Unit V: Atomic Theory

Vocabulary: atoms, ions, compounds, molecules, quantum theory, block structure of the Periodic Table, sub-atomic, particles, electrons, protons, neutrons, isotopes, radioisotopes, radioactive decay, average atomic mass vs. mass number, standard notation, plum pudding model, spectroscopy, electromagnetic spectrum, absorption and emission spectra, line spectra vs. continuous spectra, cations and anions, mass spectrometer, mass spectra, electron arrangement vs. electron configuration, ground state, excited state, photon, quantization, Planck’s equation and Planck’s constant, ionization energy (first and successive), Bohr model, valence electrons, sublevels of electrons (s,p,d,f), wave-particle duality, Heisenberg’s Uncertainty Principle, atomic orbitals, Pauli Exclusion Principle, Aufbau Principle, Hund’s third rule

1. Introduction

The idea that all matter is composed of atoms is one of the most important discoveries in human history. Although Greek and Indian philosophers had the idea 2500 years ago, we have only had experimental evidence since the 19th century. This unit is mostly theoretical, based on the following key concepts:

The everyday notion of Newtonian, classical physics does not apply at the sub-atomic level. Particles do not follow fixed paths.

Light behaves as a wave and as a particle. All matter behaves as a particle and as a wave. The wave-like properties are more significant the smaller the particle.

Electrons are not found in defined orbits but in regions of probability called orbitals.

2. Historical background

a) Dalton

John Dalton, in the 19th century, noticed that hydrogen and oxygen always combine in the same proportions to make a new substance called a compound (in this case, water). Compounds have very different properties than their component elements. Dalton proposed:

All matter is composed of tiny indivisible particles called atoms

Atoms cannot be created or destroyed

Atoms of the same element are alike in every way

Atoms of different elements are different

Atoms can combine together in small numbers to form molecules

*Note: the term “molecule” is now usually used for covalent compounds i.e. where e- are shared between non-metals.

b) Thompson

J.J. Thompson, at the end of the 19th century, discovered that different metals produce a stream of negatively charged particles when a high voltage is applied across two electrodes. This device is called a cathode ray tube (CRT) and is the technology behind tube TVs and other devices. He called these negative particles “corpuscles”, now called electrons.

The electrons were the same regardless of the metal used, so Thompson concluded they were found in all atoms, and created the Plum Pudding model of the atom:

c) Rutherford

ACT – “Atomic Target Practice”

Rutherford’s Scattering Experiment a.k.a. “Gold Foil Experiment” was one of the most important experiments in history. The key findings:

The mass of an atom is concentrated in a very small area (nucleus) and atoms are mostly empty space

The nucleus has a positive charge

Draw:

“It was quite the most incredible thing that has happened to me. It was if you had fired a (artillery) shell at a piece of tissue paper and it came back to hit you.” – Ernest Rutherford (1871-1937)

3. Subatomic particles:

Particle Relative mass Relative charge

Proton 1 +1

Electron 0.0005 -1

Neutron 1 0

Atomic number: the # of protons in the nucleus.

Mass number: the # of protons plus the # of neutrons in an atom

Standard notation:

4. Isotopes

Isotopes: atoms of the same element with different mass numbers.

On the periodic table, the relative atomic masses (a.k.a. molar mass) are often given as decimal (e.g. chlorine is 35.45) because that number is an average relative atomic mass for all of the naturally occurring isotopes.

Scientists can measure the masses of atoms, and therefore find the percent abundance of isotopes, using a mass spectrometer.

ACT – Bean Bag Isotopes

Radioisotope: an isotope that is radioactive. Having too few or too many neutrons can result in unstable nuclei. These nuclei decay (called nuclear decay) into a more stable nucleus. Recall from Grade 10 the types of decay: alpha, beta and gamma.

e.g. iodine-131, a product of nuclear fission in nuclear power plants, was found in Vancouver seaweed shortly after the Japanese power plant disaster of 2011.

Radioisotopes can also be of great benefit when used safely.

e.g. cobalt-60 is used to kill cancer cells via gamma decay, commonly called “radiation therapy”

Other uses of radioisotopes:

iodine-131 and iodine-125 used to treat cancer

sterilizing surgical instruments

preserving food

fighting crime

detecting cracks in structural materials

5. Ions

When an atom loses or gains electrons it becomes an ion. Ions have charges because the number of protons (positive charges) is no longer equal to the number of electrons (negative charges).

