Properties & Structure

Atomic Theories

The stuff you already know!

All matter occupies volume and has mass.All matter is comprised of particles.These particles are organized into atoms,

molecules & compounds.Each atom is the fundamental component

of an element.Each element has characteristic physical

& chemical properties.

The stuff you already know!

The atom has subatomic particles called;Electrons –responsible for the element’s

chemical behaviour.Protons – responsible for the element’s

physical and chemical properties.Neutron – responsible for the variance in the

population of an element’s atoms’ masses. Isotopes of the same element have identical properties.

The stuff you already know!

Particle Mass (u) Electrical Charge Location

Electron 0.0005485712 1- Outside nucleus

Proton 1.00727252 1+ Nucleus

Neutron 1.008665 0 Nucleus

The stuff you already know!

There are about 110 elements of which about 23 are synthetic or man-made.

They are named using conventions adopted by the IUPAC.

They are represented with symbols.1,2 or 3 letters of which the first is the only

one capitalized.The symbols are internationally recognized as

the name varies with the local language.

The stuff you already know!

The elements are noted with a “standard notation”.

A the mass number (protons + neutrons)Z the atomic number (protons)X the chemical symbol

Isotopes are atoms with the same number of protons but different numbers of neutrons.

XAZ

Stuff you might remember…

Early Greek Theories of Matter500 BCGreek philosopher Democrities believed that

substances were composed of indivisible particles called atoms (Greek for “indivisible”)

Atoms of different sizes, regular shapes, and are in constant motion.

Empty space between atoms.

Stuff you might remember…

Aristotle (384-322 BC)Criticized Democrities’s theory

All matter is made up of four basic substances:EarthAirFireWater

Each basic substance had different combinations of four specific qualities:

Dry, hot, moist & cold

Aristotle’s opinion prevails among scholars!

Dalton’s Atomic Theory (1805)Experimental Work

Theoretical Explanation

Atomic Theory

Law of definite proportions: elements combine in a characteristic mass ratio.

Each atom has a particular combining capacity.

Matter is composed of indestructible, indivisible atoms, which are identical for one element, but different from other elements.Law of multiple

proportions: there may be more than one mass ratio.

Some atoms have more than one combining capacity.

Law of conservation of mass: total mass remains constant

Atoms are neither created nor destroyed in a chemical reaction.

Thomson Atomic Theory (1897)Experimental Work Theoretical

ExplanationAtomic Theory

Arrhenius: the electrical nature of chemical solutions.

Atoms may gain or lose electrons to form ions in solution.

Matter is composed of atoms that contain electrons (negative particles) embedded in a positive material. The kind of element is characterised by the number of electrons in the atom.“The cookie model for the atom.”

Faraday: quantitative work with electricity & solutions.

Particular atom and ions gain or lose a specific number of electrons.

Crookes: qualitative studies with the cathode ray.

Electricity is composed of negatively charged particles.

Thomson: quantitative studies with the cathode ray.

Electrons are a component of all matter.

Millikan: charged oil drop experiment

Electrons have a specific fixed electric charge.

Rutherford Atomic Theory (1911)Experimental Work

Theoretical Explanation

Atomic Theory

Rutherford: A few positive alpha particles are deflected at large angles when fired at gold foil.

The positive charge in the atom must be concentrated in a very small volume of the atom.

An atom is comprised of a very tiny nucleus which contains positive charges and most of the mass of the atom. Very small negative electrons occupy most of the volume of the atom.

Most materials are very stable and do not self destruct or disintegrate.

A very strong force holds the positive charges together within the nucleus. (Strong nuclear force)

Rutherford: Most alpha particles pass straight through gold foil.

Most of the atom is empty space.

Quantum Theory

The next generation in the knowledge on matter!

Protons, Neutrons & IsotopesExperimental Work Theoretical

ExplanationAtomic Theory

Rutherford (1914): The lowest charge on an ionized gas particle is from the hydrogen ion.

The smallest particle of positive charge in the proton.

Atoms are comprised of protons, neutrons and electrons. Atoms of the same element have the same number of protons and electrons, but may have varying number of neutrons, resulting in isotopes of the element.

Soddy (1913): Radioactive decay suggests different atoms of the same element.

Isotopes of an element have fixed number of protons but varying stability and mass.

Aston (1919): Mass spectrometer work indicates different masses for some atoms of the same element.

The nucleus contains neutral particles called neutrons.

Radiation is produced by bombarding elements with alpha particles.

A little background . . . . .At the turn of the century, Physics and Chemistry

become entwined as the study of matter and energy become the focus of the scientific community.

Light energy is of primary interest – does it travel as a particle or a wave?

