Arrangement of Electrons in Atoms - images.pcmac.orgimages.pcmac.org/SiSFiles/Schools/TN/GreenevilleCity... · the arrangement of electrons around the nucleus. •The first # indicates

ATOMIC

THEORY 10/25/2012

WHAT ARE ATOMS? • The smallest units of matter.

• Make up elements

• Contain subatomic particles:

• A nucleus containing protons and

neutrons, and surrounded by electrons

• The number of protons and electrons is

unique for every element.

What’s in an atom?

Subatomic

Particles

Particles Charge Mass (kg) Location in

atom

Proton +1 1.67 x 10 -27 In Nucleus

Neutron 0 1.67 x 10 -27

In Nucleus

Electron -1 9.11 x 10 -31 Moving

around

outside of the

nucleus

What’s in an atom?

THE PERIODIC TABLE

• Horizontal (------) rows are called

periods.

• # of protons and electrons increases

by one as you move from left to right

across a period.

• Vertical (up and down) columns are

called groups or families.

• Atoms in the same group usually have

the same # of valence electrons

THE PERIODIC TABLE

• Atomic # (Z)=the number of protons and

electrons in a neutral atom

• Mass # (A)= the number of protons + the

number of neutrons

• # of Neutrons=A-Z

• # of Electrons= Z

• # of Protons= Z

FORMING IONS

• An ion is formed when an atom gains or

loses one or more electrons

• Cation(+)=When an atom LOSES one

electron, it gets a “+1” charge

• Anion(-)=When an atom GAINS an

electron, it gets a “-1” charge

ARRANGEMENT OF

ELECTRONS IN

ATOMS

HISTORICAL

MODELS • 4th century B.C.—Democritus

suggested that atoms existed as indivisible particles.

• 1808—John Dalton proposed the atomic theory

• 1897—Thomson—charge to mass ratio of electrons—plum pudding model

• 1911—Rutherford—nuclear model with electrons outside of the nucleus

DALTON’S THEORY

• 1. All matter is composed of indivisible

particles called atoms.

• 2. All atoms of a given element are

identical. Atoms of different elements have

different properties.

• 3. Chemical reactions involve the

combination of atoms, not the destruction of

atoms

• 4. When atoms react to form compounds,

they react in defined, whole-number ratios

JJ THOMSON

• —experiments with cathode ray

tubes (CRT’s) in 1897

CATHODE RAY

EXPERIMENT

Applying voltage causes a glow

to travel from the negative to

positive end glass tube in which

a partial vacuum exists.

What can we learn from a ray??

CATHODE RAY

OBSERVATIONS AND

CONCLUSIONS

• Ray is deflected by a magnetic

field.

• Has electromagnetic properties

• Ray travels away from negative

and toward positive

• Must be negatively charged

DEFLECTION OF CATHODE RAYS BY AN

APPLIED ELECTRIC FIELD.

CATHODE RAY

OBSERVATIONS AND

CONCLUSIONS

• Any metal will produce a ray.

• All atoms must contain these particles.

• Atoms are electrically neutral.

• Some positive particle must be present to balance the negative charge.

CHARGE VS. MASS

• Thomson was able to calculate

the charge to mass ratio = 1.7 x

1011

• VERY LARGE charge for a

very small mass

NEW MODEL

• Plum pudding—atoms are spherical masses of positive charge with electrons scattered throughout

• Since electrons have a small mass, the mass must come from something else.

MILLIKAN

• 1909 –Oil Drop Experiment

• Determined the mass and charge of

an electron

RUTHERFORD • 1911 –”Planetary Model”—Nuclear

atom with a nucleus at the center

• Suggested that electrons orbited the

nucleus like planets orbiting the sun

• Named protons and also suggested

the neutrons existed

MOSELEY

• 1914 – studied x-rays to provide

evidence that the atomic number of

elements is the same as the number

of protons in that element

BOHR

• 1922 –The Bohr Model of the atom

• Electrons only existed in specific energy

states

• They (e-) emit or absorb light at

specific wavelengths

• This accounted for line spectra

CHADWICK

• 1932 – Discovered that neutrons

existed in the nucleus of the atom.

NEW MODEL

• Modern Theory—Electrons

orbit the nucleus like waves on

a vibrating string.

• Based on observations of

absorption and emission of

electromagnetic radiation by

atoms

http://9-

4fordham.wikispaces.com/Electro+Magnetic+Spectrum+and+li

ght

ELECTROMAGNETIC

RADIATION

• Energy that travels in waves and can behave like particles

EM RADIATION

• Wave-like properties

• Wavelength—l—the distance between two corresponding points on successive waves

• Velocity—c—speed of light = 3.00 x 108 m/s

EM RADIATION

• Frequency—v—the number

of waves that pass a given

point in one second

• For any wave: c = lv

PROPORTIONALITY

• Since speed is constant (speed of light), frequency and wavelength vary inversely

• Frequency increases; wavelength decreases

• Frequency decreases; wavelength increases

PROPORTIONALITY

• Energy and wavelength vary inversely

• wavelength increases; energy decreases

• Energy and frequency vary directly

• Frequency decreases; Energy decreases

WHICH HAS HIGHER

ENERGY?

