Chemistry Chapter 5
frequency (nu)
# of wave cycles to pass a point per unit of time; measured in s^-1 or 1/s
Bohr Model/Planetary model of the atom
-electron revolves around the nucleus in a circular orbit -orbits have specific distances from the nucleus (energy levels) -energy level repres. by n -ground state is the most stable state
Magnetic quantum number (m with a small l)
3D orientation in space of the energy sublevel
Visible spectrum
All the colors combined give you white light. What is this spectrum called?
Hund's rule
Degenerate orbitals must be filled before you move to another
Electromagnetic radiation
Energy waves that travel at 3.0 x 10^8 m/s in a vacuum
Clinton Davisson (Nobel) and Lester Germer's experiments
Experiments by who confirmed that matter moves like a wave?
Niels Bohr
Explained the emission spectrum of hydrogen
Albert Einstein (Nobel Prize in 1921)
Explained the photoelectric effect with quantum theory
low/short
High frequency corresponds to _______ wavelength
Bohr's model describes electrons in fixed orbits rather than in a cloud and with fixed energy levels.
How did Bohr's model differ from that of Rutherford?
They emit energy and move down an energy level.
How do electrons become stable again?
E=hv
How much is one quantum of energy?
It must emit energy (in the form of light)
How will an electron move down an energy level?
It must gain energy/become excited/reach and excited state
How will an electron move up an energy level?
Wave particle duality (De Broglie, 1924)
If light can behave like a particle then particles can behave like waves.
high/long
Low frequency corresponds to _______ wavelength
Discrete energy levels and energy is absorbed/emitted
Name 2 key ideas from Bohr's model.
Hertz (Hz)
SI Unit of cycles per second (s^-1 or 1/s)
Einstein's experiment on the photoelectric effect
Shone light on alkali metals and the electrons were knocked out of the metal
Excited states
States higher than the ground state
Newton
Studied the continuous spectrum of the sun
Electron configurations
Tells you where electrons of an atom are
Max Planck
The scientist that tried to explain color change with temperature increase
False: There are no quarter or half energy levels. Jumps occur from energy level to energy level. There is no in between.
True or False: Electrons can be in a half energy level.
True: lower energy does correspond to greater stability
True or False: Lower energy = greater stability
True (they will always have a different ms)
True or False: No 2 electrons have the same address in a particular element
False: Two electrons in an orbital can go in opposite directions
True or False: Two electrons in an orbital can go in the same direction
False: Black lines correspond to wavelengths that the element absorbed. Colored/bright lines correspond to wavelengths that the element emitted.
True or false: Black lines correspond to wavelengths that the element emitted.
False; an element's atomic emission spectrum is characteristic to that element
True or false: Some elements have the same atomic emission spectrum.
True (wavelength times frequency equals the speed of light always)
True or false: The speed of all electromagnetic radiation is the same/constant
False: Light moves like a wave but transfers energy like a stream of particles in photons/quanta
True or false: light moves like a particle but transfers energy like a wave
True
True or false: mass needs to be small for wavelength to be large enough to observe
They bombarded metals with beams of electrons. The electrons reflected and produced curious patterns. Th electrons reflected like waves.
What did Davisson and Germer do in their experiments?
Photons
What did Einstein call light quanta?
The Quantum Mechanical Model
What did Schrodinger's equation lead to?
It couldn't explain why extra lines appear when a magnetic field is applied, and it couldn't account for atoms of more than one electron.
What did the Bohr model fail to explain?
Where an electron is likely to be found.
What do orbitals describe?
3.0 x 10^8 m/s (wavelength x frequency)
What does the speed of light equal?
Why heated objects only emit certain colors; photoelectric effect
What does the wave property of light not explain?
The frequency/electricity of the light is important; how much light and how long it is shone on the metal surface does not effect the phenomenon
What factors are important in the photoelectric effect?
Water droplets forming a rainbow because the droplets refract the light.
What is a real life example of a prism that produces a spectrum?
4 lines (purple, blue, green, red)
What is the emission spectrum of hydrogen like?
energy=planck's constant x nu
What is the equation for energy?
400 to 700 nm
What is the wavelength range of visible light?
Wavelengths of spectral lines
What makes up an element's atomic emission spectrum?
They couldn't describe the microscopic properties of atoms (stability, electron configurations, etc.)
What was the failure of 19th century physicists?
Only hydrogen (1 electron); no explanation of bonding in molecules
What was the failure of the Bohr model?
They were unable to explain the behavior of electrons in an atom
What were classical physicists unable to explain?
Unstable
When electrons in an atom absorb energy and become excited they move to a higher energy level and become...
Chromium (24) and Copper (29)
Which elements are exceptions to the Aufbau principle?
