ATOMIC STRUCTURE

Electrons in Atoms

• In the early 1900’s scientists studying the behavior of atoms observed that certain elements emitted visible light when heated by a flame or if subjected to a high voltage spark– Much of our understanding of how electrons

behave in atoms comes from the study of light– In order to understand this, you must know

about the nature of light…

LIGHT• Visible light (the light we see with

our eyes) is a type of electromagnetic radiation

• All other electromagnetic radiation is invisible

ELECTROMAGNETIC RADIATION• Electromagnetic (EM) radiation is the

transmission and emission of energy in the form of electromagnetic waves– Visible light, X-rays, microwaves , infrared waves

(IR), ultraviolet waves (UV), radio waves, etc.

ELECTROMAGNETIC SPECTRUM

• The electromagnetic spectrum encompasses all forms of electromagnetic radiations

CHARACTERISTICS OF WAVES

• A wave can be thought of as a vibrating disturbance by which energy is transmitted

• Waves can be characterized by their:– Wavelength (lambda): distance between

identical points on a successive wave • nanometers (nm)

– Frequency (nu): the number of waves that pass through a particular point in 1 second• hertz (Hz); 1 Hz = 1 cycle/s

– Amplitude: the vertical distance from the midline to the peak or trough

Low Frequency

High Frequency

SPEED OF LIGHT• Another important property of waves is

their speed– dependent on type of wave and medium

it’s traveling through• The speed of a wave is a product of

its wavelength and frequency:c =

• Notice and are inversely proportional

• In a vacuum, All electromagnetic waves travel at the speed of light (3.00 x 108 m/s)

PLANCK’S QUANTUM THEORY• According to the theory atoms and molecules can emit

or absorb energy only in discreet quantities, which Planck called “quantum”

• quantum (meaning “fixed amount”) is the smallest quantity of energy that can be gained or lost by an atom

• The energy E of a single quantum of energy is given by

E =hwhere h is called Planck’s constant (6.63 x 10-34

J•s) and is the frequency of radiation• According to Planck’s theory energy is always released

or absorbed by matter in whole-numbers of h, 2h, 3h …

THE PHOTOELECTRIC EFFECT• Albert Einstein used

Planck’s quantum theory to explain the photoelectric effect

• When light shines on a clean metal surface the surface emits electrons (can be converted to electrical energy)

• For each metal there is a minimum frequency of light needed to emit electrons– red light in incapable of

releasing electrons from sodium metal even if intense

– faint violet light releases electrons easily

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DUAL NATURE OF LIGHT• In explaining the photoelectric effect Einstein

proposed that light (electromagnetic radiation) has dual nature:– Can behave like a wave– Can behave like a stream of particles (photons)

• PHOTON - a particle of electromagnetic radiation

– A photon has no mass

– A photon carries a quantum of energy

• Einstein calculated that a photon’s energy depends on its frequency

Ephoton = h

ATOMIC EMISSION SPECTRA• Light of a neon sign is produced by

passing electricity through a tube of neon gas

• The atoms in the tube absorb energy and become excited and unstable

• They become stable by releasing the energy as light

ATOMIC EMISSION SPECTRA

• If the light emitted from neon is passed through a prism neon’s atomic emission spectrum is produced

ATOMIC EMISSION SPECTRA• The atomic emission spectrum of an

element is its “fingerprint” • It’s the set of frequencies EM waves the

element’s atoms emit (give off)• Each element’s atomic emission spectrum is

unique and can be used to identify the element– Hydrogen’s emission spectrum:

violet 410 nm

blue-violet434 nm

blue-green486.1 nm

red656.2 nm

Light Bulb(white light)

Hydrogen Bulb

THE BOHR MODEL OF THE ATOM

• Bohr proposed:– Electrons move around the nucleus in circular

orbits (“rings”) with distinct energy levels• smaller orbits have lower energy, larger orbits

higher energy – In other words, electrons found closer to the

nucleus has less energy than electrons found at greater distances from the nucleus

• Bohr assigned a quantum number (n) to each level

+n = 7

n = 6

n = 5

n = 4n = 3 n = 2 n = 1

This model is often called the planetary model

BOHR’S MODEL OF THE ATOM

Energy

-

BOHR’S ATOM CONTINUED

• The lowest energy state of an atom is its ground state

• When an atom gains energy (through heating for example) it is in an excited state– in an excited state the electron absorbs the

energy & jumps to higher energy level when it jumps back down to its ground state it releases excess energy in the form of light• Even though hydrogen contains only one

electron, it can have many excited states

• Because electrons jump between orbitals that have specific energy levels only certain frequencies of electromagnetic radiation can be given off (only certain colors can be emitted):

• If an excited electron drops from n=3 to n=2 red light is emitted

• If an excited electron drops from n=4 to n=2 blue-green light is emitted

• 5-2: blue light• 6-2: violet light• This is how Bohr explained hydrogen’s emission

spectrum

BOHR MODEL CONTINUED

E = h

E = h

Wait!• Bohr’s model explained the emission spectrum

of Hydrogen, but it did not explain the emissions of any other element!

• It was eventually found that Bohr was incorrect:– Electrons do not travel in circular orbits

around the nucleus• In 1924, a young French scientist Louis de

Broglie (1892-1987) proposed an idea that eventually accounted for the fixed energy levels of Bohr’s model and better explained the behavior of electrons

THE QUANTUM MECHANICAL MODEL OF THE ATOM

• If waves can act like particles, particles can act as waves!

