Chem Unit 9 - The Development of a New Atomic Model

The Development of a New Atomic Model

J. J. Thomson’s “plum pudding” model

• electrons are surrounded by a "soup" of positive charges to balance the electron’s negative charge,

• like negatively-charged “plums” surrounded by positively charged “pudding”.

Past Models

The Rutherford model

• Ernest Rutherford's model was an improvement over previous models, but it was incomplete.

• It did not explain how the atom’s negatively charged electrons are arranged in the space surrounding its positively charged nucleus.

In the early twentieth century, a new atomic model evolved as a result of investigations into the absorption and emission of light by matter.

• The studies revealed a relationship between light and an atom’s electrons.

Early 20th Century

Before 1900, scientists thought light behaved only as a wave.

• This belief changed when it was later discovered that light also has particle-like characteristics.

Visible light is a kind of electromagnetic radiation (form of energy that exhibits wavelike behavior as it travels through space).

Properties of Light

• Other examples include X-rays, ultraviolet and infrared light, microwaves, and radio waves.

• The electromagnetic spectrum consists of all types of electromagnetic radiation

Properties of a Wave

All forms of electromagnetic radiation move at a speed of 3.0 x 10⁸ m/s through air.

• Wavelength – λ (m, cm, or nm) and frequency – ν (wave/second) are measurable properties of wave motion.

  • One wave/second is called a hertz (Hz).

• The relationship between wavelength (λ) and frequency (ν) is c = λν where c = speed of light.

Fill in the Blank

Which wave has the higher frequency?  Explain your answer.

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Fill in the Blank

Which wave has the longer wavelength? Explain your answer.

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Photoelectric Effect

In the early 1900s, scientists conducted two experiments involving interactions of light and matter that could not be explained by the wave theory of light.

• One experiment involved a phenomenon known as the photoelectric effect.

  • The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal.

​electromagnetic radiation strikes the surface of the metal, ejecting electrons from the metal and causing an electric current.

The Wave Theory of Light

The wave theory of light predicted that light of any frequency could supply enough energy to eject an electron.

Scientists couldn’t explain why the light had to be of a certain frequency in order for the photoelectric effect to occur.

• The German physicist Max Planck proposed an explanation for the photoelectric effect.

  He proposed that a hot object does not emit electromagnetic radiation continuously, as would be expected if the energy emitted were in the form of waves.

Max Planck

Max Planck proposed that objects emit energy in small, specific amounts called quanta (1900).

• a quantum is the minimum quantity of energy that can be lost or gained by an atom.

• The relationship between a quantum of energy and the frequency of radiation is illustrated by the following equation:

  • E = hν

    • E is the energy, in joules, of a quantum of radiation, ν is the frequency in s^-1 of the radiation emitted, and h is a physical constant now known as Planck’s constant. (6.626 × 10^-34 m² kg/s)

Einstein

Einstein proposed that electromagnetic radiation has dual wave-particle nature (1905). These particles are called photons.

• A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy.

  In order for an electron to be ejected from a metal surface, the electron must be struck by a single photon possessing the minimum energy and frequency to knock it loose.

• - The energy of a particular photon depends on the frequency of the radiation. E_photon = hν

Hydrogen Atom Line-Emission Spectrum

The ground state is the lowest energy level of an atom. When it has higher potential energy an atom is in its excited state.

• When an excited atom returns to its ground state it gives off energy in the form of colored light. (Example: Neon lights.)

Multiple Choice

1. What is the difference between the ground state and the excited state of an atom?

• ground state means high potential energy, excited means low potential energy

• they are the same thing (there's no difference)

• ground state is the lowest energy level of an atom. excited states have higher potential energy

Multiple Choice

1. true or false: When an ground state atom jumps to an excited state it gives off energy in the form of colored light.

• true

• false

Experimenting with Light

excited hydrogen atoms emit a colored glow. When the visible portion of the emitted light is passed through a prism, it is separated into specific wavelengths that are part of hydrogens line-emission spectrum.

Experimenting with Light

When doing experiments with hydrogen gas, it was found that hydrogen atoms emit only specific frequencies of light.

• The fact that hydrogen atoms emit only specific frequencies of light indicated that the energy differences between the atom’s energy states were fixed.

• This suggested that the electron of a hydrogen atom exists only in very specific energy states (led to quantum theory).

Bohr Model of the Hydrogen Atom

Niels Bohr proposed a model of the hydrogen atom that showed that the electron can circle the nucleus only in allowed paths (orbits) (1913).

• he suggested that the electron could jump to a different orbit by absorbing or emitting a photon of light (indicating certain allowed circular orbits).

Does it work all of the time?

Bohr's model fit the hydrogen atom very well, but when applied to other atoms it did not work.

• it is important to understand that the current theory of atomic structure is not the same as the Bohr model.

• electrons do NOT move around the nucleus in circular orbits like planets in the solar system orbiting the sun.

  • we will learn more about the current model in a later lesson.