Atomic Structure

1 / 43

# Atomic Structure - PowerPoint PPT Presentation

Atomic Structure. Chapter 7: Describe the properties of electromagnetic radiation . Understand the origin of light from excited atoms and its relationship to atomic structure. Describe the experimental evidence for wave-particle duality. Describe the basic ideas of quantum mechanics .

I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.

## PowerPoint Slideshow about 'Atomic Structure' - keona

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.

- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

### Atomic Structure

Chapter 7:

Describe the properties of electromagnetic radiation.

Understand the origin of light from excited atoms and its relationship to atomic structure.

Describe the experimental evidence for wave-particle duality.

Describe the basic ideas of quantum mechanics.

Define the three quantum numbers and their relationship to atomic structure.

• List forms of electromagnetic radiation:

_______________ ___________

_______________ ___________

• Maxwell Theory (1831-1879): describe all forms of radiation in terms of ________

________________________________.

• Einstein Theory (1879-1955): light has _______________________________.
Wave Properties

wavelength

Visible light

Frequency – hertz (s-1)

Speed = wavelength (m) x frequency (s-1)

c = l x v

What is the frequency of orange light, which has a wavelength of 625 nm?

Students should be familiar with conversion of units and conversion between l and v.

The Visible Spectrum of Light
• Long wavelength --> ______ frequency

_____ energy

• Short wavelength --> _____ frequency

_____ energy

Energy and Frequency
• Max Planck (1858-1947): the energy of a vibrating systems is proportional to the frequency of vibration.
• The proportionality constant

h = Planck’s constant

= 6.6260693 x 10-34 J s

E = h v

Radiation given off by a Heated Body
• Planck solved the “___________________”.
• Vibrations are _________ – only vibrations with specific frequencies are allowed.
• There is a distribution of vibrations in a object.
Quantization of Energy
• An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA.
• Energy of radiation is proportional to frequency.

Light with large l (small v) has a _____ E.

Light with a short l (large v) has a ____ E.

E = h v

Photoelectric Effect
• Experiment demonstrates the _______ _____________________________.

No e- observed until light of a certain minimum E is used.

Photoelectric Effect
• Classical theory said that E of ejected

electron should increase with increase

in light frequency—not observed!

• No e- observed until light of a certain

minimum E is used.

• If the frequency is above the minimum,

the number of e- ejected depends on

light intensity.

• Einstein explained the photoelectric effect: light consists of “__________” particles called PHOTONS – _______________.
• The energy of each photon is proportional to the ______________of radiation (Planck’s relation).
• The greater the intensity of light, the more photons are available to strike per unit of time.
Show that the energy of a mol of blue photons (l = 400 nm) is higher than the energy of a mol of red photons (l=685 nm)

~

= h c v

Using Planck’s Equation
• As frequency (v) increases, energy (E) __________.
• As wavelength (l) decreases, energy (E) _________.

E = h v

v = c/l

E = h v = h c

l

(wavenumber)

Students should be familiar with frequency, wavelength, and energy calculations.

Photosynthesis
• Chlorophylls absorb blue and red light and carotenoids absorb blue-green light, but green and yellow light are not effectively absorbed by photosynthetic pigments in plants; therefore, light of these colors is either reflected by leaves or passes through the leaves. This is why plants are green.
Spectrum of Excited Hydrogen Gas
• Excited atoms emit light of only certain wavelengths. –Evidence of ____________________.
• Line Emission Spectra of Excited Atoms.
• The wavelengths of emitted light depend on ______________________________.

1

l

1

22

1

n2

(

)

= R

Which Mathematical Expression represents the Regular Patterns of Emission?
• Johann Balmer (1825-1898) and Johannes Rydberg (1854-1919) developed an equation:
• Rydberg equation – to calculate the _________________

__________________

__________________.

• Rydberg constant = R

R = 1.0974 x 107 m-1

when n > 2

n = 3 , l =red line

n = 4 , l = green line,

Etc. Balmer Series

Atomic View of the Early 20th Century

An electron (e-) traveled about the nucleus in an orbit.

