The Periodic Law

History of the Periodic Table

By 1860, more than 60 elements had been discoveredMasses were not accurately known

In 1860, the first International Congress of Chemists met in GermanyStanislao Cannizzaro presented a

convincing method for accurately measuring the relative masses of atoms

Chemists were able to agree on standard masses

Dmitri Mendeleev was writing a chemistry textHe hoped to organize the elements based

on their propertiesHe placed the individual elements and their

properties on cardsHe noticed that when the elements were

arranged according to mass, certain properties appeared at regular intervals

Mendeleev created a table in which elements with similar properties were grouped together (1869)He reversed Te & I based on their propertiesHe left several empty spaces for

undiscovered elementsHe successfully predicted the discovery and

properties of Sc, Ga, and Ge

In 1911, Henry Mosely was working with Ernest RutherfordThey were examining the spectra of 38

different elementsThis was related to the number of protons,

he noticed that the properties fit this pattern better than mass

This explained why Te & I needed to be reversed

Modern periodic law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers

The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group

William Ramsay discovered argon in 1894

Helium had been discovered in the sun’s emission spectrum in 1868

To accommodate these elements, Ramsay had to create the column for the Noble gases

In 1898, he also discovered krypton and xenon

Radon was discovered in 1900 by Friedrich Dorn

The lanthanides were recognized as a group of similar elements in the early 1900’s

The actinides were also identified and to save space, they are placed with the lanthanides down below the main body of the table

Electron Configuration and the Periodic Table

Generally the electron configuration of an atom’s highest occupied energy level governs the atom’s chemical properties

Vertical columns are called groups or families and share similar physical properties

There are 7 horizontal periods in the table

There are group names for many individual families

The d block elements are known as the transition metals Typical metallic properties

The p block element properties vary greatlyThe metals are usually harder and denser

than s block metals, but softer and less dense than d block metals

The metalloids (semi metals) have properties of both metals and nonmetals

The actinides are all radioactive

Group 17 is known as the halogen group Most reactive nonmetals React with metals to form salts

Radiant Energy

Much of our understanding of how electrons behave comes from studies of how light interacts with matter

Until the 1800’s scientists believed that light was a beam of energy moving through space in the form of waves

In the 1900’s they found that light also behaved like a stream of tiny, fast-moving particles

Light travels in electromagnetic waves form of electromagnetic radiation

An electromagnetic wave consists of electric and magnetic fields oscillating at right angles to each other and the direction of the wave

All waves, whether they are water waves or em waves, can be described in terms of four characteristicsamplitude - height of a wave as measured

from the origin to its crest or peakwavelength - (l) distance between

successive crests of the wavevisible light is 400-750 nm (can see these w/

eye) frequency - (n) number of time the wave

cycles up and down in 1 secondexpressed as cycles per second (hertz - Hz)

1/s or s-1 (FM radio is in MHz)

speed - (c)in a vacuum is 3.00 x 108 m/s

The relationship between wavelength and frequency is:

= l c / = n n c / l

The visible spectrum is an example of a continuous spectrumone color fades into the next colorviolet has the shortest wavelength, highest

frequency red has the longest wavelength, lowest

frequencyvisible light is only a small part of the

total electromagnetic spectrum the rest is invisible to the human eye

Vis

UV

Xray

Quantum Theory

At the beginning of the 20th century, the wave model of light was universally accepted

several observations brought this acceptance into questionwhy do hot objects emit light of different

colors as they heat up (red>yellow>white)why do elements burn with different colors

In 1900, Max Planck was able to predict accurately how a spectrum changes w/ T to do this he proposed that there is a

fundamental restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy a quantum

He related the frequency of wavelength with its energyE = hnh is Planck’s constant =

6.6262 x 10-34 Js (Joule-seconds)

Using Planck’s theory, scientists can determine the temperature of far off objects (ex: stars) by observing their l

Energy is absorbed or emitted in quantahard to imagine in our world (ex: car)each quantum is very small, so energy

seems continuous to us these quanta can be significant on the

atomic levelPlanck’s discovery did not attract much

early attention

Albert Einstein saw a way to explain a puzzling phenomenon called the photoelectric effectFor each metal, a minimum frequency of

light is needed to release electrons regardless of the light intensity

Einstein said that when a photon (quanta) strikes a surface, it transfers it’s energy to an electronelectrons take all the energy or nonehigher frequency light has higher energy

photons that can release electrons to flow low frequency (low energy) photons cannot

release electrons

This is why x-rays are damagingand radio waves are not

Einstein won the Nobel Prize in1921 for explaining this in 1907

In 1923 Arthur Compton convincingly proved that light consists of tiny particles (photons)He demonstrated that a photon could

collide with an electronLight therefore has a dual nature

Particles & Waves

Another Look at the Atom

A spectrum that contains only certain colors, or wavelengths, is called a line spectrumevery element has a unique line spectrum

when it is heated or electricity is passed through it

also called an emission spectrumIncandescent light gives the complete

spectrum

Emission Spectrum

Absorption Spectrum

Spectra of different star typesThe hottest stars are on top, coolest on the bottomOur sun is toward the lower middle

