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Atomic Structure and Periodicity Advanced Chemistry Ms. Grobsky.

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1 Atomic Structure and Periodicity Advanced Chemistry Ms. Grobsky

2 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”

3 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!

4 The Planetary Model is Doomed! Electron crashes into the nucleus!? Since this does not happen, the Rutherford model could not be accepted!

5 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

6 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 10 8 m/s, which is symbolized by the letter “c”

7 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

8 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 10 8 m/s  Speed of light in a vacuum is a constant c = λ· ν

9 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

10 Relationship of EM Wave Properties c = λ· ν

11 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???

12 Max Planck’s Obituary  Take 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

13 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

14 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 = x J· s

15 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

16 Practice, Practice, Practice  Calculate the energies of one photon of UV (λ = 1 x m), visible (λ = 5 x m), and IR (λ = 1 x m).

17 And then there was a problem… In the early 20 th 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

18 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

19 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  large ) Vary wavelength, fixed amplitude What if we try this ? electrons emitted ? No Yes, with low KE Yes, with high KE Increase energy by increasing amplitude “Classical” Method electrons emitted ? No

20 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

21 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 CollisionAfter Collision

22 The Dual Nature of Light A “Waveicle”

23 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 dealing with a single photon, or a few In this case, you cannot use the wave description of light It doesn’t work!

24 The Dualism of Light  Dualism is not such a strange concept  Consider the following picturepicture  Are the swirls moving, or not, or both?

25 But How is This Related to the Atom?

26 Light and the Dilemma of Atomic Spectral Lines  Experiments 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

27 Seeing Atomic Spectral Lines  Use your diffraction grating to observe the atomic spectra of:  Hydrogen  Oxygen  Neon

28 hydrogen (H) mercury (Hg) neon (Ne)

29 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

30 Just a Thought…  But if electrons are going from high-energy state to a low- energy state, where is all this extra energy going?

31 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

32 Just for Thought  How is energy related to wavelength?

33 Summary of Quanta and Atomic Spectra

34 Just for Thought…  Can we map the electrons by using these energy relationships from the emission spectrum?

35 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

36 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

37 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

38 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!

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