1. Any orbit should be possible and so is any energy.

2. But a charged particle moving in an electric field should emit energy.

End result should be matter self-destruction!

Bohr Model
• Niels Bohr (1885-1962) connected the observation of the spectra of excited atoms with the quantum ideas of Planck and Einstein.
• Based on Rutherford’s work – electrons are arranged in space outside the atom.
• Bohr model shows electrons moving in a circular orbit around the nucleus.
• Bohr postulated:

1.- An electron could occupy only __________ ___________or energy levels in which it is stable.

2.-The energy of the electron in the atom is ______________.

Atomic Spectra and Bohr
• n ___________ quantum number
• n is a _________________ having values of 1, 2, 3 and so on.
• The energy of attraction between oppositely charged bodies (negative electron and positive nuclear proton) has a negative value. The value becomes more negative as the bodies move closer together (Coulomb’s law).
• As the value of n increases, the energy becomes less negative, the distance of the electron from the nucleus increases.

Rh c

n 2

Potential energy of electron

in the nth level

= En = -

n = 2

2

E = -C (1/ 2

)

n = 1

2

E = -C (1/1

)

Atomic Spectra and Bohr
• Only orbits where n = integral number are permitted.

If e-’s are in quantized energy states, then ∆E of states can have only certain values. This explain sharp line spectra.

Ground State and Excited State
• Ground state: The state of an atom in which all electrons are in the ______________________.
• Excited state: The state of an atom in which at least one electron is ______________________ ____________________.
CC alculateDE for an e- of the H atom “falling” from high energy level (n = 2) to low energy level (n = 1).
Atomic Spectra and Bohr
• The amount of energy that must be absorbed by the atom so that an electron can move from the first to the second energy state is 3/4RhC or 984 kJ/mol of atoms – no more or less – energy levels in the H atom are quantized – only certain amounts of energy may be absorbed or emitted.
• When an electron “falls” from a level of higher n to one of lower n, ________ energy. The negative sign indicates energy is _________, 984 kJ must be _______ per mole of H atoms.
• The energy ________ is observed as ______ – This is the source of the lines observed in the emission spectrum of H atoms. – The basic explanation holds for the spectra of other elements.
Atomic Spectra and Bohr

1

1

• The origin of atomic spectra is the movement of _________ between quantized energy states.
• Electron is excited from a lower energy state to a higher one – Energy is ________.
• Electron moves from a higher energy state to a lower one – Energy is _________.

(

)

-

∆E = Efinal – Einitial = -R h c

n2final

n2initial

Electronic Transitions in an Excited H Atom
• If electrons move from energy states n >1 to the n =1 state – emission lines have energies in the UV region (Lyman series).
• If electrons move from energy states n >2 to the n =2 state – emission lines have energies in the VIS region (Balmer series).
• If electrons move from energy states n >3 to the n =3 state – emission lines have energies in the IR region.
Calculate the wavelength of the photon emitted if an electron in the H atom moves from n = 4 to n =2
Flaws in Bohr’s Theory
• Bohr’s model of the atom explained only the spectrum of H atoms and of other systems having one electron (such as He+).
• The idea that electrons are particles moving about the nucleus with a path of fixed radius, like that of the planets about the sun, is no longer valid.
Wave Mechanics

Louis de Broglie (1892-1987) proposed that all moving objects have _______ _________________(1924).

