Download The Structure of the Atom- Chapter 4, 3

Unit 2 Atomic Structure & Nuclear Chemistry
Section 1
Obj.2
Counting Subatomic Particles
The ATOM is defined as the smallest particle of an element that retains the properties of that element.
Subatomic Particles & Their Properties
Particle
Symbol
Location
Relative Mass
Actual mass in
Electrical Charge
in a.m.u*
grams
1 amu
+1
1.6710-24 g
+
p
1 amu
1.6710-24 g
0
.00055 amu
9.1110-28 g
-1
n0
e*a.m.u : atomic mass unit
1 amu (“atomic mass unit”) = 1.67  10-24 g
**The mass of the proton and neutron are very similar**
**The mass of the proton & neutron are both about 2000 times larger than the mass of the electron**
What are the 2 regions of the atom as of now?
 Nucleus: small, dense center containing ___________________ and _____________________.
 Electron Cloud: region surrounding nucleus that contains the electrons. It’s mostly empty space.
Counting Subatomic Particles
Atomic Number
 Every atom has a different number of protons,
which is represented by the atomic number.
 The number of protons determines the identity
of the atom
Atomic Number = Number of protons
Ex: Nitrogen’s atomic number is 7
Calculating Electrons in a Neutral Atom (A.P.E)
 A neutral atom has the SAME number of
protons and electrons.
# protons = # electrons if neutral
Mass Number
 The number of protons and neutrons in a
nucleus of a specific atom is called the mass
number.
Calculating Electrons in a Ion
2 Options:
[1] Add the number of given protons & neutrons
together
Mass # = protons + neutrons
[2] USE ONLY when no element information has
been given: Round atomic mass to a whole
number to get an element’s most likely most
common mass number.
o
Nitrogen’s atomic mass = 14.01 amu
o
Nitrogen’s mass number is
____________
Calculating Neutrons (M.A.N)
 To calculate the number of neutrons, subtract
the atomic number from the mass number
 Mass number – Atomic number = # Neutrons
Nitrogen:
14 – 7 = 7 neutron
 An ion is an atom with a charge. The number of
protons is NOT EQUAL to the number of
electrons.
 Positive charged ions called CATIONS have
more protons than electrons & negatively
charged ions are called ANIONS have more
electrons than protons.
Example:
[1] How many electrons does Br-1 have?
[2] How many electrons does Al+3 have?
Overall Charge = Protons – Electrons
Element Information:
Nuclear Symbol
LABEL THE PARTS BELOW
A
A
Z
Z
X
X
C
C
SELF CHECK
What are the number of protons, neutrons & electrons in each atom?
Hyphen Notation: LABEL THE PARTS!
Copper – 65 ___________
__ p __e __ n
You Try: (Be Careful: Never use the given atomic mas on the periodic table to calculate mass number unless
no other information has been provided. )
Nuclear
Symbol
Hyphen
Notation
Atomic #
Magnesium-25
Mass #
Charge
Proton
Neutron
Electron
126
82
+2
82
Section 2
Obj.2 cont. & Obj. 4
How Do Atoms Differ: Isotopes
 Isotopes are atoms of the same element that have the SAME NUMBER OF PROTONS but a DIFFERENT
NUMBER OF NEUTRONS
 Most elements contain a mixture of 2 or more isotopes. Each one having its own mass and abundance.
 Some isotopes are radioactive and unstable. (See nuclear chem. Section)
 You can identify an isotope by the different mass number of the same element
Carbon-12
Carbon-13
12
13
C
6
C
6
Fill in the chart using the diagram in the presentation
Atomic
Isotope
Protons
number
Neutrons
Electrons
Mass (a.m.u)
Lithium-6
Lithium-7
Lithium-8
You Try ?
What the number of protons, neutrons and electrons in each?
19
18
F
9
203
F
9
Hg
80
194
Hg
80
How to calculate Average Atomic Mass?
