copyright©2000 by houghton mifflin company. all rights reserved. 1 ch 7 atomic structure

Copyright©2000 by Houghton Mifflin Company. All rights reserved.

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Ch 7 Atomic Structure


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Rutherford Model

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Figure 7.2Classification of Electromagnetic

Radiation


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Electromagnetic Radiation

Radiant energy that exhibits wavelength-like behavior and travels through space at the speed of light in a vacuum.


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Light is a wave

Wavelength Distance between 2

similar points (meters)

Frequency Number of waves in a second (frequency (s1) or hertz (Hz))

Speed v (m/s)

Light: energy that travels like a wave through spaceWave properties:


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Wavelength and Frequency


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Light is a wave

All light travels the same speed:

high , has short low , has long

high = high energy


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Light is a wave

Created by movement of electric charge

An electric field and magnetic field perpendicular to each

Self-propagating

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Wavelength and frequency can be interconverted.

= c/(C =

= frequency (s1) or hertz (Hz)

= wavelength (m)

c = speed of light (m s1)

(2.9979 x 108 m/s)


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Figure 7.2Classification of Electromagnetic

Radiation


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Matter is not what it appears to be.

Before 1900:

Matter particle

Light a wave

Max Planck:

not all energies were emitted from objects heated incandescence


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Planck’s Constant

E = change in energy, in Jh = Planck’s constant, 6.626 1034 J s = frequency, in s1

= wavelength, in m

Energy gained or lost only in Energy gained or lost only in whole number whole number multiples. Transfer of energy is quantized: multiples. Transfer of energy is quantized: occur in discrete units, called quanta.occur in discrete units, called quanta.

E = nhnhcn = 1, 2,3,..


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Light is a particle

Einstein: theorizes that light made of photons.

Gets Nobel Prize QuickTime™ and aTIFF (Uncompressed) decompressorare needed to see this picture.


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Photoelectric Effect

ExperimentExperiment: light of different frequency shone on metal

Results:Results: e- ejected only at minimum

No e- ejected if too low, EVEN IF light intensity is increased.

WHY?WHY? Why some frequency and not others?


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- High photon = high E photon

- One photon hits one electron - If photon E not = to e- E nothing happens (even if bright)

-Giant example

Light made of photons






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Photoelectric Effect




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Light made of photons

Einstein: electromagnetic radiation is

quantized:

Ephoton = h = hc


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Energy and Mass

Einstein’s special theory of relativity: (1905) Energy has mass:

E = mc2

Or

m=E/c2

E = energym = massc = speed of light


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Energy and Mass

(Hence the (Hence the dualdual nature of light.) nature of light.)

Does a photon have mass?

for a photon with wavelength m = E = hc/ mh

c2 c2 c Ephoton= hc/


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Figure 7.4Dual Nature of light


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If light can be a particle can a particle (e-) be a wave?

m = h /m = h /vh / m h / m v

= wavelength, in m

h = Planck’s constant, 6.626 1034 J s = kg m2 s1

m = mass, in kg

v = velocity in m/s

Louis de Broglie’s Louis de Broglie’s Equation 1920Equation 1920


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Wave interference




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Water wave interference


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Interference in water waves








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Water interference




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Interference patterns




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Light interference


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Interference Pattern


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Electron interference

Diffraction patterns caused by interferenceX-rays passing through NaCl crystal are diffracted.Electrons passing through NaCl crystal are diffracted x-rays


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Diffraction using NaClcrystal


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So, Debroglie was right: all matter show both wave like and

particle like behavior


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How does all this stuff relate to the e- and the atom?

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Bohr model: electrons are at set distances from the nucleus (energy levels)


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How do we know this?:we can use Einstein’s ideas to

explain bright-line spectra Observe the light coming from the hydrogen emission tube.

“Excited” atoms only emit certain frequencies (colors) of light Why not all frequencies of light? Look at hydrogen emission. Each line is one frequency of light.


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Figure 7.6A Continuous Spectrum (a) and A Hydrogen Line Spectrum (b)

Absorption, emission, and energy

Ground State: electrons in lowest energy state.

Excited State: when one electron absorbsabsorbs one photon and jumps to higher energy level.

When electron falls back to G.S. it

emits emits one photon.

electron can only absorb photons or emit photons of just the right energy because levels are fixed.

Absorption

photon

Emission

photon


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The emitted photons are seen as light of specific frequency (i.e. colors).

What color emitted if electron could go anyway?

A: the contiuous spectrum


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What is the energy difference between levels?

Must be equal to the energy of the photon emitted. Energy levels are quantized:

∆E = h= hc


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Figure 7.7A Change between Two Discrete

Energy Levels


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Figure 7.8Electronic Transitions in the Bohr Model for the Hydrogen Atom


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The Bohr Model

E = energy of the levels in the H-atom

z = nuclear charge (for H, z = 1)

n = an integer

E = 2.178 10 J (18 2− × − z n/ )2

The electron in a hydrogen atom moves around the The electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits. The nucleus only in certain allowed circular orbits. The energy of each level is given by this equation:energy of each level is given by this equation:


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The Bohr Model

Ground State: The lowest possible energy state for an atom (n = 1).


