Atomic Emission Spectra and the Bohr Model 1.2

Atomic Emission Spectra and the Bohr Model 1.2


A light wave is defined by its amplitude, wavelength, and frequency. 


Describing Light as Waves

  • The amplitude of a wave is the wave’s height from the center of its oscillation to the top, or crest

  • Wavelength (λ) is the distance between two wave crests

  • Frequency (v)  is the number of wave cycles to pass a given point per second (cycles per second)

  • The Sl unit of cycles per second is called the hertz (Hz)


Comparing Light Waves

  • The mathematical relationship between frequency (v) and wavelength (λ) is an inverse relationship. 

  • This means that as the wavelength of the wave decreases, the frequency increases

  • Wavelength decreases, frequency increases

  • Wavelength increases, frequency decreases

  • Wavelength and frequency will never both increase or decrease


The Electromagnetic Spectrum

  • Electromagnetic radiation consists of waves of energy that travel through a vacuum at a speed of 2.998 x 10^8 m/s in a vacuum. 

  • Includes all of the types of electromagnetic radiation (In lowest to greatest frequencies)

  • Radio waves

  • Microwaves

  • Infrared waves

  • Visible light (Lowest energy - red) (Highest energy - violet)

  • Ultraviolet waves

  • X-rays

  • Gamma rays 

  • Different types of electromagnetic radiation have different ranges of wavelengths

  • The mathematical relationship between frequency and wavelength of light using the electromagnetic spectrum: c= λv

  • The product frequency and wavelength equals a constant (c), the speed of light

  • Because light is constant, if the wavelength of a wave increases, the frequency of the wave must decrease.

  • Electromagnetic radiation with longer wavelengths (low- frequency) is used for communication

  • High-frequency (short wavelengths) waves carry enough energy to cause damage to living tissue


Emission Spectra

  • When an element absorbs energy in a gaseous state, electrons become excited, causing the atom to become unstable

  • When electrons release energy, the atom can emit light as it returns to its more stable state

  • The light an element emits is composed of specific wavelengths of light. 

  • The atomic emission spectrum is the pattern formed when light emitted by an element is separated into the different wavelengths of light it contains. 

  • Different wavelengths of light result in different colors of lights.


Separating Colors

  • A spectroscope is a tool that uses a prism to separate light into its individual colors.

  • Each color of light has a characteristic wavelength

  • A spectroscope shows different outputs for different inputs.

  • When the input is white light from a incandescent bulb, the output from the spectroscope is a continuous spectrum of all colors and visible wavelengths

  • When the input is light from a hydrogen lamp, the output is a few discrete lines. 

  • The pattern of the wavelengths of these lines, its atomic emission spectrum, is unique to hydrogen.

Elemental Fingerprints

  • Each element emits specific colors of light that correspond to specific wavelengths.

  • The atomic emission spectrum of each element is like a person’s fingerprint. 

  • No two elements will have the same atomic emission spectrum

  • No matter where a sample of a given element is collected, the spectrum will be the same.

  • Atomic emission spectra can be used to identify unknown samples. 

  • Scientists study atomic emission spectra of the stars to learn their elemental composition

  • Each spectral line corresponds to one wavelength of visible light emitted by that element.


Energy Levels in Atoms

  • In 1913, Niels Bohr, a young Danish physicist, furthered the development of modern atomic theory by constructing a model (Bohr’s Nuclear Atom)

  • This model described electrons as moving in circular orbits around the nucleus. 

  • Each orbit in his model has a fixed energy.

  • The fixed energies an electron can have are called energy levels.

  • An electron can move, or transition, from one energy level to the other, but it cannot exist between levels.

  • To move from one energy level to another, an electron must gain or lose just the right amount of energy.

  • The amount of energy required to move an electron from one level to another is a quantum of energy

  • This is the reason why the energy of an electron is said to be quantized

  • The size of a quantum of energy can vary

  • The amount of energy an electron gains or loses in an atom is not always the same because energy in an atom are not equally spaced

  • Higher energy levels are closer together than lower energy levels


Energy Levels/ Electron Shells

  • Energy Level 1 - 2 electrons

  • Energy Level 2 - 8 electrons

  • Energy Level 3 - 18 electrons

  • Energy Level 4 - 32 electrons


Continuous Energy

  • A person climbing a ramp can take big or small steps

  • If energy levels in atoms were continuous, electrons could absorb or give off any amount of energy as they move up/down energy levels

Quantized Energy

  • A person climbing stairs can only take a certain sized step to move up/down the stairs

  • The size of the step is quantized 

  • Electrons must absorb/give off a certain amount of energy to move between energy levels

  • The size of a step is similar to a quantum of energy


Explaining Emission Spectra

  • Bohr constructed a conceptual model of the atom

  • This model correctly related the frequencies and wavelengths of light to the quantization of energy in the emission spectrum of hydrogen

  • Bohr become one of a diverse group of scientists who had an impact on scientific thought around atomic structure

  • Bohr’s model of the nuclear atom was as important step in the development of modern atomic theory


Understanding Bohr’s Atomic Model

Why do hydrogen atoms emit specific wavelengths of light according to Bohr's Model?

  • Bohr developed his atomic model to explain the hydrogen emission spectrum

  • His model could not explain the emission spectra of other elements because electrons do not move in a circular paths (as he thought)

Ground State

  • The electron is in the lowest possible energy level

-An atom absorbs a specific amount of energy, causing it to jump to a higher energy level

Excited State

  • The electron has gained energy and is in a higher energy level

- The atom emits a specific amount of energy as the electron returns to a lower energy level.

- The amount of energy associated with the emitted light is related to the frequency and wavelength of the light wave.

Emission Spectra Lines

  • Hydrogen atoms absorb/emit light with an energy exactly equal to the energy difference between energy levels in the atom

  • Each spectral line has a specific wavelength and frequency related to the energy as an electron transitions to a lower energy level


The Quantization of Energy

  • The idea for quantized energy came from the German physicist Max Planck

  • Explained mathematically why an atomic emission spectrum consists of distinct wavelengths of light in 1900

  • Calculated through trial and error that 6.626 x 10^-34 Joule-seconds defined the smallest wavelength of a photon

  • This value (Planck’s constant (h)) relates the magnitude of a quantum of energy released as an electron changes energy levels to the corresponding frequency of the radiation (v) in the electromagnetic spectrum

E= hv

  • Since the spacing between electron energy levels varies based on the element, each element emits its own distinct frequencies of light


Bohr Model Representations of Atoms

  • The Bohr Model is not completely correct

  • Scientists have since learned more about the atom’s structure through experiments and calculations

  • Many people still prefer to use the Bohr model today for a few reasons

  • The Bohr model is a simplified picture of an atom with many features that are nearly correct

  • Many important properties of atoms can be exemplified using this model