There is insurmountable evidence to support the wave theory of light and mathematical models supporting the observations.

Maxwell and others define and study the electromagnetic spectrum and the characteristics of its energy. Light is included in this spectrum.

Further investigation of the spectrum uncovers discrepancies that can’t be explained with conventional theories of the time.

Enter Max Planck (Plonk not Plank!)

Max Planck (1858-1947) His teacher, Gustav Kirchhoff, investigated the

energy emitted from blackbodies. Planck continued in this vein.

Blackbodies -objects that absorb light then radiate energy– Heat up a piece of metal and it turns “red hot”, then

orange, the blue, then “white hot”.Turn on the stove element and it soon emits heat (IR

radiation) and red light (light energy) The belief of the time was – as you increase the

intensity of the incident energy source the emitted energy would increase in energy, as found in the EMS spectrum.

Max Planck (1858-1947)In

ten

sity

of

inci

den

t en

erg

y

IR ROYGBV UV

Emitted energy from Blackbody

Classical Physics prediction

Observed Results(“The Blackbox Catastrophe”)

Max Planck (1858-1947)

The results lead to such a revolutionary concept in energy that Planck himself doubted it.

Energy was transmitted and absorbed in “bundles”, “packets” or QUANTA rather than a continuum of energy.

Matter can absorb or emit only discrete quantities of energy called quantum of energy.

(Einstein refines the terminology!) Energy transfer is like working on a staircase

rather than a ramp!

More with Max & the Boys!E

ner

gy

Intensity

E=hf

Quantum of energy

The Photoelectric Effect The curiosity around light was how its energy was

transmitted – Particle or Wave. This environment lead to much research in the properties of light.

Heinrich Hertz discovered the Photoelectric Effect when a piece of charged zinc was hooked up to an electroscope and then UV light was shone on the zinc. The charge on the electroscope was reduced.

The UV light generated an electric current. The use of a photosensitive cathode ray tube allowed

scientists to study and quantify the photoelectric effect. Photoelectric effect – light energy releases electrons

from a photosensitive surface. Einstein produces a model for the photoelectric effect

that relates the frequency of the photon to the energy of the emitted photoelectron.

Creating Quantum TheoryExperimental Work Theoretical

ExplanationQuantum Theory

Kirchhoff (1859): Identifies and investigates blackbody radiation

Planck (1900): The energy from a blackbody is quantized; i.e. emitted energy is restricted to whole number multiples of a finite quantity of certain energy

Electromagnetic energy is not infinitely subdivisible; energy exists as packets or quanta, called photons. A photon is a small packet of energy corresponding to a specific frequency of light. (E=hf )

Hertz (1887): Investigates the photoelectric effect

Einstein (1905): The size of a quantum of electromagnetic energy depends directly on its frequency; one photon of energy ejects one electron

Bohr Atomic Theory (1913)The problem with Rutherford’s model was

in the orbiting electrons.Moving electrical charges produce

electromagnetic energy.The emission of such energy would result in a

net loss of energy with the moving electrons.The reduction in the electrons’ energy would

cause them to fall into the nucleus and thereby, the destruction of the atom.

A better way to define the energy of the electrons was needed.

Bohr’s Atomic Line Spectra Spectroscopy is the study of light and its composite

“colours” or wavelengths ( IR, Visible light, UV, X-ray, etc.)

Kirchhoff & Bunsen had developed a procedure as early as 1859.

An emission spectrum of gas is when light from a “glowing” gas is passed through a prism to separate the emitted light into its component colours.

The result is a BRIGHT-LINE SPECTRA. An ABSORPTION SPECTRUM is formed by placing a

gas in the path of a continuous spectrum (white light). The gas absorbs the light energy and dark lines are found in the continuous spectrum.

Bohr utilized these techniques to look at matter and explain the unique spectra produced with each element.

Bohr’s postulates1. Electrons do not radiate or emit energy

as they orbit the nucleus. They are in a constant state of energy called a “ground state” (stationary state).

2. Electrons can change their energy only through the absorption of sufficient energy to “jump” to another stationary state. This is called “transition”.

Analysis of Bohr’s atomic modelIt explained the observed atomic line

spectrum for Hydrogen very well.It predicted the number of electrons in

each energy level well.# electrons in an energy level = 2n2

He predicted the IR and UV spectra for hydrogen correctly.

The success he had with Hydrogen was not found with other elements.

Creating the Bohr Atomic Theory (1913)Experimental Work Theoretical

ExplanationBohr’s Atomic Theory

Mendeleev (1869-1872): Periodic law – Elements expressed periodic trends in their physical and chemical properties.