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Energy Changes in the Hydrogen Atom

E = Efinal state Einitial state

= hcEΔ


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Standing waves in spring: - exist only at specific and ( .5, 1, 1.5, 2.0, 2.5, ect.) - are quantized.

Electron waves: - exist only at certain

and (and energy) - only form certain

distances from nucleus. - are quantized


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Figure 7.10The Hydrogen Electron Visualized as a Standing Wave Around the Nucleus


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Quantum Mechanics

Schrodinger’s equations: Based on the wave properties of the atom

= wave function

= mathematical operator

E = total energy of the atom

A specific wave function is often called an orbital.

$H E =

$H


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What is the orbital for H when n=1?

1s orbital

Orbitals are not the Bohr orbits.

Where is the electron in the orbital?


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Heisenberg

There is a limit to how precisely we can both the position and momentum of a an electron at a given timeQuickTime™ and a

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Heisenberg Uncertainty Principle

x mvh

⋅ ( ) ≥ 4π

• x = position• mv = momentum• h = Planck’s constant• The more accurately we know a

particle’s position, the less accurately we can know its momentum.


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What is an orbital?

If we don’t know the motion of an electron, what is an orbital?

square of the wave function gives the probability of finding an electron at a given position.

--> (a) probability distrib.

for H 1s orbital


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Figure 7.12Radial Probability Distribution

(the probability distribution in each spherical shell.)


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Bohr radius

Turns out that for H 1s orbital, the max radial probability is 5.29x10-2 nm,

= Bohr’s innermost “orbit”.


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How big is the 1s?

Probab. Decreases with radius, but never goes to zero.

Size definition: Size of the orbital is the radius of the sphere that encloses 90% of the electron’s probability.


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Quantum Numbers (QN)Schrodinger equation has many wave function solutions.

Each are described by quantum numbers1. Principal QN (n = 1, 2, 3, . . .) - related to size and energy of the orbital.

(this gives the”rings” or “shells)

2. Angular Momentum QN (l = 0 to n 1) - relates to the shape of the orbital. (ex. s p d f also called subshells)

1. Magnetic QN (ml = l to l ) - relates to orientation of the orbital in space relative to other orbitals. Gives you the number of each type of orbital.

(ex.: px py pz)

4. Electron Spin QN (ms = +1/2, 1/2) - relates to the spin states of the electrons.

(see pg 310, table 7.2)


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Orbitals when n=1 (principal quantum #):

(l = 0 to n-1)

(ml = l to - l ): l =0 ml=0 1s orbital only.

Only 1 or 2 electrons are described by an orbital.

Total electrons at n=1? : .


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When n= 2,(l = 0 to n-1) ml = l to - l ):

l = 0 and l = 1

When l = 0, ml = 0 2s orbitalWhen l = 1 ml = -1, 0, 1 giving three 2p orbitals:

2px 2py 2pz

Total orbitals:Total electrons:


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The P orbitals (energy level 2 and up)




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When n=3 (l = 0 to n-1) ml = l to - l ):

l = 0 , 1 , 2

When l = 0, ml = 0 giving one 3s orbital

When l = 1 ml = -1, 0, 1 giving three 3p orbitals:

3px 3py 3pz

When l = 2 ml = -2,-1, 0, 1,2 giving 5 d orbitals

Total orbitals level 3:Total electrons level 3:


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When n=4

l = 0 , 1 , 2, 3 when l = 0 , 1 , 2 : 4s 4px 4py 4pz 4 d’s (5 of them)

when l =3 4f orbitals (7 of them)Total orbitals: Total electrons:


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electron spin quantum number:ms = +1/2 or - 1/2


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Pauli Exclusion Principle

In a given atom, no two electrons can have the same set of four quantum numbers (n, l, ml, ms).

Therefore, an orbital can hold only two electrons, and they must have opposite spins.


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Quantum Model

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Hydrogen orbitals are degenerate

All H orbitals with the same n have the same energy.


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Polyelectronic Atoms

• How does it work after Hydrogen?

• Shielding (e- repel, feel less attraction to +)

• Hydrogen orbitals: degenerate

• Hydrogenlike orbitals: NOT degenerate!

• Ens < Enp < End < Enf …etc.

• Why?


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Figure 7.20A Comparison of the Radial Probability Distributions of the 2s and 2p Orbitals

2p appears closer to nucleus? Less energy? No, look at small 2s hump. “2s penetrates to the nucleus” penetration effect


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Figure 7.21The Radial Probability Distribution for the 3s, 3p, and 3d Orbitalsso,

E3s<E3p<E3d


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History of the Periodic Table7.10

Dobereiner: triads

Newlands: octaves

Meyer / Mendeleev: arrangements by atomic masses.

