Electronic Structure and Periodic Properties

Electronic Structure and Periodic Properties

Wave Nature of LightModels of the Atom

Bohr ModelQuantum Mechanical Model

Atomic OrbitalsElectron ConfigurationsPeriodic Properties of Elements

Electronic Structure of Atoms--Introduction

Elements in the same group exhibit similar chemical and physical properties.Alkali Metals:

softvery reactivemetal

Noble Gasesgasesinert (unreactive)

Why???


When atoms react, their electrons interact.

The properties of elements depend on their electronic structure.the arrangement of electrons in an

atomnumber of electronsdistribution of electrons around the atom

energies of the electrons


Understanding the nature of electrons and the electronic structure of atoms is the key to understanding the reactivity of elements and the reactions they undergo.

Much of our knowledge of the electronic structure of atoms came from studying the ways elements absorb or emit light.

The Wave Nature of Light

Light is a type of electromagnetic radiation a form of energy with both

electrical and magnetic components

Wavelength () the distance between successive peaks

Frequency () the number of complete wavelengths that pass a given point in 1 sec

x c c = 3.00 x 108 m/s (speed of light)


The electromagnetic spectrum:


Different types of electromagnetic radiation have different properties because they have different and .Gamma rays

wavelength similar to diameter of atomic nucleiHazardous

Radio waveswavelength can be longer than a football field

Quantized Energy and Photons

Classical physics (mechanics) suggests that both electromagnetic radiation and matter can have any energy:

A car rolling down a hill can have any potential energy (energy of position) depending on its position on the hill.


Classical mechanics is not correct, however.Max Planck suggested the idea that

energy is transferred in “packets” called quanta (plural).

Quantum: the smallest quantity of energy that can be emitted or absorbed as electromagnetic energy


Planck proposed that the energy of a single quantum is directly proportional to its frequency:

E = h

where E = energy = frequencyh = Planck’s constant (6.63x10-34

J-s)


According to Planck’s theory, energy is always emitted or absorbed in whole number multiples of h (i.e h, 2h, 3h)

According to Planck’s theory, the energy levels that are allowed are ‘quantized.’restricted to certain quantities or

values


In order to understand quantized energy levels, compare walking up (or down) a ramp versus walking up (or down) stairs:

Ramp: continuous change in height

Stairs: quantized changed in height You can only stop on the stairs, not between them


If Planck’s quantum theory is correct, why don’t we notice its effects in our daily lives?

Planck’s constant is very small (6.63 x 10-34 J-s).A quantum of energy (E = h) is very small.

Gaining or losing such a small amount of energy is: insignificant on macroscopic objects

very significant on the atomic level


In 1905 Einstein used Planck’s quantum theory to explain the photoelectric effect.Light shining on a clean metal surface

causes the surface to emit electrons.

The light must have a minimum frequency in order for electrons to be emitted.


Einstein explained these results by assuming that the light striking the metal is a stream of tiny energy packets of radiant energy (photons).

The energy of each photon is proportional to its frequency.

E = h


When a photon strikes a metal surface:Energy is transferred to the electrons

in the metalIf the energy is great enough, the electron can overcome the attractive forces holding it to the metal.

Any extra energy above the amount required to “free” the electron simply increases the kinetic energy of the electron.


Einstein’s explanation of the photoelectric effect led to a dilemma.Is light a wave or does it consist of particles?

Currently, light is considered to have both wave-like and particle-like properties.

Matter also has this same dual nature.

Atomic Models

Two models are used to explain the behavior and reactivity of atoms and ions.

Bohr model

Quantum mechanical model

Bohr Model

Bohr developed an atomic model that explained the line spectrum observed for the hydrogen atom.

Highvoltag

e

H2

When an electrical current is passed thru a sample of H2 (g), energy is transferred to the H2 molecules.

The molecules are broken up. The H atoms absorb energy and “jump” to a higher energy level.

The Bohr Model of the Atom

Highvoltag

e

H2

The H atoms “relax” back to their original energy level by giving off the absorbed energy as electromagnetic radiation.


Highvoltag

e

H2

The light is analyzed in a spectrometer by separating it into its different colors.


Highvoltag

e

H2

The separated colors are recorded as spectral lines.

Atomic spectru

m


The spectrum of atomic hydrogen consists of a series of discrete lines such as the ones shown previously.

Why would an atom emit only certain frequencies of light and not all of them?


According to the Bohr Model of the atom:

Electrons move in circular orbits around the nucleus.

Energy is quantized: only orbits of certain radii

corresponding to certain definite energies are allowed

an electron in a permitted orbit has a specific energy (an “allowed energy state”)


The allowed orbits have specific energies given by the formula:

En = (-RH) 1 where n = 1, 2, 3…n2

RH = Rydberg constant = 2.18 x 10-18 J

n is called the principal quantum number


Each orbit in an atom corresponds to a different value of n.As n increases the radius of the

orbit increases (i.e. the orbit and any electrons occupying it are further from the nucleus)

n=1 is the closest to the nucleus 0.529 Angstroms for the hydrogen atom


The energy of the orbit is lowest for n=1 and increases with increasing n.Lower energy = more stableLower energy = more preferred

state


The lowest energy state of an atom is called the ground state.n = 1 for the electron in a H atom

When an electron has “jumped” to a higher energy orbit (i.e. n = 2, 3, 4…) it is considered to be in an excited state.


To explain the line spectrum for hydrogen, Bohr assumed that an electron can “jump” from one allowed energy state to another.

Energy absorbed e- “jumps” to higher energy state

e- “relaxes” back to a lower energy state energy is emitted


n=1

n=2

n=3n=4

energy


Since the energies of the orbits in an atom are quantized, transitions from one allowed orbit to another involves only specific amounts of energy.

E = Ef - Ei


Since E = h, the energy of the light emitted can have only specific values.

Therefore the of the light can have only specific values as well.

So, the line spectrum for each element will be unique and will depend on the “allowed” energy levels in that element.

The Bohr Model of the Atom The Bohr model effectively explains the

line spectra of atoms and ions with a single electronH, He+, Li2+

Another model is needed to explain the reactivity and behavior of more complex atoms or ionsQuantum mechanical model

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Electronic Structure and Periodic Properties