Name ___KEY__________ Period ___

CRHS Academic Chemistry

Unit 3 Atomic Structure and

Nuclear Chemistry

NOTES

Cr

Key Dates

Quiz Date _______ Exam Date _______

Lab Dates ________ __________

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

52 24

Mass Number

Symbol

Atomic Number

Page 2 of 16 Unit 3 Notes

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…

1803 John Dalton First 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.

1897 J.J. Thomson Plum 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)

1911 Ernest Rutherford Nuclear 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___)

1913 Neils Bohr BOHR Model

a.k.a “planetary” model

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

electrons have fixed ___orbitals_____

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

Unit 3 Notes Page 3 of 16

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 __electron___ __cloud_____ 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

Page 4 of 16 Unit 3 Notes

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)

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.

Atomic number

symbol

average atomic mass

elements name

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 ___electons___ IS EQUAL TO the # of

___protons______.

1

H 1.008

Hydrogen

Unit 3 Notes Page 5 of 16

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 __atom___

o 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 = 1

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:

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

Q: Why don’t electrons get counted in the mass of an atom?

A: The mass of an electron is negligible, about __2000_____ times

smaller, when compared to the mass of a proton or a neutron, so

electron mass is not counted in the mass number. Electron is a drop and

proton is gallon

Page 6 of 16 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:

𝐵𝑟3580 or 𝐵𝑟0

80

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-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

35

Br 79.90

Bromine

____Average_____

atomic __mass___

Br 80

35

PERIODIC TABLE (applies to all Bromine atoms)

SHORTHAND notation (applies to one Bromine isotope)

Atomic number

___Mass__ number

Ne

or

Ne

K

or

K

Cl

Cl

22

10

41

19

41 22

36

17

36

Unit 3 Notes Page 7 of 16

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

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)

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

Potassium – 39 Potassium – 40 Potassium – 41

P+ 19 P+ 19 P+ 19

E– 19 E– 19 E– 19

N0 20 N0 21 N0 22

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 ___average___ atomic mass.

Page 8 of 16 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)

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

107 47

109 47

Unit 3 Notes Page 9 of 16

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____STABLE___________ if the nucleus will NOT spontaneously decay. An isotope with an

unstable nucleus is called a radioisotope.

o Elements with atomic # __1-20_____ have at least one isotope that is very stable

1:1 ratio of proton to neutron (p+ : n0)

Example: Carbon-12 has 6 p+ and 6 n0

o Elements with atomic # ____21-82_____ have at least one isotope that is somewhat stable (still stable!)

2:3 ratio of protons to neutrons (p+ : n0)

Example: Mercury-200 has 80 p+ and 120 n0

o 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

Page 10 of 16 Unit 3 Notes

4

He

2

0 0

or e

-1 -1

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 __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. 210 4 206 Po He + Rn 84 2 82

Practice: Write the balanced nuclear equation for the alpha decay of radium-226. 226 4 222 Ra He + Rn 88 2 86

2. Beta () Decay – emission of a beta particle, a fast-moving electron 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.

Unit 3 Notes Page 11 of 16

0

γ 0

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!)

Page 12 of 16 Unit 3 Notes

Properties of Alpha and Beta Particles 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 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

Unit 3 Notes Page 13 of 16

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.

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

2. FISSION – _____Splitting____________ a nucleus

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

b. The reaction releases lots of __energy___ and many __neutrons_______ which split more nuclei

Above: Fission of Uranium 235

c. If controlled, energy is released ___slowly_____ 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)

𝑈 + 𝑛𝑒𝑢𝑡𝑟𝑜𝑛01 → 𝐾𝑟 + 𝐵𝑎 + 𝑛𝑒𝑢𝑡𝑟𝑜𝑛𝑠0

356

1443689

92235 + 𝑒𝑛𝑒𝑟𝑔𝑦

Page 14 of 16 Unit 3 Notes

3. 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

4. Nuclear Disintegration – Emission of a __proton_____ or a ___neutron________. Occurs when very small particles

hit a nucleus with enough energy to remove particles.

𝐻 + 𝐻 → 𝐻𝑒 + 𝑛𝑒𝑢𝑡𝑟𝑜𝑛01 + 𝑒𝑛𝑒𝑟𝑔𝑦2

413

1 2

Above: Fusion of Deuterium and Tritium

Unit 3 Notes Page 15 of 16

Page 16 of 16 Unit 3 Notes