Electron Configuration

Electron Configuration Chemistry

Chapter 4 – Section 1 • The Development of a New Atomic Model

Rutherford’s Model • A. He shot positive particles (alpha particles) at a piece of thin gold foil • B. He thought the mass and charge of the atom was uniform throughout the atom (Thomson’s model – plum pudding) • C. He expected the particles to go straight through with only slight deflections (for the most part, this did occur) • D. He was shocked to discover that 1 in about 8000 particles actually deflected backwards (as if hitting something solid)

E. He compared this backwards deflection to shooting a cannonball at a piece of tissue paper and having it bounce back and hit you

F. He reasoned that atoms had a small densely packed, positively charged center of matter (nucleus) • 1.If a marble was a nucleus of an atom • 2.Then the entire size of the atom would be a football field

G. Limitations to his model were that it did not describe how the electrons were distributed around the nucleus • 1.Later studies revealed a relationship between light, energy, matter, and atomic structure • 2.Before 1900 scientists thought of light as a wave only • 3.Then, it was discovered that light has characteristics of particles, also

II. Wave Properties of light • A. Visible light is a type of electromagnetic radiation, or a form of energy that acts like a wave as it goes through space • B. Other forms of electromagnetic radiation include X-rays, ultraviolet rays, infrared rays, microwaves, radio waves, and others, which all together make the electromagnetic spectrum • C. All electromagnetic radiation travels at the speed of 3.00 x 108 m/s through a vacuum (this is also known as the speed of light)(this is the approximate speed through air, since it is mostly space)

D. A wave is a rhythmic disturbance that transfers energy (rhythmic means it is repetitive) • 1.Two characteristics of a wave are its wavelength and its frequency • 2.Wavelength (λ) is the distance between two identical points on two adjacent waves (meters) • 3.Frequency (v) is the number of waves that pass a certain point in a certain period of time (usually a second) [waves/second or Hertz (Hz)]

E.. Wavelength and frequency are inversely related (when one increases, the other decreases) • 1.The mathematical relationship is c = λv • 2.c (speed of light in m/s) • 3.λ (wavelength in m) • 4.V (frequency in s-1) • 5.Since c is the same for all electromagnetic waves, λv is always a constant (therefore, if one increases, the other must decrease proportionally)

III. The Photoelectric Effect • A. In the early 1900’s two experiments showed properties of light that could not be explained by the wave theory • B. Photoelectric Effect - electrons are given off from metal when light shines on it • C. For any metal, no electrons were given off if the frequency of light was low enough, no matter how bright the light was (brightness refers to the amount of energy in the light) • D.Before this scientists believed it was the amount of energy that knocked the electrons off metal

E. Max Planck studied light given off by hot objects, and discovered that energy is not given off continuously as in a wave • 1.Instead, it is given off in small packets, which he called quanta • 2.Quantum – the smallest possible amount of energy that can be lost by an atom • F. Planck’s mathematical relationship between quantum and energy is E = hv • 1.E - energy of a quantum (Joules) • 2.v – frequency (s-1) • 3.h – Planck’s constant (6.626 x 10-34 J • s)

4.Some metals hold their electrons more tightly than others, so different frequencies are needed to produce the photoelectric effect in different metals

G.. Einstein expanded on Planck’s theory to say light (or electromagnetic radiation) has a dual nature; it is both wave and particle • 1.He called a quantum of energy a photon – a packet of electromagnetic radiation that has zero mass, but carries a quantum of energy • 2.The energy of the photon depends on the frequency of the radiation • 3.If the frequency of a photon is not great enough, it will not cause the electron to come off the metal

IV. Hydrogen Atom Line Emission Spectrum • A.When electric current passes through a gas, the potential energy of the gas atoms increases • B.The lowest energy state of an atom is its ground state • C.When the energy level of an atom is higher than its ground state, this is called an excited state • D.When an atom is in an excited state, and then returns to its ground state, the energy from the excited state is released as electromagnetic radiation (an example occurs in a neon sign)

E.When electricity passed through Hydrogen gas, it gave a pink glow • F.When pink ray was passed through a prism, it separated into 4 colors • G.These four colors are hydrogen’s line-emission spectrum(Balmer series) • H.Other line emissions were discovered outside the visible range (ultraviolet – Lyman series) (infrared – Paschen series) • I.Scientists had believed that any amount of energy would excite an atom and the full range of frequencies would be observed

J.The full range of frequencies is the continuous spectrum • K.Trying to explain why only certain frequencies are emitted led to the study of the quantum theory • L.The fact that the electron of the hydrogen atom emitted only certain frequencies, suggested that it existed in specific energy states

V. Bohr Model 1913 • A. Electrons circle the nucleus • in certain paths or orbits • B. A Hydrogen atom is in its lowest energy state (ground state) when the electron is in orbit closest to the nucleus • C. An atom is in a higher energy state (excited state) when electrons are in orbits farther from the nucleus than normal • 1.When an electron falls to a lower level, energy is given off (emission)

2.When energy is added to an electron until it has enough to move up to the next level (absorption) • D. Bohr labeled the energy levels • 1.E1 – ground state • 2.E2 – E6 were successively higher energy levels • 3.He assumed many other levels existed beyond what had been observed (E∞)

Electron Configuration