A new period begins in the periodic table when a new energy level is started.

Electrons travel in the atom in circular orbits with quantized energy & the energy is restricted to only certain discrete quantities.There is a maximum number of electrons permitted in each orbit.Electrons “jump” to higher energy levels when a photon of energy is absorbed and a photon of energy is emitted when the electron “drops” to a lower level.

Mendeleev (1872): There are two elements in the first period and eight in the second period of the periodic table

There are two electrons maximum in the first energy level & a maximum of eight in the next level. 2n2

Kirchhoff, Bunsen (1859), J. Balmer (1885): Emission and absorption line spectra rather than continuous spectra exist for gaseous elements.

Since the energy of light absorbed and emitted is quantized (E=hf), the energy of the electrons in atoms is quantized.

Bohr’s Atomic Model

Hydrogen’s

Energy Level Diagram

n = 1

n = 2

n = 3

e-

Ground State – stationary state

Energy is added to the gas sample.

e-

A quantum of energy is absorbed by the electron and

it undergoes “transition”

Excited State – stationary state

A quantum of energy, equal to the absorbed quantity,

is released as a “photon” (EMR)

when the electron returns to its ground state.

Bohr’s explanation of the bright line spectra

Hydrogen’s

Energy Level Diagram

n = 1

n = 2

n = 3 e-Emission Line

Spectrum

IR R O Y G B V UV

Modifications to Bohr’s Atomic Model As Bohr’s work progressed more scientists studied the

phenomenon of bright line spectra under specific conditions with more advanced technologies. These studies resulted in modifications to Bohr’s theory.

Sommerfeld found that the some of the single lines were actually multiple lines very close together, hence sublevels in the principal energy levels were proposed.

Zeeman found that some of the lines would separate into multiple lines when the gas discharge tube was subjected to a strong magnetic field.

Further magnetic analysis suggested that electrons could spin one of two directions – clockwise & counter clockwise.

These discoveries lead to the adoption of four quantum numbers to describe an electron’s energy and magnetic characteristics.

Quantum Numbers The Principal Quantum Number (n) is associated

with the main energy level the electron occupies.n = 1,2,3,4,. . . .

The Secondary Quantum Number (l ) defines the shape of the electron’s energy subshells.

l = 0,1,. . .,n-1 The Magnetic Quantum Number (ml ) relates to

the direction of the electron orbit and its orientation.

ml = -l,. . . ,0, . . ,l The Spin Quantum Number (ms ) relates to the

direction in which the electron spins.ms = -1/2, 1/2

Creating the Four Quantum Numbers

Experimental Work

Theoretical Explanation

Quantum Theory

Low-resolution line spectra

Principal quantum number, n

All electrons in all atoms can be described by four quantum numbers which relate their energy and magnetic properties.

High-resolution line spectra

Secondary quantum number, l

Spectra in magnetic field Magnetic quantum number, ml

Ferromagnetism and paramagnetism

Spin quantum number, ms

Summary of Quantum Numbers & restrictions

Principal quantum number, n: the main electron energy level or shell (n )

Secondary quantum number, l: the electron sublevels or subshells (0 to n-1)

Magnetic quantum number, ml: the orientation of the sublevel (-l to +l)

Spin quantum number, ms: the electron spin (-1/2 to +1/2)

Energy shell Orbital shape Orbital orientation

Electron Spin

1 0 0 +1/2,-1/2

201

0-1,0,+1

+1/2,-1/2+1/2,-1/2

3012

0-1,0,+1

-2,-1,0,+1,+2

+1/2,-1/2+1/2,-1/2+1/2,-1/2

Creating Energy-Level Diagrams The electrons occupy specific energy levels in

the atom. The energy-level diagram is a method of representing an atom’s distribution of electrons throughout the various energy shells (n) and subshells or orbitals (l ).

The energy-level diagrams can be extended to include all four quantum numbers.

Experimental information about atomic matter shows that there are relationships between orbitals and their energy levels. Conventions have been established to attend to these observations.

Conventions for Creating Energy-Level Diagrams Pauli exclusion principle –

no two electrons in an atom may have the same four quantum numbers

no two electrons in the same orbital may have the same spin

only two electrons with opposite spins may occupy an orbital

aufbau principle – (German for “building up’)each electron is added to the lowest available energy

orbital Hund’s rule –

one electron is placed in each orbital at the same energy level before the second electron is placed

Conventions for Creating Energy-Level DiagramsCircles or squares are used to represent

the orbitalsArrows are used to represent the electrons

up represents one electron rotation (clockwise) while down the other (counter clockwise)

there is no convention as to which one you must start with

Conventions for Creating Energy-Level Diagrams

1s 2s 2p 3s 3p

O (z = 8)

1s 2s 2p 3s 3p

P (z = 15)

1s 2s 2p 3s 3p

Ar (z = 18)

Conventions for Creating Energy-Level Diagrams

•Energy-level diagrams may be written in a vertical manner to exemplify the energy level subtleties.

