Unit 2 – Atomic Structure & Nuclear ChemistryPart I – Atomic Theory, Subatomic Particles, and Average Atomic Mass

Part I Key Terms• Atomic mass - The mass of an atom of a chemical element expressed in atomic mass

units. It is approximately equivalent to the number of protons and neutrons in the atom (the mass number)

Average atomic mass – Weighted average of all atoms of a particular element and is dependent on the mass of isotopes for an element and the relative population of each isotope

• Bohr model - Devised by Niels Bohr, depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus

• Dalton’s Postulates - States that matter is composed of extremely small particles called atoms; atoms are invisible and indestructable; atoms of a given element are identical in size, mass, and chemical properties; atoms of a specific element are different from those of another element; different atoms combine in simple whole-number ratios to form compounds; in a chemical reaction, atoms are separated, combined, or rearranged

• Isotope -Atoms of the same element with different numbers of neutrons

Part I Key Terms (cont.)• Isotope notation - Subscripts and superscripts can be added to an element’s symbol to

specify a particular isotope of the element and provide other important information. The atomic number is written as a subscript on the left of the element symbol, the mass number is written as a superscript on the left of the element symbol

• Mass number - The total number of protons and neutrons in a nucleus.• Subatomic particles - The three kinds of particles that make up atoms: protons,

neutrons, and electrons • Theory - An explanation supported by many experiments; is still subject to new

experimental data, can be modified, and is considered valid if it can be used to make predictions that are proven true

Early Development of Atomic Theory

• Major Contributors to Understanding Atomic Structure• Democritus – ancient Greek philosopher that originally stated all

matter consists of atoms• 1605: Francis Bacon – published the scientific method• 1803: John Dalton – Postulates of Atomic Theory• 1897: J.J. Thomson – Discovery of the negatively charged

electron and the mass to charge ratio of the electron• 1908 Robert Millikan – Determines the charge of the electron• 1911: Ernest Rutherford – Discovers positively charged

nucleus• 1913: Niels Bohr – Theorizes structure of the electron

cloud with energy levels and planetary orbits of electrons• 1932: James Chadwick – Discovers neutrons

Atomic Theory – John Dalton• John Dalton

• English physicist• Experimented extensively with multiple gases and gaseous

compounds

• Contributions – Five Postulates of Atomic Theory• 1. All matter consists of tiny particles called atoms• 2. Atoms are indestructible and unchangeable.• 3. Elements are characterized by the mass of their atoms.• 4. When elements react, their atoms combine in simple, whole

number ratios.• 5. When elements react, their atoms sometimes combine in

more than one simple whole, number ratio.

Dalton’s Model of an Atom• He made no prediction about the construction of atoms

believing them to be solid spheres.

• Conclusions made based on his experiments and postulates:• Law of the Conservation of Mass – when chemical reactions

occur, the atoms are only rearranged and there is no difference in mass following a chemical reaction

• Law of Definite Proportions – elements combine in simple, low number ratios to form compounds (examples – H20, CO2)

• Law of Multiple Proportions –elements combine in different simple, low number ratios to form different compounds (examples – H20 and H202; CO and CO2)

Atomic Theory – J.J. Thomson• Discovered the negatively charged electron and the mass to

charge ratio of the electron

• Used cathode ray tube

• Beam of electrons deflected toward positive plate indicated the electron has negative charge

• Amount of deflection indicates the mass to charge ratio

Thomson’s Experiment

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Thomson’s Model of the Atom• Plum Pudding Model

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Atomic Theory – Ernest Rutherford• Discovered positively charged nucleus

• Used gold foil & detector ring

• Fired alpha particles at foil which are positively charged

• Most went through – atom mostly empty space

• Some deflected – nucleus positively charged

• Some bounced back – solid mass indicates nuclear core

Rutherford’s Experiment

Rutherford’s Model of the Atom

• Nuclear Atomic Model

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Atom Theory – Niels Bohr• Discovered electrons reside in energy levels with discrete

amounts of energy

• Mathematic modeling

• Needed to explain why negatively charged electrons do not get absorbed into positively charged nucleus

• Used information from Balmer, Lyman, & Paschen series

• Emission spectra for Hydrogen explained by Rydberg equation

Bohr’s Model of the Atom• Electron Shell Model

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2 Regions of the AtomNucleus

Contains the protons and neutronsAccounts for virtually all of the mass, but only a very small

portion of the volume of the atom.Has a positive charge equal to the number of protons.

Electron Cloud Contains the electrons in orbitalsHas virtually no mass, but accounts for virtually all of the volumeHas a negative charge equal to the number of electrons.

