V.Montgomery & R.Smith - 2.1 Atomic Structure complete OK

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Atomic Structure

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Classical Model

Democritus Dalton

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Democritus

Circa 400 BC

Greek philosopher

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Dalton’s Atomic Theory (1808)

1. All matter is made of tiny indivisible particles called atoms.

2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element.

3. Atoms of different elements can combine with one another in simple whole number ratios to form compounds.

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J.J. Thomson (1897)

Determined the charge to mass ratio for electrons

Applied electric and magnetic fields to cathode rays

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Rutherford’s Gold Foil

Experiment (1910)

Alpha particles (positively charged

helium ions) from a radioactive source was directed toward a very thin gold foil.

A fluorescent screen was placed behind the Au foil to detect the scattering of

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Rutherford’s Gold Foil

Experiment (Observations)

Most of the _-particles passed through

the foil.

Many of the _-particles deflected at

various angles.

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Rutherford’s Gold Foil

Experiment (Conclusions)

Rutherford concluded that most of the mass of an atom is concentrated in a core, called the atomic nucleus.

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Shortfalls of Rutherford’s

Model

Did not explain where the atom’s negatively charged electrons are located in the space surrounding its positively charged nucleus.

We know oppositely charged particles attract each other

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Bohr Model (1913)

Niels Bohr (1885-1962), Danish scientist working with Rutherford

Proposed that electrons must have enough energy to keep them in

constant motion around the nucleus

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Planetary Model

The planets are attracted to the sun by gravitational force, they move with

enough energy to remain in stable orbits around the sun.

Electrons have energy of motion that enables them to overcome the

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Think about satellites….

We launch a satellite into space with enough energy to orbit the earth

The amount of energy it is given, determines how high it will orbit

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Electronic Structure of Atom

Waves-particle duality Photoelectric effect

Planck’s constant Bohr model

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Radiant Energy

Radiation _ the emission of energy in

various forms

A.K.A. Electromagnetic Radiation

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Electromagnetic Radiation _ radiation that consists of wave-like electric and

magnetic fields in space, including light, microwaves, radio signals, and x-rays

Electromagnetic waves can travel

through empty space, at the speed of

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Waves

Waves transfer energy from one place to another

Think about the damage done by waves during strong hurricanes.

Think about placing a tennis ball in your bath tub, if you create waves at one it, that energy is transferred to the ball at the other = bobbing

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Wave Characteristics

Wavelength, _ (lambda) _ distance between

successive points

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Wave Characteristics

Frequency, _ (nu) _ the number of

complete wave cycles to pass a given point per unit of time; Cycles per

second

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Units for Frequency

1/s s-1

hertz, Hz

Because all electromagnetic waves travel at the speed of light, wavelength is determined by frequency

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Waves

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Waves

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Electromagnetic Spectrum

Radio & TV, microwaves, UV, infrared, visible light = all are examples of

electromagnetic radiation (and radiant energy)

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Electromagnetic Spectrum

1024 1020 1018 1016 1014 1012 1010 108 106

Gamma Xrays UV _Microwaves_{FM AM} IR

Frequency Hz

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Notes

Higher-frequency electromagnetic

waves have higher energy than lower-frequency electromagnetic waves

All forms of electromagnetic energy

interact with matter, and the ability of these different waves to penetrate

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What is your favorite radio

station?

Radio stations are identified by their frequency in MHz.

We know all electromagnetic

radiation(which includes radio waves) travel at the speed of light.

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Velocity of a Wave

Velocity of a wave (m/s) = wavelength (m) x frequency (1/s)

c = _

c= speed of light = 3.00x108 m/s

My favorite radio station is 105.9 Jamming Oldies!!!

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Wavelength of FM

c = _

 = 105.9MHz or 1.059x108Hz

 = c/ =3.00x108 m/s = 2.83m

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What does the electromagnetic

spectrum have to do with electrons?

It’s all related to energy – energy of

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States of Electrons

When current is passed through a gas at a low pressure, the potential energy (energy due to position) of some of the gas atoms increases.

Ground State: the lowest energy state of an atom

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Neon Signs

When an excited atom returns to its ground state it gives off the energy it gained in the form of electromagnetic radiation!

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White Light

White light is composed of all of the colors of the spectrum = ROY G BIV When white light is passed through a prism, the light is separated into a

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Line-emission Spectrum

When an electric current is passed

through a vacuum tube containing H2

gas at low pressure, and emission of a pinkish glow is observed.

