Unit 3 Atomic Structure and Nuclear Chemistry

Unit 3 Notes Page 1 of 16

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3.1 ATOMIC STRUCTURE

Historical Development of Atomic Theory

With No scientific method, the Greek philosopher __Democritus______first used the term __ATOM__ to describe the smallest, indivisible unit of matter in around 400 BCE. Almost 2000 years later…

1803John DaltonFirst Atomic Model

Matter is made of indivisible particles called __Atoms__
Atoms of one element are __identical____
Atoms of different elements are __different____
The atom is a solid ___indivisible______mass.

1897J.J. ThomsonPlum Pudding Model

Identified the ___electron_____ as a particle
Used a Crooke’s tube to examine electrons
__plum__-___pudding___ model
Atom is a clump of __positively___ charged material

(pudding) with electrons scattered throughout (plums)

1911Ernest RutherfordNuclear Model

__Gold__ __Foil___ experiment
Shot particles through paper thin gold foil
Most passed thru (atom is mostly _empty space____)
Very few deflected greatly (dense + charged __nucleous___)

1913Neils BohrBOHR Model

a.k.a “planetary” model
electrons are arranged in concentricorbits (like rings)

around the sun

electrons have fixed ___orbitals_____
an energy level is the region around the nucleus where

electrons are moving

1925Quantum Mechanical Model

currently accepted model
first proposed by Werner Heisenberg
Many physicists & chemists contributed to model
Mathematical model derived by Max Schrödinger
the __electron___ __cloud_____ is the space

where probability of finding electron is high

Other notable discoveries related to Atomic Theory……..

1897Marie Curie Radioactivity

Investigated radiation and 1stperson to use term “radioactivity”
Proved that atom is not stable, contrary to common belief at time
Isolated radioactive elements including radium (0.1 g from 1000 kg)
Shared two Nobel prizes for her work (1st women to win nobel prize)

1932James ChadwickDiscovery of Neutron

Researchers saw that mass of nucleus greater than mass of protons
Idea of neutral particle first proposed by Ernest Rutherford
Chadwick used Curie’s method of detecting particles and identified neutron

Atomic Structure

An atom is the ___most basic_____ (smallest unique) unit of matter.

There are two regions of an atom that contain particles of matter, the rest is empty space.

The nucleus, at the CENTER of the atom, holds:

PROTONS ( _+__ charge) and;
NEUTRONS ( __0__ charge)

Theelectron cloud isa region SURROUNDING the nucleus where ELECTRONS ( _-_ charge) are found.

How Atoms Differ – Atomic Number and Mass Number

Label Hydrogen’s entry on the Periodic Table.

Atomic number

symbol

average atomic mass

elements name

TheATOMIC NUMBERis the number of PROTONSin an atomand:

Is unique to each element
Is THE SAME for all atoms of an element

IDENTIFIES an element.
In a neutral atom (equal # of negative and positive particles), the # of ___electons___ IS EQUAL TO the # of ___protons______.

The MASS NUMBERof an element is the number of PROTONS plus the number of NEUTRONSin an atom and is the same as the mass of the __atom___

Atoms of the same elements can have different number of neutrons and these are called ISOTOPES and have a distinct Mass Number.

Atomic Mass Units

The mass of atoms is measured in _amu__, or atomic mass units.

1 amu = the mass of 1 atom of carbon(carbon with 6 protons and 6 neutron and therefore mass # of 12)

Fill in the missing information about each subatomic particle:

Particle / Charge / Where
found? / Mass
(amu) / In one element,
can the # vary?
proton / + / Nucleus / 1 / No!
electron / – / e-1 cloud / Ca. 0 / Yes, ions!
neutron / 0 / nucleus / 1 / Yes, isotopes!

Fill in the following information about the selected atoms:

Element / Symbol / Atomic # / Mass # / # of protons / # of neutrons / # of electrons
Sodium / Na / 11 / 23 / 11 / 12 / 11
Flourine / F / 9 / 19 / 9 / 10 / 9
Selenium / Se / 34 / 79 / 34 / 45 / 34
Chromium / Cr / 24 / 52 / 24 / 28 / 24
Gallium / Ga / 31 / 70 / 31 / 39 / 31

Shorthand Notation

Shorthand notation allows us to write a single isotope simply. When shorthand notation is used, it will appear one of the following ways:

Example: Bromine atom with a mass number of 80 amu can be written:

Bromine has an atomic number of 35. The 80, above, is the mass number of this atom of bromine. SO, we now know that this bromine isotope has 35 protons and 45 neutrons. *79.90 on the periodic table is the average mass of all known Bromine atoms.

