Name ______________ Period ___ CRHS Academic Chemistry

Unit 3 Atomic Structure and Nuclear Chemistry NOTES

52 Cr 24

Mass Number Symbol Atomic Number

Key Dates Quiz Date

_______

Lab Dates ________

Exam Date _______ __________

Notes, Homework, Exam Reviews and Their KEYS located on CRHS Academic Chemistry Website: https://cincochem.pbworks.com

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Unit 3 Notes

3.1 ATOMIC STRUCTURE Historical Development of Atomic Theory With _______scientific method, the Greek philosopher _______________ first used the term ________ to describe the smallest, indivisible unit of matter in around 400 BCE. Almost 2000 years later…



1803

John Dalton 

Matter is made of indivisible particles called _________ 

Atoms of one element are _________________



Atoms of different elements are _________________





First Atomic Model

The atom is a solid ________________________ mass.

1897

J.J. Thomson 

Plum Pudding Model

Identified the _____________________ as a particle 

Used a Crooke’s tube to examine electrons



__________-_______________ model



Atom is a clump of _______________ charged material (pudding) with electrons scattered throughout (plums)



1911

Ernest Rutherford 





Nuclear Model

__________ ___________ experiment 

Shot particles through paper thin gold foil



Most passed thru (atom is mostly ________________)

Very few deflected greatly (dense + charged ___________)

1913

Neils Bohr 

BOHR Model

a.k.a “planetary” model





electrons are arranged in concentric orbits (like rings) around the sun



electrons have fixed ___________________ an energy level is the region around the nucleus where electrons are moving

Unit 3 Notes



1925

Quantum Mechanical Model 

currently accepted model 



first proposed by Werner Heisenberg



Many physicists & chemists contributed to model



Mathematical model derived by Max Schrödinger

the _________________ ______________ is the space where probability of finding electron is high

Other notable discoveries related to Atomic Theory……..



1897

Marie Curie

Radioactivity

 Investigated radiation and 1st person 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)



1932

James Chadwick

Discovery 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

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Unit 3 Notes

Atomic Structure An atom is the __________________ (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 ( _____ charge)

The electron cloud is a 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. 1

H 1.008

Hydrogen

 The ATOMIC NUMBER is the number of PROTONS in an atom and: o

Is unique to each element

o

Is THE SAME for all atoms of an element

o

IDENTIFIES an element.

o

In a neutral atom (equal # of negative and positive particles), the # of __________________ IS EQUAL TO the # of ___________________.

Unit 3 Notes

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 The MASS NUMBER of an element is the number of PROTONS plus the number of NEUTRONS in an atom and is the same as the mass of the ___________ o

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

Q: Why don’t electrons get counted in the mass of an atom? A: The mass of an electron is negligible, about ______________ times smaller, when compared to the mass of a proton or a neutron, so electron mass is not counted in the mass number.

Atomic Mass Units

The mass of atoms is measured in ______, or atomic mass units. 1

1 amu = 12 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: Where Particle Charge found?

Mass (amu)

In one element, can the # vary?

proton

+

electron

–

Yes, ions!

neutron

0

Yes, isotopes!

Fill in the following information about the selected atoms: Element Symbol Atomic # Mass # Sodium

Na

11

23

9

19

Se

# of protons

# of neutrons

# of electrons

11

12

11

45

Chromium

28 Ga

70

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Unit 3 Notes

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: 80 35𝐵𝑟

80 0𝐵𝑟

or

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.

___________ number 35

Br 79.90

Atomic number ________________

Bromine

80 35

Br

atomic __________

PERIODIC TABLE

SHORTHAND notation

(applies to all Bromine atoms)

(applies to one Bromine isotope)

You will also see isotopes written in this format: Flourine-19. In this example, Flourine-19 refers 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 – 22

2) Potassium – 41

3) Chlorine – 36

Unit 3 Notes

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3.2 ISOTOPES AND AVERAGE ATOMIC MASS Isotopes Isotopes are atoms of the same element that have different numbers of __________________. 