Ord

er o

f fil

ling

orbi

tals

1s

2s

3s

4s

5s

6s

2p

3p

4p

6p5d

4f

5p4d

3d

2e-

2e-

6e-

2e-

6e-

2e-10e-6e-

2e-10e-6e-

2e-14e-10e-6e-

2e-

8e-

8e-

18e-

18e-

32e- •As the number of energy levels and orbitals increase, so too does the complexity of the energy-level diagram.

•The diagram indicates nicely the order in which the orbitals are filled

Energy-Level Diagrams for ions The energy-level diagram is created in the same

manner as the regular atom. However, a surplus or deficit of electrons are included.For anions – add the proper number of electrons

using regular conventions

For cations – remove the correct number of ions in the proper manner

1s 2s 2p 3s 3p

S 2- (z = 16)

1s 2s 2p 3s 3p

Al 3+

(z = 13)

Electron configurations The energy-level diagram is the best way to visualize

the energy relationships between electrons. However it may prove cumbersome.

Electron configurations show the same information in a more concise manner.

Electron configurations

becomes . . .

Cl: 1s2 2s2 2p6 3s2 3p5

In the shorthand form of electron configurations, the previous noble gas structure is used to reflect the lower energy level electrons. So . . .

Cl: 1s2 2s2 2p6 3s2 3p5

becomes . . .

Cl: [Ne] 3s2 3p5

1s 2s 2p 3s 3p

Cl (z = 17)

Explaining the Periodic Table •As the evolution of the quantum theory is impressive in its experimental

analysis and explanation of atomic spectra. However, it must also attend to the years of chemical experimentation that are reflected in the Periodic Table and its established merits.

•Quantum theory does indeed support the structure of the periodic table.

•Analysis of the Electron Subshells and the Periodic Table reveal remarkable corroboration. Period # of elements Electron distribution

groups :1-2 13-18 3-12 -orbitals: s p d f

Period 1 2 2

Period 2 8 2 6

Period 3 18 2 6 10

Period 4-5 18 2 6 10

Period 6-7 32 2 6 10 14

Explaining the Periodic Table

4f5f

1s2s3s4s5s6s7s

1s2p3p4p5p6p7p

3d4d5d6d

s-block

2 groups

f-block

14 groups

d-block

10 groups

p-block

6 groups

1s

2s

3s

4s

5s

6s

2p

3p

4p

6p5d

4f

5p4d

3d

2e-

2e-

6e-

2e-

6e-

2e-10e

-

6e-

2e-10e

-

6e-

2e-14e

-

10e-

6e-

2e-

8e-

8e-

18e-

18e-

32e-

Period 1

Period 2

Period 3

Period 4

Period 5

Period 6

Group1,2 13-18 3-12 La & Ac series

•As the orbitals fill up with electrons a balance between the lowest energy orbital and the tendency to fill each orbital with a single similar spinning electron.

•Completely filled orbitals are more stable than half filled orbitals which are more stable than partially filled orbitals.

> >

•As the principle energy level increases the difference between each orbital’s energy level becomes reduced to the point where the d-orbitals energies are very close to the s-orbital’s.

•In situations like this the “s-” and “d-orbitals” may be treated as a similar energy level and the application of Hund’s rule may be applied over the range of orbitals rather than the single second quantum orbital.

•The stability of having each orbital containing one electron is greater than having a partially filled group of orbitals and a filled orbital of only minimally less energy.

Anomalies in Electron Configurations

Anomalies in Electron Configurations

•Vanadium has the electron configuration of –

V (z=23) [Ar] 4s2 3d3

•Chromium has one extra electron and conventions would predict it to be added to the next open d-orbital.

Cr (z=24) [Ar] 4s2 3d4

•However, the stability of having one electron in each orbital of similar takes president and the true electron configuration becomes

Cr (z=24) [Ar] 4s1 3d5

n=3

n=4

s p d

Explaining the multivalent ions

•Many transitional metals have the ability to have multiple charges as ions and the explanation for this behaviour has yet to be explained.

•As orbitals are being filled there are varying levels of stability due to the interaction of forces or electrostatic repulsion between electrons and force associates with the magnetic field due to the electron’s spin.