For light: (1) E = mc2

(2) E = h v = h c / l

Wave Mechanics –Calculate the Broglie Wavelength

Baseball (115 g) at 100 mph

e- with velocity = 1.9 x 108 cm/sec

It is possible to observe wave-like properties only for particles of extremely __________, such as protons, neutrons, and electrons.

l= h

m v

The Uncertainty Principle
• Erwin Schrödinger, 1887-1961 : developed ________________or ______________.
• Werner Heisenberg, 1901-1976 : The uncertainty principle – it is impossible to fix both the ______________ electron in an atom and its ________ with any degree of certainty.
• Max Born, 1882-1970 : if the energy of an electron in an atom is known with a small uncertainty, there will be large uncertainty in its position in the space about the atom's nucleus.
• We can assess only the likelihood, or probability, of finding an electron with a given energy within a given region of space.
Schrödinger's Wave Functions
• The behavior of the electron in the atom is best described as a standing wave – In a vibrating string, only certain vibrations can be observed = only certain wave functions are allowed for the electron in the atom.
• Each wave function () is associated with an allowed energy value, En, for the electron.
• Then, from 1 and 2, the energy of the electron is quantized – only certain values of energy.

Wave motion:wave length and nodes

4. In contrast to Bohr’s theory – quantization is imposed as a postulate.

Schrödinger's Wave Functions

5. The is related to the probability of finding the electron within a given region of space = _______________.

6. Energy is known precisely – position is given by a probability. The region of space in which an electron of a given energy is most probably located is called its _______________.

7. The solution to the Schrödinger's equation, for an electron, in a 3-D space, are 3 integer numbers = quantum numbers n, l, and ml. These numbers have only certain combination of values.

Quantum numbers
• n, Principal quantum number = 1, 2, 3, …

Determines the ________ of the electron. Also related to size of orbital.

En = - Z2h R / n2

Electrons with the same n value are in the same electron ______ or same electron _________.

• l, Angular Momentum quantum number = 0, 1, 2, 3, …, n-1

Determines the ______ at which electrons circulate about the nucleus. Related to orbital __________.

Electrons with the same l value are in the same _______ and have the same orbital _____ (______). All orbitals in the same subshell have the same ___________.

• ml, Magnetic quantum number = 0, ±1, ± 2, ± 3, …, ±l

Determines the _____________ of the orbital motion of the electron. (Clockwise or counterclockwise). Related to ___________ in space of the orbitals within a subshell, this gives the ___________ of orbitals in a subshell.

See Table 7.1 (p 319)

Quantum numbers and Orbitals

Number of subshells in a shell = n

Number of orbitals in a subshell = 2l + 1

Number of orbitals in a shell = n2

l =0 (s) ; l =1 (p) ; l =2 (d) ; l =3 (f)

Name of orbital = value of n and letter code for l

If n=1 ; l = n-1 = 0 ; ml = 0

Only 1 subshell (s); only 1 orbital (1s)

If n=2 ; l = 0, 1 ; ml = +1, 0, -1

There are 2 subshells (s and p)

4 orbitals (the 2s, and three 2p (3 orientations)

Orbitals
• Electron orbitals are probabilities – represented as ____________________.
Orbitals

surface density plot

• For the s orbital, the probability of finding an electron is the same at the same distance from the nucleus – the 1s orbital is ____________ in shape.
• Quantum mechanics – electron has wave properties – the maximum amplitude of the electron wave occurs at 0.053 nm from the nucleus.
• Bohr’s radius = 0.059 nm
Orbitals
• The p orbitals have 1 nodal surface – zero probability of finding an electron.
• Number of nodal surfaces = value of l
• There are three p orbitals in each p subshell: ml = +1, 0, -1
• Refer to orbitals according to the axes along which the lobes lie: px, py, pz
Orbitals
• The d five orbitals, l=2 have 2 nodal surfaces (may not be flat).
• What type of orbital is designated n = 4, l = 3, ml =-3?

a. 4s

b. 4p

c. 4d

d. 4f

e. none

Orbitals

Students should be familiar with definitions of quantum numbers and orbital types.

Practice
• Which of the following represent valid sets of quantum numbers?
• n=3, l=3, ml= +1
• n=5, l=1
• n=6, l=5, ml=1
• n=4, l=3, ml=-4
Remember
• Go over all the contents of your textbook.
• Practice with examples and with problems at the end of the chapter.
• Practice with OWL tutors.
• W ork on your assignment for Chapter 7.
• Practice with the quiz on the cd or online service.