High frequency/energy light such as blue light
Which type of light can cause the photoelectric effect?
Max Planck (different colors of heated metals)
Who analyzed data on radiation emitted by solids at various temperatures?
Louis de Broglie (1924)
Who asked if particles of matter can behave as waves?
Planck
Who began the work of Quantum Theory?
Thomas Young (1801)
Who observed a pattern of dark and lights stripes when light was passed through 2 adjacent slits?
James Clerk Maxwell (1873)
Who proposed that light consists of electromagnetic waves?
Einstein
Who proposed that light was a quanta of energy that behaved like particles?
Thomas Young
Who proved that light is wave-like?
De Broglie
Who said that matter is both wave-like and particle-like?
Newton
Who showed that light is made of a rainbow of colors?
Einstein (1905)
Who studied the photoelectric effect?
Isaac Newton
Who tried to explain what was known about the behavior of light by assuming that light consists of particles?
Because it is unique for each element
Why is the atomic emission spectrum a fingerprint?
emission spectrum
a spectrum of bright lines with a black background
absorption spectrum
a spectrum of dark lines with a rainbow background
Line spectrum
a spectrum with parts that are black (emission/absorption)
X-ray crystallography
a tool used for identifying the atomic and molecular structure of a crystal, in which the crystalline atoms cause a beam of incident X-rays to diffract into many specific directions
White light
consists of light with a continuous range of wavelengths and frequencies; combination of all light
Classical Mechanics
described motion of large bodies
Principal Quantum number (n)
describes the energy level
Principal Quantum Number n
describes the energy level on which the orbital resides; size and energy of the orbitals
wavelength (lambda)
distance between crests (top of one crest to the top of another); peak to peak distance measured in meters
Lyman series (ultraviolet)
due to transition form higher energy levels to n=1 in the emission spectrum of hydrogen
Balmer series (visible)
due to transition from higher energy levels to n=2
Paschen series (infrared)
due to transition from higher energy levels to n=3 (or higher)
p orbitals
dumbbell shape; l=1
Photoelectric effect
electrons are ejected when light shines on a metal
Hund's rule
electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible
energy level
fixed energies and electron can have
wavelength of a particle = h/m x v (De Broglie)
h/m x v
Planck's constant for E=h x nu
h=6.626 x 10^-34 J s
Interaxial orbitals
lobes that are between axes
Erwin Schrodinger (1926)
made and solved an equation describing the behavior of the electron in a hydrogen atom
Atomic orbitals
mathematical expression from the Schrodinger equation describing the probability of finding an electron at various locations around the nucleus
Bohr's atomic model
model in which electrons orbit around the nucleus in fixed energy levels
The Quantum Mechanical Model
modern mathematical description of electron behavior (does not specify path electrons must take)
Quantum mechanics
motion of subatomic particles as waves
William Wallaston
observed dark lines in the sun's spectrum
Degenerate orbitals
orbitals of equal energy
Aufbau Principle
orbitals of lowest energy 1st; orbitals of greater energy are higher on the diagram; sublevels of a principal energy level are of lowest energy
Axial orbitals
orbitals that lie along the axes
spin
quantum mechanical model property of electrons; clockwise or counterclockwise
Heisenberg Uncertainty Principle
says that it is impossible to know the position and velocity of a particle at the same time
s orbitals
spherical shape; l=0
M and little s
tells you the direction in which the electron is spinning (clockwise or anticlockwise)
Azimuthal/Angular momentum Quantum Number (italicized l)
tells you which sub energy level the electron is in; defines shape of the orbital (range from 0 to n-1)
quantum
the amount of energy required to move an electron from one energy level to another; minimum amount of energy that can be gained/lost by an electron
Pauli Exclusion Principle
the atomic orbital describes up to 2 electrons with opposite spins
electromagnetic radiation
the emission and transmission of energy in the form of electromagnetic waves
threshold value
the minimum amount of energy needed to knock out the electrons
Rutherford's atomic model
the model in which electrons circle around the nucleus like the planets and sun; didn't explain chemical properties of elements
electricity
the movement of electrons
Valence shell
the outermost energy level/shell
Spectroscopy (used to determine the elements in stars)
the science of using spectral lines to figure out what something is made of
Continuous spectrum (has no lines of absorption)
the sun's atomic emission spectrum
Electron configurations
the ways in which electrons are arranged in various orbitals around the nuclei of atoms
wave
vibrating disturbance that moves outward from a disturbance; destructive and constructive interference
Red light
visible light with the longest wavelength and lowest frequency
amplitude
wave's height from 0 to crest; peak height above the mid-line measured as A
Ground state
when an electron has its lowest possible energy level (principal quantum number = 1)
Node
where there is 0 probability of finding an electron (nucleus)