1. Wavelength of 400 nm or 700 nm?

2. Frequency of 108 Hz or 1018 Hz?

3. Wavelength of 1 m or 1 cm?

4. Frequency of 106 Hz or 1012 Hz?

CALCULATIONS

• What is the wavelength of EM

radiation with a frequency of

1016 Hz? (Hz= 1/s)

• What type of radiation is it?

(See electromagnetic

spectrum.)

WORKSPACE

PRACTICE

QUESTIONS

5. What wavelength of

electromagnetic radiation has

a frequency of 3.27 x 104 Hz?

6. What type of radiation is

this? (See electromagnetic

spectrum.)

CALCULATIONS

• What is the frequency of light with

a wavelength of 400 nm?

(400 nm = 4 x 10-7 m)

• What color is this light?

WORKSPACE

PRACTICE

QUESTIONS

7. What is the frequency of light

that has a wavelength of 6.84

x 10-12 m?

PARTICLE BEHAVIOR OF

LIGHT • Max Planck

• Objects do not emit light at continuous frequencies; rather the energy comes out in specific amounts or quanta

• Quantum—the minimum amount of energy an atom can gain or lose

ENERGY OF A

QUANTUM

• E = hv • E = energy (J)

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

• v = frequency (1/s)

CALCULATIONS

• What is the energy of a photon

of electromagnetic radiation with

a frequency of 1.56 x 1012 Hz?

WORKSPACE

CALCULATIONS

• What is the energy of a photon

of electromagnetic radiation with

a wavelength of 4.15 x 10-7 m?

WORKSPACE

PRACTICE

PROBLEM

8. What is the energy of a

photon which has a

wavelength of 8.24 x 10-4 mm?

PHOTOELECTRIC

EFFECT

• Emission of electrons from a metal when light shines on the metal

• If light is a wave, it should have enough energy to remove electrons at any frequency, but only certain frequencies of light will remove electrons.

EINSTEIN’S

CONTRIBUTION

• Light behaves as a particle and as a wave—wave particle duality

• Photon—a particle of light that carries a quantum of energy

• Energy of photon=energy of quantum—use same equation

EXCITED

ELECTRONS

• Electrons that have absorbed energy become excited and move to higher energy states in an atom

• When they return to lower energy states, they emit the energy

ATOMIC EMISSION

SPECTRA

• H always and only emits

radiation at specific

wavelengths

• Other elements behave the

same way, but patterns are

more complex due to larger # of

electrons

LINE EMISSION

SPECTRA

• Very specific—Act like

fingerprints by which an

element can be identified

***Emission Tubes/

Spectrometers***

DO THE MATH

• A mathematical relationship

exists between the

wavelengths and the emission

spectra of hydrogen.

• The spectral emission of

elements is always the same.

http://www.daviddarling.info/encyclopedia/H/hydrogen_spectru

m.html

BOHR MODEL • Energy of

emission equals the difference in energy levels

• Added orbits for electrons to the existing atomic model

http://www.robotplatform

.com/knowledge/Atomic

%20Theory/atom-

bohr_en.jpg

ENERGY LEVELS

• Electrons may only go to certain

places (energy levels) within the

atom.

• Evidence: wavelengths in

emission spectra of atoms

LIMITATIONS OF BOHR’S

MODEL

• Since Hydrogen has only

one electron, Bohr’s model

worked well for it.

• The model did not work as

well when applied to atoms

with more electrons.

DRAWING BOHR

MODELS • Only certain numbers of

electrons are allowed in each level

• Level 1: 2

• Level 2: 8

• Level 3: 18

• Level 4: 32

EXAMPLES

•Ca

•F

•Li

•C

S orbitals of a

hydrogen

atom—

electrons can

transition

between levels,

but cannot stay

in nodes http://www.avon-chemistry.com/electron_lecture.html

DON’T WRITE THIS!! It requires the same amount of

energy every time you

• walk from your seat to the café.

• walk from your seat to your locker.

• walk from your seat to the door.

If these were the only three places

you could go, you would only use

energy in those amounts.

ASSIGNMENT

• Draw Bohr Models for the 18

elements on page 13 of your

packet and answer the two

following questions.

THE QUANTUM

MODEL

THE QUANTUM

MODEL

• De Broglie: electrons have properties of both waves & particles—interference & diffraction.

