enriched chemistry chapter 4 – arrangement of electrons in atoms
DESCRIPTION
The Wave Description of Light Visible light is a kind of electromagnetic radiation; Together, all forms electromagnetic radiation form the electromagnetic spectrum; All forms e.r. move at a constant speed of 3.00 x 108 m/s through a vacuum and at slightly slower speeds through matter;TRANSCRIPT
Enriched Chemistry Chapter 4 Arrangement of Electrons in
Atoms
Section One The Development of a New Atomic Model The Wave
Description of Light
Visible light is a kind of electromagnetic radiation; Together, all
forms electromagnetic radiation form the electromagnetic spectrum;
All forms e.r. move at a constant speed of 3.00 x m/s through a
vacuum and at slightly slower speeds through matter; The
Electromagnetic Spectrum
HIGH ENERGY LOW ENERGY Describing Waves Wavelength () - length of
one complete wave
Frequency () - # of waves that pass a point during a certain time
period hertz (Hz) = 1/s Amplitude (A) - distance from the origin to
the trough or crest crest origin trough A c = c: speed of light
(3.00 108 m/s) : wavelength (m, nm, etc.)
Frequency & wavelength are inversely proportional c = c:speed
of light (3.00 108 m/s) :wavelength (m, nm, etc.) :frequency (Hz)
Calculate the wavelength (in meters) of radiation a frequency of
5
Calculate the wavelength (in meters) of radiation a frequency of
5.00 x 1014 s1. When certain frequencies of light strike a metal,
electrons are emitted.
Photoelectric effect refers to the emission of electrons from a
metal when light shines on the metal. Scientists observed that for
a specific metal, no electrons were emitted if the lights frequency
was below a certain minimum, regardless of the intensity. This
puzzled scientists because it was not predicted by the wave theory
of light. https://www.youtube.com/watch?v=v-1zjdUTu0o The Particle
Description of Light
In 1900, German physicist Max Planck suggested that hot objects
emit energy in small, specific packets called quanta. A quantum of
energy is the minimum quantity of energy that can be lost or gained
by an atom. The energy of a photon is proportional to its
frequency.
E = h E: energy (J, joules) h: Plancks constant ( Js) : frequency
(Hz) In 1905, Einstein introduced the idea of photons, which are
particles of e.r. having zero mass and carrying a quantum of
energy. Electrons exist only in very specific energy states for
every atom of each element.
Ground state the lowest energy state of an atom. Excited state the
atom has a higher potential energy than it has in its ground state.
Basically, when an atom absorbs energy it moves to an excited
state; When that atom returns to its ground state, it releases
energy in the form of e.r. Neon signs are an example. Hydrogens
Line-Emission Spectrum
Investigators passed electric current through a vacuum tube
containing hydrogen gas at low pressure, they observed the emission
of a characteristic pinkish glow. When a narrow beam of the emitted
light was shined through a prism, it was separated into four
specific colors of the visible spectrum. The four bands of light
were part of what is known as hydrogens line-emission spectrum.
Scientists had expected to observe the emission of a continuous
range of frequencies of electromagnetic radiation, a continuous
spectrum. Hydrogens Line-Emission Spectrum Bohr Model of the
Hydrogen Atom
Niels Bohr proposed a hydrogen-atom model that linked the atoms
electron to photon emission. According to the model, the electron
can circle the nucleus only in allowed paths, or orbits. The energy
of the electron is higher when the electron is in orbits that are
successively farther from the nucleus. When an electron falls to a
lower energy level, a photon is emitted, and the process is called
emission. Energy must be added to an atom in order to move an
electron from a lower energy level to a higher energy level. This
process is called absorption. video Photon Emission and Absorption
Section Two The Quantum Model of the Atom
In the same way that no two houses have the same address, no two
electrons in an atom have the same set of four quantum numbers. In
this section, you will learn how to use the quantum-number code to
describe the properties and locations of electrons in atoms.
Electrons have wave-like properties.
French scientist Louis de Broglie suggested in 1924 that electrons
be considered waves confined to the space around an atomic nucleus.
It followed that the electron waves could exist only at specific
frequencies. According to the relationship E = h, these frequencies
corresponded to specific energiesthe quantized energies of Bohrs
orbits. Heisenbergs Uncertainty Principle
German physicist Werner Heisenberg proposed that any attempt to
locate a specific electron with a photon knocks the electron off
its course. Electrons are detected by their interactions with
photons. The Heisenberg uncertainty principle states that it is
impossible to determine simultaneously both the position and
velocity of an electron or any other particle. Atomic Orbitals and
Quantum Numbers
https://www.youtube.com/watch?v=9E3QaRxqXZc Atomic Orbitals and
Quantum Numbers
Quantum numbers specify the properties of atomic orbitals and the
properties of electrons in orbitals. The principal quantum number,
symbolized by n, indicates the main energy level occupied by the
electron. Quantum numbers (cont)
The angular momentum quantum number, symbolized by l, indicates the
shape of the orbital. We will learn about the s, p, d and f orbital
shapes. N = 1 has one sublevel (s) N = 2 has two sublevels (s, p) N
= 3 has three sublevels (s, p, d) N = 4 has four sublevels (s, p, d
and f) Atomic Orbitals and Quantum Numbers, continued
The magnetic quantum number, symbolized by m, indicates the
orientation of an orbital around the nucleus. The spin quantum
number has only two possible values(+1/2 , 1/2)which indicate the
two fundamental spin states of an electron in an orbital. Lets
review two terms. Orbital a single allowed location for electrons
capable of holding two electrons of opposite spin states. Sublevel
includes all of the similarly shaped orbitals in an energy level.
Shapes of s, p, and d Orbitals Electrons Accommodated in Energy
Levels and Sublevels Electron Configurations
The arrangement of electrons in an atom is known as the atoms
electron configuration. The lowest-energy arrangement of the
electrons for each element is called the elements ground-state
electron configuration. Rules Governing Electron
Configurations
According to the Aufbau principle, an electron occupies the
lowest-energy orbital that can receive it. According to the Pauli
exclusion principle, no two electrons in the same atom can have the
same set of four quantum numbers.