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Hydrogen’s Emission

Spectrum

The pink light consisted of just a few specific frequencies, not the whole range of colors as with white light

Scientists had expected to see a continuous range of frequencies of electromagnetic

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Bohr’s Model of Hydrogen

Atom

Hydrogen did not produce a continuous spectrum

New model was needed:

 Electrons can circle the nucleus only in allowed

paths or orbits

 When an e- is in one of these orbits, the atom has

a fixed, definite energy

 e- and hydrogen atom are in its lowest energy

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Bohr Model Continued…



Orbits are separated by empty space,

where e- cannot exist



Energy of e- increases as it moves to

orbits farther and farther from the nucleus

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Bohr Model and Hydrogen

Spectrum

While in orbit, e- can neither gain or lose energy

But, e- can gain energy equal to the difference between higher and lower orbitals, and therefore move to the higher orbital (Absorption)

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Bohr’s Calculations

Based on the wavelengths of

hydrogen’s line-emission spectrum,

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Photoelectric Effect

An observed phenomenon, early 1900s

When light was shone on a metal, electrons were emitted from that metal

Light was known to be a form of energy,

capable of knocking loose an electron from a metal

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Photoelectric Effect pg. 93

Light strikes the surface of a metal (cathode), and e- are ejected.

These ejected e- move from the cathode to the anode, and current flows in the cell.

A minimum frequency of light is used. If the frequency is above the minimum and the

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Photoelectric Effect

Observed: For a given metal, no

electrons were emitted if the light’s frequency was below a certain

minimum, no matter how long the light was shone

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Explanation….

Max Planck studied the emission of light by hot objects

Proposed: objects emit energy in small, specific amounts = quanta

(Differs from wave theory which would say objects emit electromagnetic radiation continuously)

Quantum: is the minimum quantity of

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Planck’s Equation

E _radiation= Planck’s constant x frequency of radiation

E = h_

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

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Einstein Expands Planck’s

Theory

Theorized that electromagnetic

radiation had a dual wave-particle nature!

Behaves like waves and particles

Think of light as particles that each carry one quantum of energy =

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Photons

Photons: a particle of electromagnetic radiation having zero mass and carrying a quantum of energy

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Back to Photoelectric Effect

Einstein concluded:

 Electromagnetic radiation is absorbed by

matter only in whole numbers of photons

 In order for an e- to be ejected, the e-

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Example of Planck’s Equation

CD players use lasers that emit red light

with a _ of 685 nm. Calculate the

energy of one photon.

 Different metals require different minimum

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Answer

Ephoton = h

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

c = _

= (3.00x108 m/s)/(6.85x10-7m)

=4.37x10141/s

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Wave Nature of Electrons

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de Broglie’s Equation

A free e- of mass (m) moving with a velocity (v) should have an associated

wavelength: _ = h/mv

Linked particle properties (m and v)

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Example of de Broglie’s

Equation

Calculate the wavelength associated

with an e- of mass 9.109x10-28 g

traveling at 40.0% the speed of light.

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Answer

C=(3.00x108m/s)(.40)=1.2x108m/s

 = h/mv

 = (6.626 x 10-34 J•s) =6.06x10-12m

(9.11x10-31kg)(1.2x108m/s)

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Wave-Particle Duality

de Broglie’s experiments suggested that e- has wave-like properties.

Thomson’s experiments suggested that e- has particle-like properties

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Quantum mechanical model

SchrÖdinger

Heisenberg Pauli

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Where are the e- in the atom?

e- have a dual wave-particle nature

If e- act like waves and particles at the same time, where are they in the atom? First consider a theory by German

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Heisenberg’s Idea

e- are detected by their interactions with photons

Photons have about the same energy as

e-Any attempt to locate a specific e- with a photon knocks the e- off its course

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Heisenberg’s Uncertainty

Principle

Impossible to determine both the

position and the momentum of an e- in an atom simultaneously with great

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Schr

Ö

dinger’s Wave Equation

An equation that treated electrons in atoms as waves

Only waves of specific energies, and therefore frequencies, provided

solutions to the equation

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Schr

Ö

dinger’s Wave Equation

Solutions are known as wave functions Wave functions give ONLY the

probability of finding and e- at a given place around the nucleus

e- not in neat orbits, but exist in regions

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Schr

Ö

dinger’s Wave Equation

Here is the equation

Don’t memorize this or write it down

It is a differential equation, and we need calculus to solve it

-h (ә2 Ψ )+ (ә2Ψ )+( ә2Ψ ) +Vψ =Eψ

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Probability _ likelihood

Orbital _ wave function; region in space

where the probability of finding an electron is high

SchrÖdinger’s Wave Equation states

that orbitals have quantized energies But there are other characteristics to describe orbitals besides energy

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Quantum Numbers

Definition: specify the properties of atomic orbitals and the properties of electrons in orbitals

There are four quantum numbers The first three are results from

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Quantum Numbers (1)

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Quantum Numbers

Principal Quantum Number, n

 Values of n = 1,2,3,… _

 Positive integers only!

