Chapter 71 Atomic Structure Chapter 7. 2 Electromagnetic Radiation -Visible light is a small portion of the electromagnetic spectrum.

Chapter 7 1

Atomic StructureAtomic Structure

Chapter 7Chapter 7

Chapter 7 2

Electromagnetic RadiationElectromagnetic Radiation

- Visible light is a small portion of the electromagnetic spectrum

Chapter 7 3

Frequency (v, nu – The number of times per second that one complete wavelength passes a given point.

Wavelength (lambda) – The distance between identical points on successive waves.

v = c

c = speed of light, 2.997 x 108 m/s


Chapter 7 4

- When talking about atomic structure, a special type of wave is important:

Standing Wave: A special type of wave with two or more stationary point with no amplitude.


Chapter 7 5

- We can also say that light energy is quantized- This is used to explain the light given-off by hot

objects. - Max Planck theorized that energy released or

absorbed by an atom is in the form of “chunks” of light (quanta).

E = h vh = planck’s constant, 6.63 x 10-34J/s

- Energy must be in packets of (hv), 2(hv), 3(hv), etc.

Planck, Einstein, Energy and PhotonsPlanck, Einstein, Energy and Photons

Planck’s EquationPlanck’s Equation

Chapter 7 6


The Photoelectric EffectThe Photoelectric Effect

Chapter 7 7

The Photoelectric EffectThe Photoelectric Effect- The photoelectric effect provides evidence for the

particle nature of light.- It also provides evidence for quantization.- If light shines on the surface of a metal, there is a

point at which electrons are ejected from the metal.- The electrons will only be ejected once the

threshold frequency is reached .- Below the threshold frequency, no electrons are

ejected.- Above the threshold frequency, the number of

electrons ejected depend on the intensity of the light.


Chapter 7 8

The Photoelectric EffectThe Photoelectric Effect

- Einstein assumed that light traveled in energy packets called photons.

- The energy of one photon, E = h.


Chapter 7 9

Bohr’s Model of the Hydrogen Atom Bohr’s Model of the Hydrogen Atom

Line SpectraLine Spectra

Chapter 7 10



Chapter 7 11



Line spectra can be “explained” by the following equation:

- this is called the Rydberg equation for hydrogen.

21

22

18 1110179.2

1

nnx

Chapter 7 12


Bohr’s ModelBohr’s Model- Assumed that a single electron moves around the

nucleus in a circular orbit. - The energy of a given electron is assumed to be

restricted to a certain value which corresponds to a given orbit.

k = 2.179 x 10-18J z = atomic numbern = integer for the orbit

2

2

n

kzE

Chapter 7 13


Bohr’s ModelBohr’s Model- Assumed that a single electron moves around the

nucleus in a circular orbit. - The energy of a given electron is assumed to be

restricted to a certain value which corresponds to a given orbit.

n = integer for the orbit ao = 0.0529 angstroms z = atomic number

z

anradius o

2

Chapter 7 14


Bohr’s Model – Important FeaturesBohr’s Model – Important Features- Quantitized energy and angular momentum- The first orbit in the Bohr model has n = 1 and is

closest to the nucleus.- The furthest orbit in the Bohr model has n close to

infinity and corresponds to zero energy.- Electrons in the Bohr model can only move between

orbits by absorbing and emitting energy in quanta (h).

Chapter 7 15


Bohr’s Model – Line SpectraBohr’s Model – Line Spectra

Ground State – When an electron is in its lowest energy orbit.

Excited State – When an electron gains energy from an outside source and moves to a higher energy orbit.

Chapter 7 16



12)( EElightE

Chapter 7 17



21

2

22

2

12)(

n

kz

n

kz

EElightE

Chapter 7 18



21

22

2

21

2

22

2

12

11

)(

nnkz

n

kz

n

kz

EElightE

Chapter 7 19



21

22

18

21

22

2

21

2

22

2

12

1110179.2

11

)(

nnx

nnkz

n

kz

n

kz

EElightE

Chapter 7 20


Bohr’s ModelBohr’s Model- Since the energy states are quantized, the light emitted

from excited atoms must be quantized and appear as line spectra.

