Atomic Structure and PeriodicityAdvanced ChemistryMs. Grobsky

Food For Thought Rutherford’s model became

known as the “planetary model”

The “sun” was the positively-charged dense nucleus and the negatively-charged electrons were the “planets”

The Planetary Model is Doomed! The classical laws of motion and gravitation

could easily be applied to neutral bodies like planets, but NOT to charged bodies such as protons and electrons According to classical physics, an electron in orbit

around an atomic nucleus should emit energy in the form of light continuously because it is continually accelerating in a curved path

Resulting loss of energy implies that the electron would necessarily have to move close to the nucleus due to loss of potential energy Eventually, it would crash into the nucleus and the

atom would collapse!

The Planetary Model is Doomed!

Electron crashes into the nucleus!?

Since this does not happen, the Rutherford model could not be accepted!

The Bridge Between the Planetary Model and the Bohr Model Atomic structure was often elucidated

by interaction of matter with light Classical wave theory of light described

most observed phenomenon until about 1900

But What Exactly is Light? Light is a form of ELECTROMAGNETIC

RADIATION A form of energy that exhibits wavelike

behavior as it travels through space Does not require a medium to travel

through In a vacuum, every electromagnetic

wave has a velocity (speed) of 3.00 x 108 m/s, which is symbolized by the letter “c”

Some Properties of Waves Wavelength (λ)

Distance between two consecutive peaks or troughs in a wave

Measured in meters (SI system)

Frequency (ν) Number of waves that

pass a given point per second

Measured in hertz (sec-1) Speed ( c )

Measured in meters/sec Amplitude (A)

Distance from maximum height of a crest to the undisturbed position

Relationships of EM Wave Properties The wavelength and frequency of light are

inversely proportional to each other As wavelength increases, frequency

decreases As wavelength decreases, frequency

increases Wavelength and frequency are related via

the speed of light in a vacuum (c) c = 3.00 x 108 m/s

Speed of light in a vacuum is a constantc = λ· ν

The Electromagnetic Spectrum

Electromagnetic spectrum is the range of all possible frequencies of electromagnetic radiation

The highest energy form of electromagnetic waves is gamma rays and the lowest energy form is radio waves

Relationship of EM Wave Properties

c = λ· ν

Max Planck• Around the year 1900, a physicist

named Max Planck introduced his hypothesis of the quantum behavior of radiation

• A major turning point in physics!

• But what is quantum???

Max Planck’s ObituaryTake a few minutes to survey the reading (skim/scan the text) Turn headings and subheadings into questions related to who, what, where, when, why, or how? Write these questions on the LEFT side of the reading Write down key ideas found during reading on RIGHT side of reading (answers to questions, significant information, reflections) Summarize key ideas at BOTTOM of reading

Planck and Quanta Planck studied the energy given off by

heated objects until they glow He made a wild assumption that there is

a fundamental restriction on the amounts of energy that an object emits or absorbs He called these pieces “quanta”

To understand quantization, consider walking up a ramp versus walking up the stairs For the ramp, there is a continuous change

in height whereas up stairs, there is a quantized change in height

More on the Idea of Quanta The energy possessed by the wave is

only related to the frequency of the wave

The frequency of an electromagnetic wave can be converted directly to energy by:

h = Planck’s constant = 6.626 x 10-34 J· s

Planck’s Constant Planck’s constant, h, is just like a penny

Planck determined that all amounts of energy are a multiple of a specific value, h

This is the same as saying that all currency in the US is a multiple of the penny

Practice, Practice, Practice Calculate the energies of one photon of

UV (λ = 1 x 10-8 m), visible (λ = 5 x 10-7 m), and IR (λ = 1 x 10-4 m).

And then there was a problem…

• In the early 20th century, several effects were observed which could not be understood using the wave theory of light

• Every element emits light when energized either by heating the element or by passing electric current through it

• Elements in solid form glow when they are heated

• Elements in gaseous form emit light when electricity passes through them

Einstein and the Photoelectric Effect Another observation that could not be explained via the

wave theory of light: The Photoelectric Effect Electrons are attracted to the (positively charged) nucleus by

the electrical force

In metals, the outermost electrons are not tightly bound, and can be easily “liberated” from the shackles of its atom It just takes sufficient energy

If light was really a wave, it was thought that if one shined light of a fixed wavelength on a metal surface and varied the intensity (made it brighter and hence classically, a more energetic wave), eventually, electrons should be emitted from the surface

Photoelectric Effect

• No electrons were emitted until the frequency of the light exceeded a critical frequency, at which point electrons were emitted from the surface! (Recall: small l large n)

Vary wavelength, fixed amplitude

What if we try this ?

electrons emitted ?

