Chapter 5: Electrons in Atoms Chemistry

5.1 Light and Quantized Energy Rutherford’s Nuclear Atomic Model was incomplete, did not explain How electrons were arranged Why electrons were not pulled into the nucleus The differences in chemical behavior

Wave Nature of Light In 1900s, observation of light when elements were heated were observed Analysis of light revealed an elements chemical behavior is related to the arrangement of electrons in its atoms

Electromagnetic Radiation Form of energy that exhibits wavelike behavior as it travels through space Includes visible light which is only a small portion of EMR Visible Spectrum – ROY G BIV

Electromagnetic Spectrum Chart

Wave

Wavelength Symbol: λ (Lambda) Shortest distance between equivalent points (Peak to Peak) Units = m, cm, nm nm=nanometers (10-9)

Frequency Symbol: ν (nu) Number of waves that pass a given point per second Units: hertz (Hz or 1/sec) 652 Hz=652 waves/sec=652/sec=652s-1

Waves (Continue) Wavelength and frequency have an inverse relationship As wavelength increases, frequency decreases Amplitude: Height of a Wave Speed and Amplitude are not affected by wavelength

Light All electromagnetic waves travel at 3.00 x 108 m/s in a vacuum This is known as the speed of light Symbol is c Speed of light is a product of wavelength and frequency c = λν All light travels at the same speed but with different wavelength and frequency

Practice Problems A helium-neon laser emits light with a wavelength of 655 nm. What is the frequency of the light? What is the wavelength of x-rays having a frequency of 4.80 x 1017 Hz? An FM radio station broadcasts at a frequency of 98.5 MHz. What is the wavelength of the station’s broadcast signal?

Electromagnetic Spectrum Electromagnetic Spectrum: classifies light based on wavelength and frequency Encompasses all forms of electromagnetic radiation Visible Spectrum: Roy G. Biv Only difference in types of radiation is wavelengths and frequencies Energy increases with greater frequency Violet light has shorter wavelength, greater frequency than red light, therefore violet light has more energy that red light.

Particle Nature of Light Wave model could not explain Why heated objects emit light at given temperatures Why some metals emit electrons when light shines on them Max Planck studied light emitted from heated objects, Discovered that matter can gain or lose energy in small, specific amounts called quanta Quantum: the minimum amount of energy that can be lost or gained by an atom Proposed emitted light (from glowing objects) is quantized Equantum= energy h= Planck’s constant= 6.626 x 10-34 J.s J= Joules = Kg .m2/s2 (SI unit of energy) ν= frequency

Plank’s Equation Equantum=hν Planck’s Theory= for a given frequency matter can emit or absorb energy only in whole number multiples ( 1hν, 2hν, 3hν, 4hν) As energy increases, frequency increases Equation explains why violet light has more energy than red light

Photoelectric Effect Einstein reasoned that light acted both wave like and particle like Reasoned that light acted like a stream of particles, knocking electrons out of atoms Photoelectric Effect = electrons called photo electrons are emitted from metals surface when light of a certain frequency shines on surface

Photoelectric Effect Light of a certain minimum frequency ejects electrons from the surface of metal

Planck and Einstein An energy increases, frequency increases

Practice Problems Calculate the energy of a gamma ray photon whose frequency is 5.02 x 1020 Hz. What is the difference in energy between a photon of violet light with a frequency of 6.8 x 1014 Hz and a photon of red light with a frequency of 4.3 x 1014 Hz? Calculate the energy of a photon of ultraviolet light that has a wavelength of 49.0 nm.

Atomic Emission Spectra Atomic Emission Spectra: when atoms of an element in the gaseous state are excited by energy they emit light The light can be broken down into a spectrum consisting of discrete lines of specific frequencies, or colors Atomic Emissions Spectra is unique for each element Example: Neon signs

5.2 Quantum Theory and the Atom Bohr Model of the Atom proposed quantum or planetary model of the atom Correctly predicted frequencies of the lines in H atomic emission spectrum Related H atomic energy states to motion of e- assigned Quantum numbers n to each orbit (n = energy level) Known as the principle quantum number n = 1, 2, 3, 4, 5, 6, 7 (numbers correspond to energy level as well as quantum numbers)

Problems Electrons do not orbit in circular paths Bohr’s model only worked for hydrogen Electrons do not orbit in circular paths

Louis de Broglie Accounted for the fixed energy level of Bohr’s model Proposed/thought that if waves have particle-like behavior then particles can have wave-like behavior. De Broglie Equation Predicts all moving particles have wave characteristics Example: cars (have wavelengths to small to be seen even with sensitive equipment)

Werner Heisenberg Heisenberg Uncertainty principle  it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time. We are certain about being uncertain Bumping into an electron while trying to determine its position and movement transfers energy and disrupts the electron

Erwin Schrödinger Treated hydrogen atoms electrons as a wave Applied well to other elements Led to quantum mechanical model of an atom (atomic model in which electrons are treated as waves) Wave equation was able to predict three-dimensional regions of an e- probable location (ATOMIC ORBITALS-electron cloud, fuzzy cloud)

Atomic Orbitals Principle Quantum numbers indicate the relative sizes and energies of atomic orbitals The atom’s major energy levels are called principal energy levels Principal energy levels contain sublevels The number of energy sublevels in a principal energy level increases as the quantum number increases

Principal Quantum Number n Corresponds to energy level and contains 2n2 electrons n = 1 contains 2e- [2(1)2] n = 2 contains 8e- [2(2)2] n = 3 contains 18e- [2(3)2] n = 4 contains 32e- [2(4)2] n = 5 contains 50e- [2(5)2] n = 6 contains 72e- [2(6)2] n = 7 contains 98e- [2(7)2

Energy Levels Contain Sub-Levels Known as Quantum number l s,p,d,f (Sublevels) s = 0, p = 1, d = 2, f = 3 Number of orbitals in a sub-level = (2l +1) Number of electrons in a sub-level = 2(2l +1)

Electron Configuration Electron configuration-arrangement of electrons in an atom, usually lowest amount of energy is more stable. Three Rules or principles that define how electrons can be arranged The Aufbau Principle The Pauli Exclusion Principle Hund’s Rule

The Aufbau Principle States that each electron occupies the lowest energy orbital available

Pauli Exclusion Principle States that a maximum of 2 electrons may occupy a single orbital, but only if the electrons have opposite spins ↑↓

Hund’s Rule States that a single electron with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbital

Orbital Diagrams and Electron Configuration Notation Orbital Diagram Method Electron Configuration Notation Noble Gas Configuration

Valence Electrons Electrons in the outer most orbitals Determines the chemical properties of an element Generally the electrons in the highest energy level Elections involved in bonding

Electron-Dot Structures element symbol that represents atomic nucleus and inner electrons surrounded by dots which represents valence electrons