Sun

Neils Bohr was able to explain why elements give line spectra In 1911 he attended a lecture where Ernest

Rutherford was explaining his planetary model of the atom

Bohr realized that Planck’s idea of quantization could be applied to this model to explain the line spectrahe started with hydrogeneach electron was allowed to have only certain

orbits corresponding to different energy levelshe gave each orbit a quantum number (n) the lowest orbit was the ground state (n=1)

When an electron absorbs an appropriate amount of energy, it jumps to a higher orbit, or excited state (n=2,n=3, n=4, etc.)

radiation is then emitted as the electron falls back to the ground state

Bohr was able to use this model and Planck’s equation (E=hn) to calculate the frequencies observed in the line spectrum of hydrogen

his model worked well for hydrogen, but could not explain the spectra of atoms w/ more than one electron, except in a rather approximate way

it represented an important initial step in our current understanding of electronic structure

Until 1900, scientists believed that there was a clear distinction between matter & energyPlanck, Einstein, & Bohr showed that

waves had particle propertiesIn 1924 Louis de Broglie wondered if

matter had waves (matter waves)eventually proven correct, won the Nobel

Prize, now used in electron microscopeswe are not aware of this, because they are

so smalla golf ball at 40m/s would have a wavelength of

3X10-34 mbecome significant on the atomic level

Electron micrograph of scratch & sniff paper w/ tiny glass capsules containing the scent

Louis deBroglie

Werner Heisenberg

-In 1927, Werner Heisenberg proposed the uncertainty principle

-the position and the momentum of a moving objectcannot simultaneously be measured and knownexactly-it is not appropriate to think in terms of electrons moving in well defined orbits, because there is noway to test this idea

A New Approach to the AtomThe quantum-mechanical model

explains the properties of atoms by treating the electron as a wave that has quantized energy

It is impossible to describe the exact positions of electrons or how they are moving

Describes the probability that electrons will be found in certain locations around the nucleus

The probability of finding an electron in various locations around the nucleus can be pictured in terms of a blurry cloud of negative energy

Cloud is referred to as electron densityThe probability of finding electrons in

certain regions of an atom is described by orbitals

An atomic orbital is a region around the nucleus of an atom where an electron with a given energy is likely to be found

Orbitals have characteristic shapes, sizes, and energies

They do not describe how the electron actually moves

Rather than drawing electron clouds to represent orbitals, it is more convenient to merely draw the surface within which an electron is found 90% of the time

Different kinds of orbitals have different shapess is spherical, p is dumbell shaped, d & f

are more complex

The main energy level or principal energy level in an atom is designated by the quantum number n the higher the n, the greater the energy

levelEach principal energy level is divided

into one or more sublevels the number of sublevels is equal to the

principal quantum number the first energy level has just the s sublevel the second energy level has an s & p

sublevelEach sublevel has a specific number of

orbitalss=1, p=3, d=5, and f=7

The higher the n, the larger theorbital

The three p orbitals combineinto the sublevel at the bottom

The 5 d sublevels

*Each orbital can contain 2 electrons-these will spin either clockwise or counterclockwise-to exist in the same orbital they must be spinning in opposite directions (opposite magnetic fields)-Called the Pauli exclusion principle after Wolfgang Pauli (1900-1958) Austrian

Electron Configurations

The distribution of electrons among the orbitals of an atom is called the electron configuration

Helps chemists to understand chemical behavior

The Aufbau principle (German for building up)Electrons are added one at a time to the

lowest energy orbitals available until all the electrons of an atom have been accounted for

Pauli exclusion principlean orbital can hold a maximum of 2

electrons, must spin in opposite directionsHund’s rule

electrons occupy equal energy orbitals so that a maximum number of unpaired electrons results

Arrangement for Carbon

Open arrangement of atomic orbitals

Orbital diagrams & configurations for several elements

Notice the sum of the superscripts in the configuration is equal to the number of electrons

The electron configurations represent the ground state of the electronsheating or electric current will cause

electrons to “jump” to higher levels when they fall, they emit specific amounts of

energy, creating the line spectra discussed earlier

Cr & Cu are exceptions to the Aufbau Principle due to the unique stability of a full or half full sublevel