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Atomic Structure

- From Indivisible to Quantum Mechanical Model of

the Atom

Classical Model

- Democritus
- Dalton
- Thomson
- Rutherford

Democritus

- Circa 400 BC
- Greek philosopher
- Suggested that all matter is composed of tiny,

indivisible particles, called atoms

Daltons Atomic Theory (1808)

- All matter is made of tiny indivisible particles

called atoms. - Atoms of the same element are identical. The

atoms of any one element are different from those

of any other element. - Atoms of different elements can combine with one

another in simple whole number ratios to form

compounds. - Chemical reactions occur when atoms are

separated, joined, or rearrangedhowever, atoms

of one element are not changed into atoms of

another by a chemical reaction.

J.J. Thomson (1897)

- Determined the charge to mass ratio for electrons
- Applied electric and magnetic fields to cathode

rays - Plum pudding model of the atom

Rutherfords 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 alpha (?)

particles.

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Rutherfords Gold Foil Experiment (Observations)

- Most of the ?-particles passed through the foil.
- Many of the ?-particles deflected at various

angles. - Surprisingly, a few particles were deflected back

from the Au foil.

Rutherfords Gold Foil Experiment (Conclusions)

- Rutherford concluded that most of the mass of an

atom is concentrated in a core, called the atomic

nucleus. - The nucleus is positively charged.
- Most of the volume of the atom is empty space.

Shortfalls of Rutherfords Model

- Did not explain where the atoms negatively

charged electrons are located in the space

surrounding its positively charged nucleus. - We know oppositely charged particles attract each

other - What prevents the negative electrons from being

drawn into the positive nucleus?

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 - Analogous to the motion of the planets orbiting

the sun

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 attraction for the positive

nucleus

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 - We use energy from a rocket to boost our

satellite, what energy do we give electrons to

boost them?

Electronic Structure of Atom

- Waves-particle duality
- Photoelectric effect
- Plancks constant
- Bohr model
- de Broglie equation

Radiant Energy

- Radiation ? the emission of energy in various

forms - A.K.A. Electromagnetic Radiation
- Radiant Energy travels in the form of waves that

have both electrical and magnetic impulses

- 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 light (c3.00x108m/s) or

about 300million m/s!!!

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 - Electromagnetic waves have the same

characteristics as other waves

Wave Characteristics

Wavelength, ? (lambda) ? distance between

successive points

2m

10m

Wave Characteristics

- Frequency, ? (nu) ? the number of complete wave

cycles to pass a given point per unit of time

Cycles per second

t5

t0

t0

t5

Units for Frequency

- 1/s
- s-1
- hertz, Hz
- Because all electromagnetic waves travel at the

speed of light, wavelength is determined by

frequency - Low frequency long wavelengths
- High frequency short wavelengths

Waves

- Amplitude ? maximum height of a wave

Waves

- Node ? points of zero amplitude

Electromagnetic Spectrum

- Radio TV, microwaves, UV, infrared, visible

light all are examples of electromagnetic

radiation (and radiant energy) - Electromagnetic spectrum entire range of

electromagnetic radiation

Electromagnetic Spectrum

Frequency Hz

1024 1020

1018 1016 1014

1012 1010 108 106

Gamma Xrays UV

Microwaves FM AM

IR

10-16 10-9 10-8 10-6

10-3 100 102

Wavelength m

Visible Light

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 matter is a measure of the energy of

the waves

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. - What is the wavelength of your favorite station?

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!!! - What is the wavelength of this FM station?

Wavelength of FM

- c ??
- c speed of light 3.00x108 m/s
- ? 105.9MHz or 1.059x108Hz
- ? c/? 3.00x108 m/s 2.83m
- 1.059x1081/s

What does the electromagnetic spectrum have to do

with electrons?

- Its all related to energy energy of motion(of

electrons) and energy of light

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
- Excited State a state in which the atom has a

higher potential energy than it had in its ground

state

Neon Signs

- When an excited atom returns to its ground state

it gives off the energy it gained in the form of

electromagnetic radiation! - The glow of neon signs,is an example of this

process

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 spectrum, of all the

colors - What are rainbows?

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. - What do you think happens when that pink glow is

passed through a prism?

