thought that there were neutral particles in the nucleus called neutrons.

• The atom has a tiny, massive nucleus made up of protons and neutrons.

• Negatively charged electrons, whose charge balances the charge on the protons, move around the nucleus at a distance of about 100,000 times the radius of the nucleus.

• Rutherfords's nuclear model of the atom explained the alpha scattering results as positive alpha particles experiencing a repulsive forced from the positive nucleus

– Measurements of the percent of alpha particles passing straight through and of the various angles of scattering of those coming close to the nuclei gave Rutherford a means of estimating the size of the nucleus.

• From measurements of alpha particle scattering, Rutherford estimated the radius of an atom to be 100,000 times greater than the radius of the nucleus. This ratio is comparable to that of the (A) thickness of a dime to the (B) length of football field.

• The Bohr Model

• The Quantum Concept.– In 1900 Max Plank introduced the idea that matter emits

and absorbs energy in discrete units called quanta.– In 1905 Albert Einstein extended the quantum concept

to include light and that light consist of discrete units called photons.

– The energy of a photon is directly proportional to the frequency of vibration.

• E=hf– where E = energy

– h = Plank’s constant = 6.63 X 10-34 J•s

– f = frequency

• (A) Light from incandescent solids, liquids, or dense gases, produces a continuous spectrum as atoms interact to emit all frequencies of visible light (B) Light from an incandescent gas produces a line spectrum as atom emit certain frequencies that are characteristic of each element.

• Atomic hydrogen produces a series of characteristic line spectra in the ultraviolet, visible, and infrared parts of the total spectrum. The visible light spectra always consist of two violet lines, a blue-green line, and a bright red one.

• Bohr’s Theory– Allowed Orbitals

• An electron can only orbit around an atom in specific orbits

– Radiationless Orbits• An electron in an allowed orbit does not emit radiant energy as

long as it remains in the orbit.

– Quantum Leaps• An electron gains or loses energy only by moving from one

allowed orbit to another.• The lowest energy state is known as the ground state• Higher states are known as excited states

• Each time an electron males a "quantum leap," moving from a higher energy orbit to a lower energy orbit, it emits a photon of a specific frequency and energy value.

• An energy level diagram for a hydrogen atom, not drawn to scale. The energy levels (n) are listed on the left side, followed by the energies of each level in J and eV. The color and frequency of the visible light photons emitted are listed on the right side, with the arrow showing the orbit moved from and to.

• These fluorescent lights emit light as electrons of mercury atoms inside the tube gain energy from the electric current. As soon as they can, the electrons drop back to their lower-energy orbit, emitting photons with ultraviolet frequencies. Ultraviolet radiation strikes the fluorescent chemical coating inside the tube, stimulating the emission of visible light.

• Quantum Mechanics

• Quantum mechanics states that light and matter, including electrons, have a dual nature of both particles and waves.

• Matter Waves.– Louis de Broglie reasoned that particles must also have a

dual nature.– He reasoned that the electron should have a certain

wavelength that would fit into its orbit around the nucleus.

∀λ=h/mv

• where λ is the wavelength• h is Plank’s constant• m is the mass• v is the velocity

• (A) A schematic of de Broglie wave, where the standing wave pattern will just fit in the circumference of an orbit. This is an allowed orbit. (B) This orbit does not have a circumference that will match a whole number of wavelengths; it is not an allowed orbit.

• Wave Mechanics– Electrons do emit light in certain wavelengths based on

their energy levels (orbital radius)– Since waves spread out from the electron, the wave

mechanic model predicts an area where an electron would be found, and not a specific place where it would be found.

• The Quantum Mechanics Model– Quantum mechanics describes the energy levels of an

electron wave with four quantum numbers.• distance from nucleus• energy sublevel• orientation in space.

• direction of spin

– Principal quantum number (n)• Describes main energy level of the electron in terms of its

distance from the nucleus.• n = 1, 2, 3, 4, 5, 6, 7

– Angular momentum quantum number• Defines energy sublevels within the main energy

levels

• s, p, d, or f designating the type of orbital and also the orbital shape.

• The Heisenberg Uncertainty Principle states that you cannot measure the momentum and exact position of an electron at the same time.

– What you can measure is the probability that an electron will be found in a certain area, called an orbital.

• (A)An electron distribution sketch representing probability regions where an electron is most likely to be found. (B) A boundary surface, or contour, that encloses about 90 percent of the electron distribution shown in (A). This three-dimensional space around the nucleus, where there is the greatest probability of finding an electron, is called an orbital.

– Magnetic quantum number• Defines the orientation in space of the orbitals relative

to an electrical field.• The s orbital has one orientation• The p sublevel can have 3 orientations• The d sublevel can have 5 orientations

• The f sublevel can have 7 orientations.

• (A) A contour representation of an s orbital. (B) A contour representation of a p orbital.

– Spin quantum number• Describes the direction of spin of an electron in its

orbit.• Electrons occur in pairs and each of the orientations

for a sublevel can have one electron pair.

• Experimental evidence supports the concept that electrons can be considered to spin one way or the other as they move about an orbital under an external magnetic field.

– Pauli Exclusion Principle• No two electrons can have the same set of quantum numbers.• At least one of the quantum numbers must differ.

• Electron Configuration– This is a shorthand designation for electron orientation.– The lowest possible energy level is n=1.

• If one electron already occupies this energy level, a second can only occupy it if it has a different spin quantum number.

– Electron configurations tells you the quantum numbers of the electron.

– Energy sublevel ↓

Principle quantum number 1s2 two electrons

• There are three possible orientations of the p orbital, and these are called px, py, and pz. Each orbital can hold two electrons, so a total of six electrons are possible in the three orientations; thus the notation p6.

• A matrix showing the order in which the orbitals are filled. Start at the top left, then move from the head of each arrow to the tail of the one immediately below it. This sequence moves from the lowest-energy level to the next higher level for each orbital.