Cation: a positively charged ion. Anion: a negatively charged ion

Ionic compound: a compound that results from the attraction of cations and anions to each other e.g. NaCl

6. The Mass Spectrometer

The mass spectrometer allows us to measure the masses of atoms. Therefore it also allows us to find the relative abundance of isotopes for an element.

Draw:

5 Basic Operations:

a) Vaporization – allows individual atoms of the element to be analyzed

b) Ionization – the atoms are bombarded with high-energy electrons which knock the electrons out of the atoms, producing positively charged ions

c) Acceleration – the positive ions are attracted to negatively charged plates. The electric field accelerates them and they pass through a hole in the plates

d) Deflection – the positive ions then get deflected by a magnetic field placed at right angles to their path. The amount of deflection is proportional to the charge/mass ratio. Ions with a

smaller mass get deflected more than larger ions. Also, ions with a higher charge will be deflected more because they are more affected by the magnetic field.

e) Detection – the ions are detected and a signal is sent to a recorder. The strength of the signal is a measure of the number of ions with that particular charge mass ratio.

The mass spectrometer can give us the mass of individual atoms, which are all in the range of 10 -24g to 10-22g. As these numbers are so small, it makes sense for us to use relative values. Scientists chose carbon-12 as the standard, because it is common, easy to transport, and a solid. In other words, at atom of carbon-12 was given a mass of 12, and every other atom is compared to that.

Results of a mass spectrometer are presented as a mass spectrum:

e.g. mass spectrum for copper. Calculate the average mass

e.g. Boron, used as a control for nuclear reactors, exists as 10B and 11B. Use your periodic table to find the abundances of the two isotopes.

7. Electron arrangement

ACT – flame test

a) The electromagnetic spectrum

Electromagnetic radiation comes in different forms, from high-energy gamma rays to low-energy radio waves.

All waves travel at the speed of light (c) = 3x108 m s-1

Wavelength (λ)= the distance between successive crests

Frequency (ƒ) = the number of waves that pass a point in 1 s.

The higher the frequency, the higher the energy.

Related by equation c=λƒ

ACT: Spectroscopy lab

continuous spectra – when white light passes through a prism and all the wavelengths appear ROYGBIV

emission line spectrum – when high voltage is applied to a gas, lines that correspond with distinct wavelengths of light are produced. As each element has a different emission line spectrum, they are like barcodes that identify the element.

b) Bohr models

In Grade 10 you learned to the rules for drawing Bohr models for the first 20 elements. Here you will learn why we have those rules, and the limitations of the Bohr model, but first some background:

The photons (packets of light) of light that are emitted from an element to give a line spectrum come from the energy released when an electron falls from its excited state to its ground state.

Bohr used the line spectrum of hydrogen to propose that because the light is found in discrete wavelengths, the energy of the electrons is also quantized – meaning it exists at discrete energy levels and not in between. (staircase analogy)

Ephoton = ΔEelectron

Max Planck related the energy to the frequency of the radiation using ΔE=hf

(‘h’ = Planck’s constant = 6.626068 × 10-34 m2 kg / s)

Hydrogen spectrum: Transition to n=1 produces UV light

Transition to n=2 produces visible light

Transition to n=3 or higher produces infrared light

Ionization energy: the energy re quired to remove an electron from the ground state of each atom in one mole of gaseous atoms, ions or molecules

First ionization energy: the minimum energy needed to remove one mole of electrons from one mole of gaseous atoms in their ground state

Valence electrons: the outer electrons (assume ground-state)

Image from: http://www.chemguide.co.uk/atoms/properties/moreies.html

Key points:

c) Quantum MechanicsThose “shells” of the Bohr model are referred to as the principal energy levels, or principal quantum numbers. Experiments show there are sub-levels, referred to as s, p, d and f. These patterns only make sense if we treat electrons as waves instead of particles.

Wave-particle duality:

Light can be described by its frequency (a wave property) or by the E of the individual “particles” (called photons or quanta of light). Both properties are related by Planck’s equation E = hf

The same is true of electrons, or any matter, although the significance of the wave-like properties is more significant for subatomic particles. Electrons fired from an “electron gun” through a slit make a diffraction pattern, just like waves.

Heisenberg’s uncertainty principle: we cannot know where an electron is at any given moment. The best we can hope for is a “probability picture” of where the electron might be.