You will also see isotopes written in this format: Flourine-19. In this example, Flourine-19refers to the isotope of fluorine that has an atomic mass of 19, i.e. 9 protons and 10 neutrons.

Practice: Write the shorthand notation for…

1)Neon – 222) Potassium – 413) Chlorine – 36

3.2 ISOTOPES AND AVERAGE ATOMIC MASS

Isotopes

Isotopes are atoms of the same element that have different numbers of ___neutrons (no)_____.

This means isotopes have different atomic masses, but the same atomic number
Isotopes of an element are chemically the same (because neutrons are neutral).
All elements have isotopes.
Every element found in nature is a mixture of all its isotopes

Example: Three isotopes of potassium

Potassium – 39 / Potassium – 40 / Potassium – 41
P+ / 19 / P+ / 19 / P+ / 19
E– / 19 / E– / 19 / E– / 19
N0 / 20 / N0 / 21 / N0 / 22

Average Atomic Mass

Average atomic mass is a weighted average of all isotopes of an element. The percent of each isotope in an element (all known atoms) is called its PERCENT ABUNDANCE. Every isotope has its own percent abundance.

Example:Nitrogen has two naturally occurring isotopes, nitrogen-14 and nitrogen-15. The average atomic mass of nitrogen is 14.007 amu. Which isotope is more abundant in nature? 14N – closer to average mass unit (amu)

CalculateAverage Atomic Mass in a 3 step process.

Example:lithium-7 (mass = 7.016 amu, 92.41%)

lithium-6 (mass = 6.015 amu, 7.59%)

Step 1:Change the percent abundance for each isotope to a decimal.

(Move decimal 2 places to left to convert from percent to decimal)

lithium-7 = 92.41% 0.9241lithium-6 = 07.59% 0.0759

Step 2:Multiply each abundance value by the mass of the isotope. The product is called relative mass.

9241 x 7.016 = 6.483 amu.0759 x 6.015 = 0.457 amu

Step 3:Add the relative masses to find average atomic mass. Units are amu.

6.483 + 0.457 = 6.940 amu

Example:Find the average atomic mass of boron.

boron-10 (%abundance = 19.8% and mass = 10.013 amu)

boron-11 (%abundance = 80.2% and mass = 11.009 amu)

0.198 x 10.013 = 1.98

0.802 x 11.009 = 8.83

______

10.81 amu

Example:Silver is found in nature in the following percentages:

Ag = 51.82%Ag = 48.18%

Calculate the average atomic mass of Silver.

0.5182 x 107 = 55.45

0.4818 x 109 = 52.52

______

107.97 amu

Practice:Rubidium has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium?

0.722 x 85 = 61.4

0.278 x 87 = 24.2

______

85.6 amu

3.3 ISOTOPE STABILITY AND NUCLEAR DECAY

In reality, all atomswill eventually break apart, given enough time. The time required for half of a sample of one isotope to break apart (spontaneously decay) is called its half-life.

Some isotopes have a half-life of seconds; others have a half-life ofbillions of years (longer than the age of the universe!).

When a nucleus decays, energy, and often particles (protons, neutrons and/or electrons) are ejected from the nucleus.

PREDICTING ISOTOPE STABILITY

An isotope is considered____STABLE______if the nucleus will NOT spontaneously decay. An isotope with an unstable nucleus is called a radioisotope.

Elements with atomic # __1-20_____have at least one isotope that isvery stable
1:1 ratio of proton to neutron (p+ : n0)
Example: Carbon-12 has 6 p+ and 6 n0

Elements with atomic #____21-82_____have at least one isotope that issomewhat stable (still stable!)

2:3 ratio of protons to neutrons (p+ : n0)
Example: Mercury-200 has 80 p+ and 120 n0

Elements with atomic # ___>/= 83______do not have a stable isotope and are unstable AND radioactive
1: >2 ratio of protons to neutrons (p+ : n0)
Examples: Uranium (U) and Plutonium (Pu)

The Band of Stability

NUCLEAR DECAY

An unstable nucleusdecays because it has a number of neutrons, either too many or not enough, that makes the nucleus unstable. The decaying nucleus emits energy as particles and rays and transmutates into a more stable isotope of a different element. There are many types of decay.