 Average atomic mass is the weighted average of the masses of all naturally occurring isotopes.

Average Atomic Mass Equation:
Average atomic mass = (% abundance of isotope x mass of 1st
isotope) + (% abundance of isotope x mass of 2nd isotope) +
………
Cl-35 and Cl-37
Why is the AAM = 35.45 amu ?
Example 1:
Element x has 2 natural isotopes. Calculate its average atomic mass. 1st isotope has a mass of 10.012 a.m.u.
with 19.91% abundance. 80.09% of the 2nd element has a mass of 11.009 a.m.u.
Example 2:
Calculate the average atomic mass of copper if it has 2 isotopes. 69.11% has a mass of 62.93 a.m.u and the
rest has a mass of 64.93 a.m.u.
Section 3
Obj.3
Nuclear Decay Reactions
Nuclear chemistry is the study of the changes of the NUCLEUS of an atom.
 Nuclear Reactions involve changes within the nucleus where as chemical reactions involve the loss, gain
or sharing of electrons.
The Nucleus
 Contains ________________ and neutrons. They are collectively called NUCLEONS.
Radioactivity
 An unstable nucleus will decay or break down, releasing particles and/or energy in order to become stable.
 An atom with an unstable nucleus is considered “_____________________”.
Transmutation
 Type of nuclear reaction that will change the number of PROTONS and thus will create a new & different
ELEMENT
There are several ways radioactive atoms can decay into different atoms!
Basic Types of Radioactive Decay
Particle
Type
Symbol
Alpha
Beta
***
or
or
What Happens?
Example
Penetrating Power
Atomic number
decreases by ____ and
mass number
decreases by ____
(1) LOW: Can be
blocked by
paper/clothing
Atomic number
increases by _____ but
the mass number stays
the same
(100) MEDIUM: can
penetrate the skin;
need to be
protected by
clothing/thin
metals like
aluminum
***A neutron becomes a ________________ and a high-speed electron that is discharged from the nucleus**
Gamma
(100000) HIGH:
need to be
protected by thick
concrete or metal
like lead
No change in atomic
nor mass number;
occurs with other
types of decay
Writing Balanced Nuclear Equations
The sums of both the ATOMIC NUMBER & MASS NUMBER on the left side must equal the sums of both the
ATOMIC NUMBER & MASS NUMBER.
Often problems will have 1 particle missing and you will need to identify it.
Alpha Decay of thorium-230
You Try!
1. Beta decay of zircomium-97
Beta Decay of cesium -137
3. Complete this:
2. Alpha decay of americium-241
Making New Elements and Isotopes: Bombarding the Nucleus
 All the transuranium elements (elements with atomic numbers higher than Uranium) have been made by
bombarding the nucleus with neutrons and other atoms in accelerators
Example:
Section 4
Obj.3 cont.
Half Life
 Radioactive isotopes decay at a characteristic rate measured in HALF LIVES.
 A half-life time (THL) is the time required for HALF of the amount of radioactive atoms to decay. The time
ranges from __________________ to millions of _____________________.
HOW TO’s
1. To calculate the number of half-lives, divide the half-life (THL) into the total time (T). Then cut the original
amount in half the number of time determined by the # of half lives (HL).
T / THL = # HL
2. Algebraic equation to calculate the remaining amount left over after a certain number of half-lives have
passed.
Amount remaining = (initial amount) (.5) #HL
Examples:
1. Suppose you have 20 grams of sodium-24. Its half-life is 15 hours. How much is left over after 60 hours.
2. Uranium-238 has a half-life of 4.46 x 109 years. How long will it take for 7/8th of the sample to decay?
You Try!
1. 1.5 grams of a 12.0 g sample are left after 114 s. What is the half-life of radium-222?
2. A sample of 3x107 Radon atoms is trapped in a basement that is sealed. The half-life of Radon is 3.83
days. How many radon atoms are left after 31 days?