Subatomic Particles

ElectronsCharge = -1Mass ≈ 0 amuLocation: in orbitals in the electron cloud (outside the nucleus)

ProtonsCharge = +1Mass = 1 amuLocation: Inside the nucleus

NeutronsCharge = 0Mass = 1 amuLocation: Inside the nucleus

Properties of the AtomMass

Measured in Atomic Mass Units (amu)Equal to the sum of the number of protons and neutronsRepresented by the Mass Number

ChargeNeutral unless electrons gained or lost (ionized)Number of electrons and protons is equal and, therefore balance

out

Atomic NumberEqual to the number of protonsDefine the element and its chemical properties

SymbologyExample assuming neutral atom

of FluorineAtomic number: 9

Mass Number:19Protons: 9

Neutrons: 10(mass number – atomic number)

Electrons: 9

F919

Isotopes

• Atoms of the same element with different mass due to different number of neutrons

Average Atomic Mass

Weighted average of all atoms of a particular element

Dependent on the mass of isotopes for an element and the relative population of each isotope

% mass oxygen-16: (15.99491) (.99759) = 15.9564% mass oxygen-17: (16.99913) (.00037) = 0.0063% mass oxygen-18: (17.99916) (.00204) = 0.0367Average Atomic Mass of Oxygen = 15.9994

IsotopeIsotope Atomic

Mass (amu)Population (%)

Oxygen-16 15.99491 99.7590

Oxygen-17 16.99913 0. 037

Oxygen-18 17.99916 0.20400

Naming IsotopesName of the element followed by the mass number of the

isotope

Carbon – 12 = the name of the carbon atom with a mass number of 12 (6 protons and 6 neutrons)

Carbon – 14 = the name of the carbon atom with a mass number of 14 (6 protons and 8 neutrons)

Fluorine – 19 = the name of the Fluorine atom with a mass number of 19 (9 protons and 10 neutrons)

Energy Levels

• Energy levels correspond to the energy of individual electrons. Each energy level has a discrete numerical value.

• Different energy levels correspond to different numbers of electrons using the formula 2n2 where “n” is the energy level

Energy Level Number of electrons (2n2)

1 2(12) = 2

2 2(22)= 8

3 2(32)= 18

4 2(42)= 32

n 2n2

Quantum Mechanical Model of Atomic Structure• 1900: Max Planck – Develops law correlating energy to

frequency of light• 1905: Albert Einstein – Postulates dual nature of light as both

energy and particles• 1924: Louis de Broglie – Applies dual nature of light to all

matter• 1927: Werner Heisenberg – Develops Uncertainty Principle

stating that it is impossible to observe both the location and momentum of an electron simultaneously

• 1933: Erwin Schrodinger – Refines the use of the equation named after him to develop the concept of electron orbitals to replace the planetary motion of the electron

OrbitalsImpossible to determine the location of any single electron

Orbitals are the regions of space in which electrons can most probably be found

Four types of orbitalss – spherically shapedp – dumbbell shapedd – cloverleaf shapedf – shape has not been determined

Each additional energy level incorporates one additional orbital type

Each type of orbital can only hold a specific number of electrons

Orbital Types

Orbital Type

General Shape

OrbitalSublevels

# of electrons

per sublevel

Total # of electrons

per orbital type

s Spherical 1 2 2

p Dumbbell 3 2 6

d Clover leaf 5 2 10

f unknown 7 2 14

Electron Configuration

Energy Level

Orbital Type

OrbitalSublevel

# of orbitals

per energy level (n2)

# of electrons

per orbital type

# of electrons

per energy level (2n2)

1 s 1 1 2 2

2sp

13

426

8

3spd

135

926

1018

4

spdf

1357

16

26

1014

32

Electron Configuration Notation

• Find the element on the periodic table• Follow through each element block in order by stating the

energy level, the orbital type, and the number of electrons per orbital type until you arrive at the element.

1s

2s     2p

3s     3p

4s   3d 4p

5s   4d 5p

6s 4f 5d 6p

7s 5f 6d 7p

Samples of e- Configuration

• Element Electron Configuration• H 1s1

• He 1s2

• Li 1s2 2s1

• C 1s2 2s2 2p2

• K 1s2 2s2 2p6 3s2 3p6 4s1

• V 1s2 2s2 2p6 3s2 3p6 4s2 3d3

• Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 (Note the overlap)• Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2

Noble Gas Electron Configuration Notation

Find element on the Periodic Table of Elements• Example: Pb for Lead

Move backward to the Noble Gas immediately preceding the elementExample: Xenon

Write symbol of the Nobel Gas in bracketsExample: [Xe]

Continue writing Electron Configuration Notation from the Noble GasExample: [Xe] 6s2 4f14 5d10 6p2