• Electrons behave like waves• Can’t know electrons position/path around the

nucleus:– Electrons move about in a cloud around the

nucleus in what appears to be a random pattern

• The Quantum Model only predicts where an electron is likely to be found

HEISENBERG UNCERTAINTY PRINCIPLE

• According to this principle it is fundamentally impossible to know the exact position and velocity of a particle at the same time

• locating an electron produces uncertainty in the position and motion of the electron

• It’s like trying to locate a helium filled balloon in a completely darkened room:– If you locate it by touch, you change

it’s position-Once you find it, it’s already somewhere else!

SO WHERE CAN ELECTRONS BE FOUND?• In the quantum model, the nucleus is

not surrounded by orbits, but by atomic orbitals

• Atomic Orbital: a three-dimensional pocket of space around the nucleus that the electron is most likely to be found– An electron has a 90% chance of

being found in the atomic orbital– That is the best we can do!

Electron probability density for hydrogen

e- density (1s orbital) falls off rapidly as distance from nucleus increases

Where 90% of thee- density is foundfor the 1s orbital

ATOMIC ORBITAL ORGANIZATION1. Principal energy level (n: 1-7)

• (n) indicates relative size and energy of orbital. As n increases so do energy and size

2. Energy sublevels (s, p, d, f)– sublevels labeled according to shape

– s: spherical; p: dumbbell; d/f: varied

3. Orbitals: Each sublevel has a specific number of orbitals:

– s: 1orbital– p: 3 orbitals– d: 5 orbitals– f: 7 orbitals Each orbital can hold two

electrons!!!

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increasingenergy

Energy

Each orbital can hold two electrons

2p

3p

1s

2s

3s

4s

5s4p

3d

4d

He Helium

2

4.003H

Hydrogen

1

1.008Li Lithium

3

6.941Be

Beryllium

4

9.012B

Boron

5

10.81C

Carbon

6

12.01N

Nitrogen

7

14.01O

Oxygen

8

16.00F

Fluorine

9

19.00Ne

Neon

10

20.18

ORDER OF ORBITALS (FILLING) IN MULTI-ELECTRON ATOM

ELECTRON CONFIGURATION

• An atoms electron configuration is the way an atom’s electrons are distributed among the orbitals of an atom

• The most state stable electron configuration is an atom’s ground state– Ground state: all electrons are in the lowest

possible energy state

• Electron configuration represented by writing symbol for the orbital and a superscript to indicate the number of electrons in the orbital Li: 1s2 2s1

The Pauli Exclusion Principle• The two electrons in an orbital

must spin in opposite directions

1s 2p2s

3s 3p

3d4s

paramagneticunpaired electrons

diamagneticpaired electrons

HUND’S RULE• Negatively charged electrons repel each other, so:

– Electrons won’t pair up unless they have to– Once there is one electron in every orbital…the

pairing will begin!

2s1s

2p

1.

2s1s

2p

4.

2s1s

2p

2.

2s1s

2p

3.

Add an electron:

Add an electron:

Add an electron:

ELECTRON CONFIGURATION• The periodic table can be divided into four

distinct blocks based on valence electron configuration

• electron configuration explain the recurrence of physical and chemical properties

SHORTHAND (NOBLE GAS) NOTATION

• Shows electron filling starting from previous noble gas:– Na: 1s22s22p63s1

– Noble gas configuration: [Ne]3s1

Electron Configuration for Cation and Anions

• What is the electron configuration for a sodium atom?– Na: 1s22s22p63s1

• What is the electron configuration for a sodium ion (Na+)?– Na+: 1s22s22p6 or [Ne]

TRANSITION CATIONS• When transition metals form cations

electrons are removed from the s-orbital before the d-orbital – Mn: [Ar]4s2 3d3

– Mn2+: [Ar]3d3

• This happens because d-orbitals are more stable than s-orbitals in transition metals

• The s-orbitals are always in a higher energy level

ISOELECTRONIC • What do you notice about the

following electron configurations:– F- : 1s22s22p6 or [Ne]– O2- : 1s22s22p6 or [Ne]– N3- : 1s22s22p6 or [Ne]

• They have the same number of electrons and therefore identical ground-state electron configurations– These three ions are isoelectronic

VALENCE ELECTRONS• Valence electrons are the electrons

found in the outermost orbitals of the atom (highest energy level)

• They are the electrons involved in bonding and are responsible for the chemical properties of an element

• Carbon has 4 valence electrons: 1s2 2s2 2p2

• How many does magnesium have?

VALENCE ELECTRONS & GROUP NUMBER

• One of the most important relationships in chemistry:– Atoms in the same group have similar

chemical properties because they have the same number of valence electrons!

• For the representative elements (group A elements) – group # = number of valence electrons– Exceptions: He is in group 8 but only

has 2 valence electrons

The octet/duet rule• Atoms will gain, lose or share

electrons to achieve noble gas configuration, meaning all atoms want a full outer orbital:– 2 valence electrons for He– 8 valence electrons for all other noble

gases

ELECTRON-DOT STRUCTURE• Chemists often represent

valence electrons in electron-dot structures

• Electron-Dot Structure consists of the element’s symbol surrounded by dots representing the atom’s valence electrons– valence electrons are placed one to

each of the four sides first, – when each side has one dot, you may

begin doubling up S