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+

P+

P+

E–

E–

E–

N0

N0

N0

Average Atomic Mass Q: Why aren’t the masses listed on the periodic table whole numbers and why don’t they match the mass numbers we have been using? A: Since ALL elements exist as many different isotopes (with different mass numbers), the mass on the periodic table is the ________________________ 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?

Calculate Average 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%  0.9241

lithium-6 = 07.59%  0.0759

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Unit 3 Notes

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)

Example:

Silver is found in nature in the following percentages: 107 Ag 47

= 51.82%

109 Ag 47

= 48.18%

Calculate the average atomic mass of Silver.

Practice:

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

Unit 3 Notes

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3.3 ISOTOPE STABILITY AND NUCLEAR DECAY In reality, all atoms will 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 of billions 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___________________________ if the nucleus will NOT spontaneously decay. An isotope with an unstable nucleus is called a radioisotope. o

Elements with atomic # ______________ have at least one isotope that is very stable  

o

Elements with atomic # ______________ have at least one isotope that is somewhat stable (still stable!)  

o

1:1 ratio of proton to neutron (p+ : n0) Example: Carbon-12 has 6 p+ and 6 n0

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

Elements with atomic # _______________ 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

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Unit 3 Notes

NUCLEAR DECAY An unstable nucleus decays 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 ____ protons & _____ neutrons (like a Helium nucleus). 

The charge is ________ because it has _________ protons.



Alpha decay ____________________the mass number by _______ and the atomic number by _______.



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.

Practice:

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

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

OR

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 _______ (just like an electron)



Beta decay causes ____________ change in the mass number.



The atomic number ____________________________ by 1.

Unit 3 Notes

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Example:

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

Practice:

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

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 ___________________.



__________ 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!)

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Unit 3 Notes

Alpha ()

Beta ()

Gamma ()

Helium nucleus 2p+, 2no

High energy electron

High-energy electromagnetic radiation

+

–

0

Change in Mass Number

Decrease by ____

no change

no change

Change in Atomic Number

Decrease by ____

Increase by ____

no change

4

1 1837

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

Composition Charge

Mass (amu)

Properties of Alpha and Beta Particles and Gamma Radiation

Unit 3 Notes

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3.4 NUCLEAR REACTIONS In a NUCLEAR reaction, the following will occur… 

isotopes of one element are CHANGED into isotopes of another element (_________________________________)



contents of the nucleus change



_________________________ amounts of energy are released

There are FOUR types of nuclear reactions. 1. Radioactive Decay – alpha decay, beta decay, and gamma electromagnetic radiation

2. FISSION – ______________________________ a nucleus a. A very ________________________ nucleus is split into two large fragments by a fast moving neutron.

235 92𝑈

+ 10 𝑛𝑒𝑢𝑡𝑟𝑜𝑛 →

89 36𝐾𝑟

+

144 56𝐵𝑎

+ 30 𝑛𝑒𝑢𝑡𝑟𝑜𝑛𝑠 + 𝑒𝑛𝑒𝑟𝑔𝑦

b. The reaction releases lots of _______________ and many __________________ which split more nuclei

Above: Fission of Uranium 235

c. If controlled, energy is released ________________ like in a nuclear reactor, and can be turned into electricity. d. If not controlled or control is lost, a nuclear explosion or reactor meltdown can occur e. 1st controlled nuclear reaction – 1942 (Chicago Pile-1 created by Enrico Fermi) f.

1st atomic bomb explosion – 1945 (Trinity Bomb Test in White Sands, NM)

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Unit 3 Notes

3. FUSION –________________________________ of nuclei



two ______________ nuclei combine to form single larger nucleus 2 1𝐻

+ 31𝐻 → 42𝐻𝑒 + 10 𝑛𝑒𝑢𝑡𝑟𝑜𝑛 + 𝑒𝑛𝑒𝑟𝑔𝑦



Does NOT occur under standard conditions, positively charged Hydrogen atoms ______________ each other.



advantages (compared to fission) - inexpensive, no radioactive waste



disadvantages - requires ______________ amounts of energy to start reaction and is difficult to control



examples – energy output of stars, modern thermonuclear weapons (hydrogen bombs), future nuclear reactors

Above: Fusion of Deuterium and Tritium

4.

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

Unit 3 Notes

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Unit 3 Notes