• Wave-particle duality

• Waves could only exist with wavelengths that corresponded to Bohr’s orbits

HEISENBERG

UNCERTAINTY

PRINCIPLE

• If electrons are like both waves and particles, where are they in an atom?

• Electrons are located by their interaction with photons, but that interaction knocks them off course—as soon as you locate it, it’s somewhere else!

HEISENBERG

UNCERTAINTY

PRINCIPLE

It is impossible to know both the

position and velocity of an

electron or any other particle at

the same time.

SCHRODINGER’S WAVE

EQUATION

• Only certain frequencies of waves

solve the equation

• These frequencies correspond to

quanta of energy & Bohr’s orbitals

The Shrodinger equation is:

SCHRODINGER’S

WAVE EQUATION

• Schrodinger’s wave function describes a probable space in which electrons can be found

• Orbital—a 3-dimensional region around the nucleus that indicates the probable location of an electron

ELECTRON

CONFIGURATIONS

• An electron configuration is a way of showing

the arrangement of electrons around the

nucleus.

• The first # indicates the energy level

• The letter indicates the type of orbital: s, p,

d, f

• A superscript indicates the # of electrons in

the orbital.

• Ex: 1s2 = 1st energy level, ‘s’ orbital, 2

electrons

• Electrons occupy the lowest energy orbital first

• An atom will gain or lose electrons to get 8

electrons in the outer shell (the “Octet” rule)

• ‘s’ and ‘p’ orbitals are the outer shell

LOCATION ON THE PERIODIC

TABLE ALSO TELLS ELECTRON

CONFIGURATION

http://chemed.chem.wisc.edu/chempaths/GenChem-

Textbook/Electron-Configurations-and-the-Periodic-Table-

564.html

DIAGONAL RULE This diagram

shows the order

in which

electrons fill

orbitals. Begin

with 1s and

follow the

diagonal lines. http://www.wyzant.com/Help/Sc

ience/Chemistry/Electron_Confi

guration/

EXAMPLE E-

CONFIGURATIONS

• Hydrogen( ):

• Lithium( ):

• Magnesium( ):

SPECIAL

DIRECTIONS

• “D” orbital: The energy level for

the d-orbital is always one less

than the row #

• Ex: Zinc ( )

• “F” orbital: The energy level for

the f-orbital is always two less

than the row #

• Ex: Cerium ( )

EXCEPTIONS

Cr

we would predict:

1s2 2s2 2p6 3s2 3p6

4s2 3d4

But it is actually:

1s2 2s2 2p6 3s2 3p6

4s13d5

Cu

we would predict:

1s2 2s2 2p6 3s2 3p6 4s2

3d9

But it is actually:

1s2 2s2 2p6 3s2 3p6 4s1

3d10

NOBLE-GAS

NOTATION • AKA—abbreviated ground state

• Noble gases=elements in group 18

• Electron Configurations of Noble gases

which have eight electrons in their

outermost orbitals

• Uses brackets: [Ne] or [He]

• [Ne]=1s22s22p6

• Can be used to write shorthand versions of

other electron configurations:

• Ex: S=

ORBITAL

DIAGRAMS

Used to show the

location of electrons

within their energy

levels

ORBITAL

SHAPES

http://chemwiki.ucdavis.edu/Physical_Chemistry/Quantum

_Mechanics/Atomic_Theory/Electrons_in_Atoms/Electroni

c_Orbitals

IMAGERY OF THE P ORBITALS

http://www.introorganicchemistry.com/basic.html

ORBITALS • Each orbital holds 2 electrons

• Different numbers of orbitals for each sublevel

• 1 s orbital—holds 2 __

• 3 p orbitals—hold 6 __ __ __

• 5 d orbitals—hold 10 __ __ __ __ __

• 7 f orbitals—hold 14 __ __ __ __ __ __ __

THE AUFBAU

PRINCIPLE

• Each electron occupies the

lowest energy orbital

available

•Ex: An electron will be in

2s before 2p and 4s

before 6s

HUND’S RULE • Single electrons will occupy all

orbitals of the same sublevel

before pairing.

• All single electrons must have

the same spin.

PAULI EXCLUSION

PRINCIPLE

No two electrons in a single

atom can have the same set

of four quantum numbers.

Two electrons in a single

orbital must have opposite

spins

SPINNING

ELECTRONS

ORBITAL NOTATION

(ORBITAL DIAGRAM)

• Shows electrons in orbitals

• __ __ __ = unoccupied orbitals

• __ __ __ = orbitals occupied with

electrons of the same spin

• __ __ __ = orbitals occupied with

electrons with different spins

EXAMPLES

• Write orbital diagrams for N,

Ne, and Mg.

PRACTICE

PROBLEMS • In your lab journal, write

orbital diagrams for:

• Ca

• O

• Ne

• Fe