1. Alpha () Decay – emission of an alpha particle, denoted by the symbol to the RIGHT

because contains__2_ protons __2__ neutrons (like a Helium nucleus).

The charge is __+____ because it has ___2____ protons.

Alpha decay ____decreases____the mass number by __4__ and the atomic number by __2___.
There are NO electrons in an alpha particle
All nuclear equations are balanced

Example:Write the nuclear equation for the radioactive decay of polonium-210 (Po) by alpha emission.

2104206

Po He+ Rn

84282

Practice:Write the balanced nuclear equation for the alpha decay of radium-226.

2264222

Ra He+ Rn

88286

2. Beta ()Decay – emission of a beta particle, a fast-movingelectron given by the symbols at right.

particles have insignificant mass, so mass # = 0

 decay results from the conversion of a neutron into a proton in the nucleus. In this process, a high speed electron is ejected from the nucleus.
The charge of the particle is __-1__ (just like an electron)
Beta decay causes ____No____ change in the mass number.
The atomic number _____increases______by 1.

Example:Write the nuclear equation for the radioactive decay of carbon-14 by beta emission.

14 0 14

C   + N

6 -1 41

Practice: Write the balanced nuclear equation for the reaction in which zirconium-97 undergoes beta decay.

97 0 97

Zr   + Nb

40 -1 41

3. Gamma (γ)Emission – high-energy ELECTROMAGNETIC RADIATION

denoted by the symbol at right. No particles included, only energy, so no change in

contents of nucleus.

Charge is ____0______.
__No____ effect on mass number or atomic number, so not included in nuclear reactions.
Gamma rays always accompany alpha and beta radiation.

Uses of Radioactive Isotopes

All three types of radiation are used beneficially in the following ways:

Medical imaging, treatment, research and diagnostics
Food irradiation to kill harmful bacteria
Smoke detectors
Biological research and studies
Insecticides
Energy Production
Numerous Industrial Applications

transmutation – the conversion of an atom of one element to an atom of another element (radioactive decay is one way that this occurs!)

Properties of Alpha and BetaParticles and Gamma Radiation

Alpha () / Beta () / Gamma ()
Composition / Helium nucleus
2p+, 2no / High energy
electron / High-energy electromagnetic radiation
Charge / + / – / 0
Change in Mass
Number / Decrease by _4__ / no change / no change
Change in Atomic Number / Decrease by _2__ / Increase by _1_ / no change
Mass (amu) / 4 / / 0
Tissue Penetrating power
(depth of travel) / Low
(0.05 mm) / Moderate
(4 mm) / Very High
(penetrates entire body easily)
Shielding
(to stop progress of radiation) / Sheet of paper / Wood
Metal foil / Lead
Concrete

3.4 NUCLEAR REACTIONS

In a NUCLEAR reaction, the following will occur…

isotopes of one element are CHANGED into isotopes of another element (__transmutation______)
contents of the nucleus change
__Large___ amounts of energy are released

There are FOUR types of nuclear reactions.

Radioactive Decay – alpha decay, beta decay, and gamma electromagnetic radiation

FISSION – _____Splitting______a nucleus

A very ____Large______nucleus is split into two large fragments by a fast moving neutron.

The reaction releases lots of __energy___ and many __neutrons______which split more nuclei

Above: Fission of Uranium 235

If controlled, energy is released ___slowly_____ like in a nuclear reactor, and can be turned into electricity.
If not controlled or control is lost, a nuclear explosion or reactor meltdown can occur
1st controlled nuclear reaction – 1942 (Chicago Pile-1 created by Enrico Fermi)
1st atomic bomb explosion – 1945 (Trinity Bomb Test in White Sands, NM)

FUSION –___Combining______of nuclei

two ____small___ nuclei combine to form single larger nucleus

Does NOT occur under standard conditions, positively charged Hydrogen atoms __repel__ each other.
advantages (compared to fission) - inexpensive, no radioactive waste
disadvantages - requires _large______amounts of energy to start reaction and is difficult to control
examples – energy output of stars, modern thermonuclear weapons (hydrogen bombs), future nuclear reactors

Nuclear Disintegration – Emission of a __proton_____ or a ___neutron______. Occurs when very small particles hit a nucleus with enough energy to remove particles.