No

Yes, withlow KE

Yes, withhigh KE

Increase energy by increasing amplitude

“Classical” Method

electrons emitted ?

No

No

No

No

Einstein’s Theory Einstein proposed an alternative theory to

the classical wave theory of light He used Planck’s idea of energy quanta to

understand the photoelectric effect Light exists as ‘quanta’ of energy (specific

amounts) These quanta behave like particles Light ‘particles’ are known as photons Each photon carries an amount of energy

that is given by Planck’s equation

Einstein’s Photons and the Photoelectric Effect

The light particle must have sufficient energy to “free” the electron from the atom

Increasing the Amplitude is simply increasing the number of light particles, but its NOT increasing the energy of each one!

However, if the energy of these “light particle” is related to their frequency, this would explain why higher frequency light can knock the electrons out of their atoms, but low frequency light cannot

“Light particle”

Before Collision After Collision

The Dual Nature of Light A “Waveicle”

Light travels through space as a wave

Light transmits energy as a particle

Each photon carries an amount of energy that is given by Planck’s equation

So is Light a Wave or a Particle ?

• On macroscopic scales, we can treat a large number of photons as a wave

• When dealing with subatomic phenomenon, we are often dealingwith a single photon, or a few• In this case, you cannot use the wave description of

light• It doesn’t work!

The Dualism of Light Dualism is not such a strange concept Consider the following picture

Are the swirls moving, or not, or both?

But How is This Related to the Atom?

Light and the Dilemma of Atomic Spectral LinesExperiments show that when white light is passed through a prism, a continuous spectrum results

Contain all wavelengths of light

When a hydrogen emission spectrum in visible region is passed through a prism, a line spectrum results

Only a few wavelengths of visible light pass through

Seeing Atomic Spectral Lines Use your diffraction grating to observe

the atomic spectra of: Hydrogen Oxygen Neon

hydrogen (H)

mercury (Hg)

neon (Ne)

Planck’s Quanta and Atomic Spectra To produce a line spectrum, the

electrons in an atom move between energy levels

Electrons typically have the lowest energy possible (ground state), but upon absorbing energy via heat or electricity: Electrons jump to a higher energy level,

producing an excited and unstable state Those electrons can’t stay away from the

nucleus in those high energy levels forever so electrons would then fall back to a lower energy level

Just a Thought…But if electrons are going from

high-energy state to a low-energy state, where is all this extra energy going?

Connecting Planck’s Quanta to the Atomic Model Energy does not disappear

First Law of Thermodynamics!

Electrons re-emit the absorbed energy as photons of light Difference in energy would correspond

with a specific wavelength line in the atomic emission spectrum Larger the transition the electron makes,

the higher the energy the photon will have

Just for Thought How is energy related to

wavelength?

Summary of Quanta and Atomic Spectra Atoms must somehow absorb energy and then

give the energy off in the form of light Excited electrons in an atom return to lower

energy states Each element has a unique emission spectrum

Electron movements create the specific colors that we witness

Only certain energies are possible Electron energy levels are quantized!

Thus, the electron arrangement in every element is unique!

Just for Thought… Can we map the electrons by using

these energy relationships from the emission spectrum?

Neils Bohr and the Atomic Model The answer is YES! Neils Bohr was one of the first to see

some connection between the wavelengths an element emits and its atomic structure Related Planck’s idea of quantized

energies to Rutherford’s atomic model

Bohr and the Atomic Model Bohr discovered that as the electrons in the

hydrogen atoms were getting excited and then releasing energy, only four different color bands of visible light were being emitted: red, bluish-green, and two violet-colored lines If electrons were randomly situated, as

depicted in Rutherford’s atomic model, then they would be able to absorb and release energy of random colors of light

Bohr concluded that electrons were not randomly situated Instead, they are located in very specific

locations that we now call energy levels

Bohr model of the Hydrogen Atom

Niels Bohr

• Protons and neutrons compose the nucleus

• Electrons orbit the nucleus in certain well-defined ‘energy levels’

nucleus

Many Electron Atoms Recall that because each element has a

different electron configuration and a slightly different structure, the colors that are given off by each element are going to be different Thus, each element is going to have its

own distinct color when its electrons are excited (or its own atomic spectra)

Flame Test Lab!