•Filled orbitals are most stable because the electrostatic repulsion is balanced against the magnetic attraction.

e- e-Electrostatic repulsion

e-South

Magnetic fieldMagnetic attractione-North

Magnetic field

Explaining the multivalent ions •Orbitals that are completely filled, like Noble gases, have the most stable

structure due to the balanced forces between electrostatic repulsion and magnetic attraction.

•Transitional metals often have partially filled d-orbitals and complete s-orbitals with similar energy levels. Electrons are lost to achieve the best combination of stability.

•Best stability – completely filled orbitals (2 electrons/orbital)Electrostatic repulsion and Magnetic attraction

•Next best stability – half filled orbitals (1 electron/orbital)Electrostatic repulsion with minimal crowding

Co: [Ar] 4s2 3d7

Co: [Ar]

Co2+: [Ar]

Co3+: [Ar]

- Filled 4s and half filled 3d

- Half filled 4s & 3d less stable than Co2+ so less common ion

Explaining Magnetism •Some materials exhibit strong magnetic properties naturally these are

referred to as being ferromagnetic.

•As moving charges have magnetic fields so too do spinning electrons. Those spinning in one direction would have one magnetic polarity as those spinning in the opposite direction the opposing magnetic pole.

•For example – clockwise – south pole and counter clockwise –north pole

•Those atoms that have a number of similarly spinning electrons would have similar magnetic fields.

•If these “magnetic atoms” were free to align themselves with neighbouring atoms of similar characteristics they would create regions of magnetism in the material called domains.

•The arrangement of these domains in a material results in a magnet.

Explaining Magnetism Iron, nickel and cobalt are such naturally occurring atoms.

Iron, nickel and cobalt are small enough atoms that they can realign themselves due to the magnetic properties of their surrounding and thereby create domains.

There are other such atoms that have similar electron configurations but limited ability to migrate. Hence, they are reduced in their magnetic properties.

Explaining Magnetism

Ferromagnetic – materials with strong magnetic properties. Their presence increases a magnetic field substantially.

Paramagnetic – materials with weak magnetic properties. Their presence only slightly strengthens a magnetic field.

Diamagnetic – materials that have reduced magnetic properties. Their presence weakens a magnetic field.

Wave Mechanics & Orbitals

Louis de Broglie used the accepted concept that light behaved as a particle as well as a wave to suggest that all matter might have wave properties.

Experimental evidence did show that particles (electrons) do have wave properties.

Matter waves were established!

Erwin Schrodinger applied de Broglie’s concept to explain the behaviour of electrons in the atom.

The foundation lies in the fact that electrons are in a stable standing wave pattern. This pattern requires a specific whole number of wavelengths to generate a standing wave. (quantum of energy)

He created mathematical models to emulate and explain the wave properties of electrons. Which is usually referred to as Quantum Mechanics.

In the atom’s environment the energy of the standing wave is not lost to friction and other such resistant factors.

The quantum mechanics theory and research has given viable explanations of the electron’s energy and characteristics, but little on its location or motion.

Shapes of Orbitals The act of measuring or identifying an electron

requires us to interfere with its motion and energy. This interference may result in an inaccuracy in

identifying the very characteristic we intended to measure.

The act of measuring the location of an electron requires the absorption or release of its energy, which in turn affects its speed and maybe location.

Heisenberg’s uncertainty principle indicates that it is impossible to know the exact position and speed of a particle at the same time.

Such an understanding reinforces a sense of probability in determining the location of an electron.

Shapes of Orbitals Schrodinger’s wave equations can be used to

predict the likelihood of “finding” and electron at a specific location. These probabilities can be used to plot an electron probability density.

The second quantum number, l, indicates the variety in “shape” of the orbital. (s – sharp, p – principal, d – diffuse, f – fundamental) Recent theory suggests that there may be g-orbitals!?!

An orbital is associated with a size, three dimensional shape and orientation around the nucleus.

Together the size, shape and position represent the probability of finding a specific electron at that location.

Shape of orbitals

The probability density plot can be assessed in multiple dimensions to generate a 3-d density cloud or “shape” for the orbital.

Shape of orbitals

The 2s-orbital electrons have specific characteristics and when the probability density plot for these electrons is produced, new “shapes” are suggested.

Shape of orbitals

p-orbital

s-orbital

d-orbital

Shape of orbitals

The culmination of all of the electrons produces a combined effect from all involved orbitals to generate a unique “orbital shape” for each atom. For example, the structure of Ne . . .

Shape of orbitals The diagram we used to represent oxygen is;

8 Protons

-

-

-

-

-

-

-

-

8

16O

Shape of orbitals The diagram we might currently use to

represent oxygen is;