52 Cr Mass Number 24 - Cinco Ranch Academic … Is unique to each element o Is THE SAME for all...
Transcript of 52 Cr Mass Number 24 - Cinco Ranch Academic … Is unique to each element o Is THE SAME for all...
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