Hydrogens 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 radiation,

because the hydrogen atoms were excited by

whatever amount of energy was added to them. - Lead to a new theory of the atom

Bohrs 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

state when it is in the orbit closest to the

nucleus

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 - (Similar to a person climbing a ladder)

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) - When e- falls from higher state to lower state,

energy is emitted (Emission)

Bohrs Calculations

- Based on the wavelengths of hydrogens

line-emission spectrum, Bohr calculated the

energies that an e- would have in the allowed

energy levels for the hydrogen atom

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 - Therefore, light of any frequency could supply

enough energy to eject an electron.

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 intensity

of the light is increased, more e- are ejected.

Photoelectric Effect

- Observed For a given metal, no electrons were

emitted if the lights frequency was below a

certain minimum, no matter how long the light was

shone - Why does the light have to be of a minimum

frequency?

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 energy that

can be lost or gained by an atom.

Plancks Equation

- E radiation Plancks constant x frequency of

radiation - E h?
- h Plancks constant 6.626 x 10-34 Js
- When an object emits radiation, there must be a

minimum quantity of energy that can be emitted at

any given time.

Einstein Expands Plancks 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 photons

Photons

- Photons a particle of electromagnetic radiation

having zero mass and carrying a quantum of energy - Ephoton h?

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- must be

struck by a single photon with minimum frequency

Example of Plancks 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

frequencies to exhibit photoelectric effect

Answer

- Ephoton h?
- h Plancks constant 6.626 x 10-34 Js
- c ??
- c speed of light 3.00x108 m/s
- ? (3.00x108 m/s)/(6.85x10-7m)
- ?4.37x10141/s
- Ephoton (6.626 x 10-34 Js)(4.37x10141/s)
- Ephoton 2.90 x 10-19J

Wave Nature of Electrons

- We know electrons behave as particles
- In 1925, Louis de Broglie suggested that

electrons might also display wave properties

de Broglies 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) with a wave

property (?)

Example of de Broglies Equation

- Calculate the wavelength associated with an e- of

mass 9.109x10-28 g traveling at 40.0 the speed

of light. - 1 J 1 kg m2/s2

Answer

- C(3.00x108m/s)(.40)1.2x108m/s
- ? h/mv
- ? (6.626 x 10-34 Js) 6.06x10-12m
- (9.11x10-31kg)(1.2x108m/s)
- Remember 1J 1(kg)(m)2/s2

Wave-Particle Duality

- de Broglies experiments suggested that e- has

wave-like properties. - Thomsons experiments suggested that e- has

particle-like properties - measured charge-to-mass ratio

Quantum mechanical model

- SchrÖdinger
- Heisenberg
- Pauli
- Hund

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 theoretical

physicist, Werner Heisenberg.

Heisenbergs 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 - ALWAYS a basic uncertainty in trying to locate an

e-

Heisenbergs Uncertainty Principle

- Impossible to determine both the position and the

momentum of an e- in an atom simultaneously with

great certainty.

SchrÖdingers Wave Equation

- An equation that treated electrons in atoms as

waves - Only waves of specific energies, and therefore

frequencies, provided solutions to the equation - Quantization of e- energies was a natural outcome

SchrÖdingers 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

called orbitals

SchrÖdingers Wave Equation

- Here is the equation
- Dont memorize this or write it down
- It is a differential equation, and we need

calculus to solve it - -h (?2 ? ) (?2? )( ?2? ) V? E?
- 8(p)2m (?x2) (?y2) (?z2 )
- Scary???

Definitions

- Probability ? likelihood
- Orbital ? wave function region in space where

the probability of finding an electron is high - SchrÖdingers Wave Equation states that orbitals

have quantized energies - But there are other characteristics to describe

orbitals besides energy

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 SchrÖdingers

Wave Equation

Quantum Numbers (1)

- Principal Quantum Number, n

Quantum Numbers

- Principal Quantum Number, n
- Values of n 1,2,3, ?
- Positive integers only!
- Indicates the main energy level occupied by the

electron

Quantum Numbers

- Principal Quantum Number, n
- Values of n 1,2,3, ?
- Describes the energy level, orbital size

Quantum Numbers

- Principal Quantum Number, n
- Values of n 1,2,3, ?
- Describes the energy level, orbital size
- As n increases, orbital size increases.