 Indicates the main energy level occupied

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Quantum Numbers

 Values of n = 1,2,3,… _

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Quantum Numbers

 Values of n = 1,2,3,… _

 Describes the energy level, orbital size

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Principle Quantum Number

n = 1 n=2

n=3 n=4 n=5 n=6

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Principle Quantum Number

More than one e- can have the same n value

These e- are said to be in the same e- shell

The total number of orbitals that exist

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Quantum Numbers (2)

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Quantum Numbers

Angular momentum quantum number,

l

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Quantum Numbers

l

 Values of l = n-1, 0

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Quantum Numbers

l

 Values of l = n-1, 0

 Describes the orbital shape

 Indicates the number of sublevel

(subshells)

(except for the 1st main energy level,

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Orbital Shapes

For a specific main energy level, the number of orbital shapes possible is equal to n.

Values of

_l

= n-1, 0

 Ex. Orbital which n=2, can have one of two

shapes corresponding to l = 0 or l=1

Depending on its value of l, an orbital is

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Orbital Shapes

Angular magnetic quantum number, l

If _l = 0, then the orbital is labeled s.

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Orbital Shapes

If l = 1, then the orbital is labeled p.

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Orbital Shapes

If l = 2, the orbital is labeled d.

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Orbital Shapes

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Energy Level and Orbitals

n=1, only s orbitals n=2, s and p orbitals

n=3, s, p, and d orbitals n=4, s,p,d and f orbitals

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Atomic Orbitals

Atomic Orbitals are designated by the principal quantum number followed by letter of their subshell

 Ex. 1s = s orbital in 1st main energy level

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Quantum Numbers (3)

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Quantum Numbers

Magnetic Quantum Number, m_l

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Quantum Numbers

Magnetic Quantum Number, m_l

 Values of m_l = +l…0…-l

 Describes the orientation of the

orbital

Atomic orbitals can have the same shape

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Magnetic Quantum Number

s orbitals are spherical, only one orientation, so m=0

p orbitals, 3-D orientation, so m= -1, 0 or 1 (x, y, z)

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Quantum Numbers (4)

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Quantum Numbers

Electron Spin Quantum Number,m_s

 Values of m_s = +1/2 or –1/2

 e- spin in only 1 or 2 directions

 A single orbital can hold a maximum of 2

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Electron Configurations

Electron Configurations: arragenment of e- in an atom

There is a distinct electron configuration for each atom

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Pauli Exclusion Principle

No 2 e- in an atom can have the same

set of four quantum numbers (n, l, m_l_,

m_s). Therefore, no atomic orbital can

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Aufbau Principle

Aufbau Principle: an e- occupies the

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Hund’s Rule

Hund’s Rule: orbitals of equal energy

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Electron Configuration

The total of the superscripts must equal the atomic number (number of

electrons) of that atom.

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Differentiating Electron

The differentiating electron is the electron that is added which makes the configuration different from that of the preceding element. The “last” electron.

H 1s1

He 1s2

Li 1s2, 2s1

Be 1s2, 2s2

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Orbital Diagrams

These diagrams are based on the electron configuration.

In orbital diagrams:

 Each orbital (the space in an atom that will

hold a pair of electrons) is shown.

 The opposite spins of the electron pair is

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Orbital Diagram Rules

1. Represent each electron by an arrow

2. The direction of the arrow represents the

electron spin

3. Draw an up arrow to show the first electron

in each orbital.

4. Hund’s Rule: Distribute the electrons among

the orbitals within sublevels so as to give the most unshared pairs.

Put one electron in each orbital of a sublevel before the second electron appears.

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Orbital Diagram Examples

H __

1s

Li _ __

1s 2s

B _ _ _ __ __

1s 2s 2p

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Dot Diagram of Valence

Electrons

When two atom collide, and a reaction takes place, only the outer electrons interact.

These outer electrons are referred to as the valence electrons.

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Rules for Dot Diagrams

:Xy:

. .

_{. .}

S sublevel

electrons P_x orbital

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Rules for Dot Diagrams

Remember: the maximum number of valence electrons is 8.

Only s and p sublevel electrons will ever be valence electrons.

Put the dots that represent the s and p electrons around the symbol.

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Examples of Dot Diagrams

H

He

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Examples of Dot Diagrams

C

N

O

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Summary

Both dot diagrams and orbital diagrams will be use full to use when we begin our study of

atomic bonding.

We have been dealing with valence electrons since our initial studies of the ions.

The number of valence electrons can be determined by reading the column number.