Chapter 7 21

Quantum Mechanical View of the AtomQuantum Mechanical View of the Atom- DeBroglie proposed that there is a wave/particle

duality.- Knowing that light has a particle nature, it seems

reasonable to assume that matter has a wave nature.- DeBroglie proposed the following equation to describe

the relationship:

- The momentum, mv, is a particle property, where as is a wave property.

mvh

Chapter 7 22

The Uncertainty PrincipleThe Uncertainty PrincipleHeisenberg’s Uncertainty Principle - on the mass scale of

atomic particles, we cannot determine exactly the position, speed, and direction of motion simultaneously.

- For electrons, we cannot determine their momentum and position simultaneously.

Quantum Mechanical View of the AtomQuantum Mechanical View of the Atom

Chapter 7 23

- These theories (wave/particle duality and the uncertainty principle) mean that the Bohr model needs to be refined.

Quantum Mechanics

Quantum Mechanical View of the AtomQuantum Mechanical View of the Atom

Chapter 7 24

- The path of an electron can no longer be described exactly, now we use the wavefunction().

Wavefunction () – A mathematical expression to describe the shape and energy of an electron in an orbit.

- The probability of finding an electron at a point in space is determined by taking the square of the wavefunction:

Probability density =

Quantum MechanicsQuantum MechanicsSchrödinger’s ModelSchrödinger’s Model

Chapter 7 25

Quantum Mechanics

- The use of wavefunctions generates four quantum

numbers.

Principal Quantum Number (n)

Angular Momentum Quantum Number (l)

Magnetic Quantum Number (ml)

Spin Quantum Number (ms)

Quantum NumbersQuantum Numbers

Chapter 7 26

Quantum Mechanics

Principal Quantum Number (n) - This is the same as Bohr’s n- Allowed values: 1, 2, 3, 4, … (integers)- The energy of an orbital increases as n increases- A shell contains orbitals with the same value of n


Chapter 7 27

Quantum Mechanics

Angular Momentum Quantum Number (l) - Allowed values: 0, 1, 2, 3, 4, . , (n – 1) (integers)- Each l represents an orbital type

l orbital

0 s

1 p

2 d

3 f


Chapter 7 28

Quantum Mechanics

Angular Momentum Quantum Number (l) - Allowed values: 0, 1, 2, 3, 4, . , (n – 1) (integers)- Each l represents an orbital type- Within a given value of n, types of orbitals have slightly

different energy

s < p < d < f


Chapter 7 29

Quantum MechanicsQuantum Mechanics

Magnetic Quantum Number (ml).

- This quantum number depends on l. - Allowed values: -l +l by integers.-Magnetic quantum number describes the orientation of the

orbital in space.

l Orbital ml

0 s 0

1 p -1, 0, +1

2 d -2, -1, 0, +1, +2


Chapter 7 30


Magnetic Quantum Number (ml).

- This quantum number depends on l. - Allowed values: -l +l by integers.- Magnetic quantum number describes the orientation of the

orbital in space.- A subshell is a group of orbitals with the same value of n

and l.


Chapter 7 31


Spin Quantum Number (ms)

- Allowed values: -½ +½.

- Electrons behave as if they are spinning about their own

axis.- This spin can be either clockwise or counter clockwise.


Chapter 7 32

Quantum MechanicsQuantum MechanicsQuantum NumbersQuantum Numbers

Chapter 7 33

Representation of OrbitalsRepresentation of Orbitals

The The ss Orbitals Orbitals- All s-orbitals are spherical.- As n increases, the s-orbitals get larger.- As n increases, the number of nodes increase.- A node is a region in space where the probability of

finding an electron is zero.

Chapter 7 34


The The ss Orbitals Orbitals

Chapter 7 35


The The pp Orbitals Orbitals- There are three p-orbitals px, py, and pz. (The letters

correspond to allowed values of ml of -1, 0, and +1.)

- The orbitals are dumbbell shaped.

Chapter 7 36


The The pp Orbitals Orbitals

Chapter 7 37


The The dd and and ff Orbitals Orbitals- There are 5 d- and 7 f-orbitals. - Four of the d-orbitals have four lobes each.- One d-orbital has two lobes and a collar.

Chapter 7 38


The The dd and and ff Orbitals Orbitals

Chapter 7 39

32, 34, 42

Homework ProblemsHomework Problems

Chapter 71 Atomic Structure Chapter 7. 2 Electromagnetic Radiation -Visible light is a small portion of the electromagnetic spectrum.

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