Valence Electrons

• The electrons in the highest (outermost) s and p orbitals of an atom

• The electrons available to be transferred or shared to create chemical bonds to form compounds

• Often found in incompletely filled energy levels

Valence ElectronsShortcut to finding valence electrons for main group elements

Family 1A (1) 1 valence electronFamily 2A (2) 2 valence electronsFamily 3A (13) 3 valence electronsFamily 4A (14) 4 valence electronsFamily 5A (15) 5 valence electronsFamily 6A (16) 6 valence electronsFamily 7A (17) 7 valence electronsFamily 8A (18) 8 valence electrons

Family 3-12 have multiple possibilities and shortcuts do not work

Electron Dot NotationElectron configuration notation using only the valence electrons of an

atom.

The valence electrons are indicated by dots placed around the element’s symbol.

Used to represent up to eight valence electrons for an atom. One dot is placed on each side before a second dot is placed on any side.

Valance Electrons: Sodium Magnesium Chlorine Neon

1 2 7 8

Electron Dot Notation: • • •• ••

Na Mg : Cl : : Ne : • • ••

Oxidation Numbers:+1 +2 -1 0

Part II Key Terms• Alpha particle: A helium nucleus emitted by some radioactive

substances• Beta particle: An energetic electron or positron produced as the result

of a nuclear reaction or nuclear decay• Beta radiation: Radioactive decay in which an electron is emitted• Electron Configuration Notation -Consists of an element’s symbol,

representing the atomic nucleus and inner-level electrons, that is surrounded by dots, representing the atom’s valence electrons.

• Emission spectrum: The range of all possible wave frequencies of electromagnetic radiation, waves created by the systematic interactions of oscillating electric and magnetic fields

• Energy Levels - A certain volume of space around the nucleus in which an electron is likely to be found. Energy levels start at level 1 and go to infinity.

• Excited state: The state of an atom when one of its electrons is in a higher energy orbital than the ground state.

Part II Key Terms (cont.)• Gamma radiation: Electromagnetic radiation emitted during

radioactive decay and having an extremely short wavelength• Ground state: The lowest energy state of an atom or other particle• Nuclear fission: Splitting of the nucleus into smaller nuclei• Nuclear fusion: Combining nuclei of light elements into a larger

nucleus• Nucleon: a constituent (proton or neutron) of an atomic nucleus• Planck’s constant: As frequency increases, the energy of the wave

increases• Radioactive decay: Spontaneous release of radiation to produce a

more stable nucleus• Radioactive isotope: An isotope (an atomic form of a chemical

element) that is unstable; the nucleus decays spontaneously, giving off detectable particles and energy

Electromagnetic (EM) Spectrum

• The EM Spectrum is the range of all possible wave frequencies of electromagnetic radiation, waves created by the systematic interactions of oscillating electric and magnetic fields

• The general term for all electromagnetic radiation is light• The range of the EM Spectrum is from very low frequency

known as radio waves to very high frequency known as gamma radiation

• The visible spectrum of light is in the center portion of this EM Spectrum

• All EM Spectrum travels at the same speed in a vacuum – this speed is known as the speed of light, 3.00 x 108 m/s

EM Spectrum

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Speed of Light and Frequency• Since the speed of all EM radiation is the same, there is a clear

mathematical relationship between the frequency of the light and its wavelength

• All waves travel at a speed that is equal to the product of its frequency (the reciprocal of time) and its wavelength (distance)

c = f λ• The speed of EM radiation is fixed at 3.00 x 108 m/s

• Therefore:3.00 x 108 m/s = f λ

Speed of light = frequency x wavelength• As frequency increases, wavelength decreases. As wavelength

increases, frequency decreases• Example: If frequency doubles, wavelength is cut in half

As f ↑, ↓λ : Calculations• If the wavelength of a radio wave is 15 meter, what is its

frequency?3.00 x 108 m/s = f (10 m)

(3.00 x 108 m/s) / 15 m = f2.0 x107 s-1 = f

Frequency = 2.0 x107 Hertz

• If the frequency of gamma radiation is 6.25 x 1022 Hertz, what is its wavelength?

3.00 x 108 m/s = (6.25 x 1022 s-1) λ(3.00 x 108 m/s) / (6.25 x 1022 s-1) = λ

4.80 x10-15 m = fWavelength = 4.80 x10-15 m

Planck’s Law• Max Planck determined in 1900 there was a mathematical

relationship between the energy of EM radiation and the frequency of that radiation:

As frequency increases, the energy of the wave increases

E = h f

Energy = Planck’s constant x frequencyE = (6.63 x 10-34 Joule seconds) f

Planck’s Law Calculations• Example: If the wavelength of green light is 5.21 x 10-7 meters,

what is the energy of this light?