Theory → prediction


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Figure 7.23Mendeleev’s Early Periodic Table,

Published in 1872

Prediction: Ga, Ge (see table 7.3, pg318 to see how

cool Mendeleev was)


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Aufbau Principle

As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals.

“ an electron occupies the lowest energy orbital that can receive it”


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Hund’s Rule

The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals.

“in the p, d, f, orbitals, spread out before you pair up”


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The Electron Configurations in the Type of Orbital Occupied Last for the First 18

Elements

Note: elements in same group have number valence electrons (valence: The electrons in the outermost principle quantum level of The electrons in the outermost principle quantum level of

an atom.an atom. ((Core electron: other than valence)

Ex:Cl 1s22s22p63s23p5 or [Ne] 3s23p5

# valence = ?


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Figure 7.25Electron Configurations for Potassium Through Krypton

Why doe the 4s fill before the 3d? Penetration effect

Notice Cr, Cu columns.


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Figure 7.26The Orbitals Being Filled for Elements in Various Parts of the Periodic Table

After lathanum [Xe] 6s25d1 , go to lathanide series, fill the 4fs (fig 7.27: note anomalies)

After Actinium [Rn]7s26d1, fill actinide series with 5fs.


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Figure 7.27 The Periodic Table With Atomic Symbols, Atomic Numbers, and Partial Electron

Configurations


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Figure 7.36 Special Names for Groups in the Periodic Table

Main-group elements or

representative elements:

1A 2A 3A 4A 5A 6A

7A 8A or 1,2, 13-18

M-g e: each group has same #valence


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Broad Periodic Table Classifications

Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O)

Transition Elements: filling d orbitals (Fe, Co, Ni)

Lanthanide and Actinide Series (inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es)


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Figure 7.30The Positions of the Elements

Considered in Sample Exercise 7.7


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Ionization Energy

The quantity of energy required to remove one electron from the gaseous atom or ion.

X(g) → X+(g) + e-


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Periodic Trends

First ionization energy: increases from left to right across a period (increase +, no shielding)decreases going down a group.

( increase n, more shielding)


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Figure 7.31The Values of First Ionization Energy for the

Elements in the First Six Periods


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Trends in Ionization Energies for the Representative Elements

Who has the highest IE? Who has the lowest? Do metals or nonmetals have higher IE? What does IE tell us about metal reactivity


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Periodic Trends: IE

Al(g) --> Al+ + e- I1= 580 kJ/mol

Al+(g) --> Al2+ + e- I2= 1815 kJ/mol

Al2+(g) --> Al3+ + e- I3= 2740 kJ/mol

Al3+(g) --> Al4+ + e- I4= 11,600 kJ/mol

Why the differences? Indicates something about the electron structure.


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Practice: 1s22s22p63s1

1s22s22p6

1s22s22p63s2

Which atom has the largest first I.E.?

Which one has the smallest second I.E.?

Explain.


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More Practice

IE increases across the period.

Check IE of P and S on pg 329. Explain the anomaly


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Electron Affinity

The energy change associated with the addition of an electron to a gaseous atom or ion.

X(g) + e X(g)If change is exothermic, then E.A. is

negative.


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Figure 7.33The Electronic Affinity Values for Atoms Among the First 20 Elements that Form

Stable, Isolated X- Ions


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Electron Affinity

Left to right mostly more neg kJ, more energy released. (note the missing elements: why C, but not N? write o.d. for both and t.t.y.n)Down a group, usually less neg kJ, less energy released. (see table 7.7. Notice anomaly. T.t.y.n.)


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E.A.

Who has more neg EA, metals or nonmetals?

What does this say about the reactivity of nonmetals?


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Radius trend

Data usually from distance between nuclei in a compound.


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Radius trend

Data usually from distance between nuclei in a compound.


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Figure 7.35Atomic Radii for Selected Atoms

Period trend?Group trend?

Why?


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Periodic Trends

Atomic Radii:decrease going from left to right across a period: increase + draws in valence.

increase going down a group: increase in orbital sizes with n.


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Alkali Metals Trends look at pg 335.

Note trends down the group in:

1. IE, radius

2. Density. Why?

3. mp/bp


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Alkali metals

Reaction with water. Write the balanced equation for Na(s) with water.

List alkali from most to least reactive (think I.E.) Cs > Rb > K > Na > LiBut in water,Li > K > Na , even though K loses electrons the easiest.

why? Hydration energy (see table 7.9) Li is small, higher charge density, better at attracting water.But there is more:


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Alkali metals

What we observe when they react with water:K > Na > Li

Look at mp. K and Na melt, increasing reaction rate.


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Waves

Waves have 3 primary characteristics:

1. Wavelength: distance between two peaks in a wave.

2. Frequency: number of waves per second that pass a given point in space.

3. Speed: speed of light is 3.00 108 m/s.


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Figure 7.1 The Nature of Waves

copyright©2000 by houghton mifflin company. all rights reserved. 1 ch 7 atomic structure

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