Download
1 / 31
320 likes | 411 Views
Electrons in Atoms. Wave Nature of Light. Electromagnetic radiation is a form of energy that exhibits wave-like behavior as it travels through space. Atoms give off electromagnetic radiation when the electrons that are excited drop back to their original energy level.
E N D
Wave Nature of Light • Electromagnetic radiation is a form of energy that exhibits wave-like behavior as it travels through space. • Atoms give off electromagnetic radiation when the electrons that are excited drop back to their original energy level.
Types of electromagnetic radiation include • Sunlight or visible light • Microwaves, infrared radiation, ultraviolet radiation • Radiowaves, X-rays, and gamma rays
Characteristics of Waves • Wavelength (λ): the shortest distance between equivalent points on a continuous wave. • It is measured from trough to trough or from crest to crest. • It is usually expressed in meters, centimeters, or nanometers (1 nm = 1x10-9) • Frequency (ν): the number of waves that pass a given point per second. • Hertz (Hz) is the SI unit of frequency • 1 Hz = 1 wave/second
Amplitude: the wave’s height from the origin to a crest or from an origin to a trough. • Speed: all waves travel at the speed of light (c), which is 3.00 x 108 m/s in a vacuum. • λν = 3.00 x 108 m/s
Particle Nature of Light • Max Planck (1858-1947): discovered that matter can gain or lose energy in small, specific amounts that he called quanta. • A quantum is the minimum amount of energy that can be gained or lost by an atom. • Energy of a quantum is related to the frequency of the emitted radiation. • Equantum= hν - where E = energy, h = Planck’s constant, ν= frequency - Planck’s constant = 6.626 x 10-34 Js
Photons and Energy • A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy. • A photon’s energy depends on its frequency. • Ephoton = hν - where E = energy, h = Planck’s constant, ν= frequency
Atomic Emission Spectra • The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by atoms of the element. • Each element’s atomic emission spectrum is unique and can be used to determine if that element is part of an unknown compound.
Quantum Theory of the Atom • Niels Bohr created an atomic model that showed that electron move around the nucleus in only certain allowed circular orbits. • The smaller the electron’s orbit, the lower the atom’s energy state or energy level. • Atoms at the lowest energy state are in their ground state. • Atoms that gain energy jump to a higher orbit and are in an excited state.
Bohr assigned a quantum number (n) to each orbit and the relative energy level of each orbit: n = 1, E1 n = 2, E2 = 4E1 n = 3, E3 = 9E1 n = 4, E4 = 16E1 n = 5, E5 = 25E1 n = 6, E6 = 36E1 n = 7, E7= 49E1
When an electron returns to its ground state, energy is released in the form of a photon of light. ΔE = Ehigher orbit – Elowerorbit = Ephoton = hν • Bohr’s model explain the light spectrum produced by different element, but it was not correct when placing the electrons in circular orbits.
Quantum Mechanical Model of the Atom • Louis de Broglie (1892-1987) suggested that electrons move in a wavelike motion and are restricted to circular orbits of fixed radius. • λ= h/mν wavelength = Planck’s constant mass x velocity
Heisenberg Uncertainty Principle • States that it is impossible to make any measurement on an object without disturbing the object. • It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time. • You change one or the other when you take a measurement.
Schrodinger Wave Equation • Electrons are treated as waves and limits an electron’s energy to certain values. • No attempt is made to describe the electron’s path around the nucleus. • This model of the atom is called the quantum mechanical model. • Atomic orbitals in this model describe an electron’s POSSIBLE location.
Hydrogen’s Atomic Orbitals • Orbitals do not have an exactly defined size. • Principle quantum numbers (n) are assigned that indicate the relative sizes and energies of each atomic orbital. • As n increases, the orbital becomes larger, the electron spends more time farther from the nucleus, and the atom’s energy level increases. • Therefore, n specifies the atom’s major energy levels, called principal energy levels.
Principal Energy Levels • Contain sublevels • Principal energy level 1 has one sublevel • Principal energy level 2 has two sublevels • Principal energy level 3 has three sublevels, etc. • Sublevels are orbitals labeled s, p, d, and f • s orbitals are spherical and increase in size with energy level • p orbitals are dumb-bell shaped • d and f orbitals have varying shapes
Hydrogen’s First Four Principal Energy Levels Hydrogen has one electron and it spends most of its time in the 1s sublevel when not excited.
Electron Configurations • Ground-State Electron Configuration • Low-energy system are more stable than high-energy systems. • Electrons in an atom tend to assume the arrangement that gives the atom the lowest possible energy. • The most stable, lowest-energy arrangement of the electrons in an atom is called the element’s ground-state configuration. • Three rules, or principles, define how electrons can be arranged in an atom’s orbitals.
The Aufbau Principle • Each electron occupies the lowest energy orbital available. • All orbitals within an energy sublevel have equal energy. • Energy sublevels have different energies: s, p, d, and f. • Orbitals within one sublevel can overlap orbitals of a different energy level.
The Pauli Exclusion Principle • Each electron in an atom has an associated spin and can only spin in one of two directions. • represents the electron spinning in one direction and the represents another electron spinning in the other direction. • A maximum of two electrons may occupy a single atomic orbital, but only if the electrons have opposite spins. • An atomic orbital containing paired electrons with opposite spins is written as .
Hund’s Rule • Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals. • One electron enters each of the three 2p orbitals before a second electron enters any of the orbitals.