3.00 x 108 m/s = f (5.21 x 10-7 m)(3.00 x 108 m/s) / 5.21 x 10-7 m = f

5.76 x1014 s-1 = fFrequency = 5.76 x1014 Hertz

E = (6.63 x 10-34 Joule seconds) (5.76 x1014 s-1)E = 3.82 x10-19 Joules

Implication of Planck’s Law• In order to move an electron to a higher energy level, excite an

electron, energy must be absorbed to move the electron• Since electrons exist in fixed energy levels with a specific

amount of energy, the amount of energy needed is a finite amount equal to the difference in the energy associated with the ground state of the electron and the energy associated with the level to which the electron is excited

• If the energy related to the excited electron is removed, the electron will return to its ground state and the energy released is equal to the energy absorbed to excite it

• The energy released is released as light• The overall result is that every element has a unique spectra of

light associated with it and the spectra can be used to identify the element

All nuclear reactions are based on Einstein’s Theory of Relativity

At speeds approaching the speed of light, energy and mass are interchangeable

E = mc2 Energy = mass x (speed of light)2

Mass can be converted to energy and vice versa

Nuclear Reactions

There is a difference between the mass of an atom and the various particles that make up the atom

This difference is called the mass defect of the atom

This mass defect is the binding energy of the atom

In nuclear reactions, the binding energy is released as energy (heat, light, or gamma radiation) and/or particles with measureable mass

Mass Defect

• Fission – Splitting of the nucleus into smaller nuclei

• Fusion – Combining nuclei of light elements into a larger nucleus

• Radioactive Decay – Spontaneous release of radiation to produce a more stable nucleus

Types of Nuclear Reactions

Nucleus splits into smaller nuclei when struck by a neutron of sufficient energy

Tremendous release of energy

When controlled can produce huge amounts of power in nuclear reactors

Naturally occurs in uranium and other ores in spontaneous fission

Clean source of energy with no carbon footprint

Produces radioactive nuclear waste with long term environmental and health considerations

Fission

Fission Process

Fission and Nuclear Reactors

Lighter nuclei (such as hydrogen) combined to form heavier nuclei

Tremendous release of energy

2H + 3H 4He + 1n + energy Deuterium Tritium Helium(occurs naturally in water)

Powers the sun and stars

No practical application to produce usable energy at this time

Fusion

Fusion Process

• Spontaneous release of radiation by unstable nuclei in order to increase stability

• Radiation can be either energy alone (gamma) or energy accompanied by release of a particle (all of the other forms of decay)

Radioactive Decay

Alpha decay – release of alpha particle and energy

Beta decay – release of beta particle and energy

Gamma Emission – release of electromagnetic radiation (energy)

Positron Emission – release of a positron and energy

Electron Capture – absorption of and electron and release of energy

Neutron Emission – release of a free neutron and energy

Forms of Radioactive Decay

Typically found in heavier nuclei and the means to achieve stability is to reduce mass

Nuclei shed mass in the form of a helium nucleus to become more stable

Helium nucleus that is released is ionized and called and Alpha Particle

Alpha Decay

Alpha Particle is positively charged (no electrons present)

Alpha Particles are very massive, but travel slower (low penetrating power)

Can cause significant tissue damage if not shielded

Shielding can be accomplished with clothing or paper

Alpha Decay (cont.)

Alpha Decay Process

Common in nuclei of any size where instability is caused by the number of neutrons

Neutron decays into a proton and an electron

Proton remains in the nucleus

The electron leaves the atom and is called a Beta Particle

Beta Decay

Beta Particle is negatively charged

Mass of the nucleus is unchanged

Beta particles have very low mass but are travelling at very high speed

Beta particles can penetrate through the skin and cause deep tissue damage

Beta Decay (cont.)

Beta Decay Process

Nucleus becomes more stable through the release of electromagnetic energy

No change in mass

No change in the element

The Gamma radiation can be reduced by shielding, but Gamma radiation cannot be stopped

Usually found with another type of decay, but not always

Gamma Radiation

Radioactive Decay Type

Mass Charge Penetrating Power

Transmutation

Alpha 4 amu Positive Low New Element Formed

Beta 0 amu Negative High New Element Formed

Gamma None (no particle)

None Extremely High No

Radioactivity Decay Comparison

• All nuclear reactions must conserve the overall mass of the particles involved in the reaction

• Two properties must be the same on both sides of a nuclear equation

• Total Mass Number – the sum of the mass numbers of all particles must be the same on both sides of the reaction

• Total Atomic Number – the sum of the atomic numbers of all particles must be the same on both sides of the reaction

Nuclear Reaction Mass Conservation