Principle Quantum Number

n6

n5

n4

n3

Energy

n2

n 1

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 in a

given shell n2

Quantum Numbers (2)

- Angular momentum quantum number, l

Quantum Numbers

- Angular momentum quantum number, l
- Values of l n-1, 0

Quantum Numbers

- Angular momentum quantum number, l
- Values of l n-1, 0
- Describes the orbital shape

Quantum Numbers

- Angular momentum quantum number, l
- Values of l n-1, 0
- Describes the orbital shape
- Indicates the number of sublevel (subshells)
- (except for the 1st main energy level, orbitals

of different shapes are known as sublevels or

subshells)

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 n2, can have one of two shapes

corresponding to l 0 or l1 - Depending on its value of l, an orbital is

assigned a letter.

Orbital Shapes

- Angular magnetic quantum number, l
- If l 0, then the orbital is labeled s.
- s is spherical.

Orbital Shapes

- If l 1, then the orbital is labeled p.
- dumbbell shape

Orbital Shapes

- If l 2, the orbital is labeled d.
- double dumbbell or four-leaf clover

Orbital Shapes

- If l 3, then the orbital is labeled f.

Energy Level and Orbitals

- n1, only s orbitals
- n2, s and p orbitals
- n3, s, p, and d orbitals
- n4, s,p,d and f orbitals
- Remember l n-1

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
- Ex. 4d d sublevel in 4th main energy level

Quantum Numbers (3)

- Magnetic Quantum Number, ml

Quantum Numbers

- Magnetic Quantum Number, ml
- Values of ml l0-l

Quantum Numbers

- Magnetic Quantum Number, ml
- Values of ml l0-l
- Describes the orientation of the orbital
- Atomic orbitals can have the same shape but

different orientations

Magnetic Quantum Number

- s orbitals are spherical, only one orientation,

so m0 - p orbitals, 3-D orientation, so m -1, 0 or 1 (x,

y, z) - d orbitals, 5 orientations, m -2,-1, 0, 1 or 2

Quantum Numbers (4)

- Electron Spin Quantum Number,ms

Quantum Numbers

- Electron Spin Quantum Number,ms
- Values of ms 1/2 or 1/2
- e- spin in only 1 or 2 directions
- A single orbital can hold a maximum of 2 e-,

which must have opposite spins

Electron Configurations

- Electron Configurations arragenment of e- in an

atom - There is a distinct electron configuration for

each atom - There are 3 rules to writing electron

configurations

Pauli Exclusion Principle

- No 2 e- in an atom can have the same set of four

quantum numbers (n, l, ml, ms ). Therefore, no

atomic orbital can contain more than 2 e-.

Aufbau Principle

- Aufbau Principle an e- occupies the lowest

energy orbital that can receive it. - Aufbau order

Hunds Rule

- Hunds Rule orbitals of equal energy are each

occupied by one e- before any orbital is occupied

by a second e-, and all e- in singly occupied

orbitals must have the same spin

Electron Configuration

- The total of the superscripts must equal the

atomic number (number of electrons) of that atom. - The last symbol listed is the symbol for the

differentiating electron.

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
- B 1s2, 2s2, 2p1

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

indicated.

Orbital Diagram Rules

- Represent each electron by an arrow
- The direction of the arrow represents the

electron spin - Draw an up arrow to show the first electron in

each orbital. - Hunds 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. - Half filled sublevels are more stable than

partially full sublevels.

Orbital Diagram Examples

- H ?_
- 1s
- Li ?? ?_
- 1s 2s
- B ?? ?? ? __ __
- 1s 2s 2p
- N ?? ?? ? ? ?_
- 1s 2s 2p

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. - Because of the overlaying of the sublevels in the

larger atoms, there are never more than eight

valence electrons.

Rules for Dot Diagrams

- Xy

. .

Px orbital

S sublevel electrons

. .

Py orbital

Pz orbital

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. - Use the same rule (Hunds rule) as you fill the

designated orbitals.

Examples of Dot Diagrams

- H
- He
- Li
- Be

Examples of Dot Diagrams

- C
- N
- O
- Xe

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. - Al 3 valence electrons
- Br 7 valence electrons
- All transitions metals have 2 valence electrons.