Atomic orbitals: a region around a nucleus where there is a 90% chance of finding an electron. An electron is characterized by its orbital and its spin

ACT – quantum leap lab

Principal Quantum Number , n

is roughly equivalent to the n of the energy levels in the Bohr atom. has whole number values n = 1, 2, 3, 4,… and designates what are called the main energy levels in an

atom. all orbitals with the same n value are said to be in the same shell. as n increases, electrons are generally further from the nucleus and have higher total potential

energy.

Subshell or Secondary Quantum Number , l

divides the shells into smaller groups of orbitals called subshells only certain values of l are allowed

n: l = 0, 1, 2, …, (n-1).

The total possible subshells in each level is equal to nFor n = l = 0, 1, 2, 3, …, (n-1) (subshell letters)

1 0(s)

2 0, 1(s, p)

3 0, 1, 2(s, p, d)

4 0, 1, 2, 3(s, p, d, f)

whereas the principal quantum number describes the energy and size of an orbital, the secondary quantum number determines the shape of the orbital

energies of the subshells within the same main energy level increase from s < p < d < f

Shapes: you should know the shapes of the s and p orbitals. See p. 332

Magnetic Quantum Number, m l (also known as Orbital Quantum Number)

splits the subshells into individual orbitals describes how an orbital is oriented in space relative to other orbitals restrictions on the value of ml can range from -l to +l

For l = ml = -l,…-1, 0, 1…, +1)

(total)

0 (s) 0

(one s orbital)

1 (p) +1, 0 -1

(three p orbitals)

2 (d) +2, +1, 0, -1, -2

(five d orbitals)

3 (f) +3, +2, +1, 0, -1, -2, -3

(seven f orbitals)

Spin Quantum Number, ms

electrons behave as tiny magnets and like a toy top can spin in one of two directions electron spin has two possible orientations

ms = +1/2, -1/2

Pauli exclusion principle – states that no two electrons in the same atom may have identical values for all four quantum numbers

Aufbau principle – electrons are placed into the lowest energy orbitals first.

e.g. Draw and write electron configurations for the first five elements

Hund’s Rule – when electrons are placed in a set of orbitals of equal energy, they are spread out as much as possible to give as few paired electrons as possible (separate orbitals and spins in the same direction)

8. Electron Configuration for Atoms

Note: Electron arrangement refers to the 2,8,8 etc of the Bohr models – the number of electrons in the principal energy levels only. Electron configuration refers to the specific orbitals.

e.g. write the full electron configuration for:

Argon

Fluorine

Magnesium

Energy level diagram:

As we get further from the nucleus, the energy levels get closer together. Notice that the 4s orbital will fill before the 3d.

http://users.stlcc.edu/gkrishnan/electron3.html

Block structure of the periodic table

The shape of our modern periodic table is explained by the electron configuration. If you are writing the electron configuration for an

element and finish with electrons in the p-orbitals, we call them “p-block elements.”

Use coloured pencils to identify the s, p, d, and f-block elements.

Understanding this will save you a lot of time when writing electron configurations!

Core notation: a shorthand notation for electron configuration. The “core” refers to the electron configuration of the noble gas that precedes the element you are dealing with. Write the core in square brackets, followed by the additional electrons.

e.g. Core notation for silicon: [Ne] 3s23p2

e.g. Core notation for iron:

e.g. Core notation for krypton:

Exceptions:

The 4s and 3d orbitals are so close together, that it results in a more stable atom if an e- jumps from the 4s orbital to either:

a) half-fill the 3d subshell. 3d4 3d5

or

b) completely fill the 3d subshell. 3d9 3d10

e.g. chromium

e.g. copper

9. Electron configuration for Ions

a) Anions – when adding electrons, simply follow the order of orbital filling given above

e.g. oxide

e.g. chloride

b) Cations – when removing electrons, remove from the highest principal energy level (n) first. For example, for transition metals, remove the 4s electrons before the 3d electrons.

e.g. Fe3+

e.g. S2-

10. Ionization energy

We have already learned about ionization energy as it relates to a single atom. Here you will compare first ionization energy for different elements.

ACT – Graphing Ionization Energy for the first 20 elements.

a) Explain the first ionization energy trend as you move across a period?

b) What are the exceptions?

c) What is the first IE trend as you move down a family?

d) How does the trend and the exceptions provide evidence for the existence of sub-shells?