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Transcript
Standards
 1a. Students know how to relate the position of an
element in the periodic table to its atomic number
and atomic mass.
 1b. Students know how to use the periodic table to
identify metals, semi-metals {metalloids}, nonmetals, and halogens
 1e. Students know the nucleus of the atom is much
smaller than the atom yet contains most of its
mass.
Purpose
 We will use this information to build our chemistry
knowledge.
 We will use this information as the foundation to
calculate limiting reagent problems.
 The standardized exams in the spring will test you on
this information.
Objectives
 Know the 3 particles of the atom and where they
reside
 Know the difference between atomic number and
mass number
 Know how to write nuclide symbols
 Know the three isotopes of hydrogen
 Know how to calculate atomic mass
 Know how to calculate percent composition
First Some Questions…
 What are atoms made up of?
 Protons, Neutrons and Electrons
 Where do you find these particles?
 Protons and Neutrons are located in the nucleus
 Electrons are located in the outer rings, outside the
nucleus.
Vocabulary
 Atom- from the Greek atomos=indivisible. The atom is the
smallest particle of an element that retains the properties
of that element.
 Nucleus: the center of the atom; composed of neutrons
and protons. Because the mass of the proton and the
neutron is much larger than that of electrons, almost all
the mass is located in the nucleus.
 Ion: a charged particle; # protons ≠ # electrons
 Electrons occupy most of the volume of an atom
outside/around the nucleus.
Fundamental Particles
 Proton
 A positively charged particle located in the nucleus.
 Neutron
 A neutral particle located in the nucleus.
 Electron
 A negatively charged particle located outside the
nucleus.
Question
 What differentiates one atom from another atom?
 The number of PROTONS
Atomic Number (Z)
A
Z
El
 Number of protons in the nucleus of an atom
 This number is found on the Periodic Table
 Atomic Number identifies an element
 Always a positive number (b/c it is a counting
#)
 Tells number of electrons in a neutral atom
 An atom is electrically neutral
What does it mean to be
electrically neutral?
 The atom has no charge
 The number of protons = the number of electrons
Question
 What observations can you make about atomic
numbers on the periodic table?
 Atomic Number increases as you go across the rows from
left to right.
Questions
 What is the atomic number of Chlorine?
 What can you tell me about its protons and electrons?
 What element has 20 protons?
 What is the relationship between the # protons and
the atomic number?
 They’re equal.
Complete the Chart
Element
Symbol
Atomic #
# Protons
Potassium
K
19
19
Boron
B
5
5
Sulfur
S
16
16
Yttrium
Y
39
39
Mass Number (A)
A
Z
El
 Total number of protons and neutrons in the
nucleus of an atom
 Always a positive number
 You can determine the nuclear composition
of an atom from its mass number and atomic
number
Question
 What do the atomic number and the mass number
have in common?
 Both Positive integers
 Both have the same # of protons
How to find # of Neutrons
 Mass # - Atomic#= # Neutrons
 Or
 # protons + # neutrons= Mass #
 (atomic number + # neutrons)=Mass #
Complete the Chart
Atomic#
Mass#
#Protons
#Neutrons #Electrons Chemical
Symbol
9
19
9
10
9
F
14
29
14
15
14
Si
22
47
22
25
22
Ti
25
55
25
30
25
Mn
6
12
6
6
6
C
Isotopes
 Atoms of the same element with differing numbers of
neutrons
 Atoms with the same atomic number but different
mass number
 Isotopes of an element have different masses
 Chemical properties of different isotopes are virtually
the same
Nuclide Symbol
A
Z
 A=Mass #
 Z= Atomic #
El
Nuclide
 A specific kind of atom
 Specification of an element in terms of its
nuclear composition/structure
 Tells number of protons and number of
neutrons
Chemical
Symbol
C
Nuclide Symbol
13
6
C
# protons
# neutrons
# electrons
6
[6, 7, 8]
6
6
7
6
Complete the Chart
Atomic#
Mass#
#Protons
#Neutrons #Electrons Chemical
Symbol
Nuclide
Symbol
9
19
9
10
9
F
19
9
14
29
14
15
14
Si
29
14
22
47
22
25
22
Ti
47
22
25
55
25
30
25
Mn
55
25
6
12
6
6
6
C
12
6
F
Si
Ti
Mn
C
3 Isotopes of Hydrogen
Isotope
Of
Hydrogen
Protium
Deuterium
Tritium
Nuclide
Symbol
# protons
# neutrons # electrons
H
1
0
1
2
1
H
1
1
1
3
1
H
1
2
1
1
1
Nuclides
 By specifying the nuclear structure, then you call it a
nuclide.
 But if you say Carbon atom, you do not know which
Carbon atom it is, therefore you don’t know how many
neutrons it has
 Example: Brothers and Sisters You are members of the Jones family, but you have not
specified which Jones member you are referring to.
Write the nuclide name and nuclide symbol
# protons # neutrons # electrons Nuclide
Name
17
20
18
Chlorine-37
Anion
20
20
18
Calcium-40
Cation
92
146
92
Uranium-238
Nuclide
Symbol
37
17
Cl
40
20
Ca
238
92
U

2
Atomic Mass
 A weighted average of the atoms in a naturally
occurring sample of the element.
 Naturally occurring: no matter where you get the sample
from, it will have the same percentages of isotopes.
Construct a Fruit Basket
 Fruit Type
Weight of Each Piece
2 grapefruit
14 oz
4 apples
10 oz
3 pears
7 oz
1 kiwi
3 oz
 What is the Average Weight?
Fruit Basket
 Average weight=9.2oz
 Each type of fruit makes a different contribution to the
overall weight
 How many pieces of fruit actually weigh 9.2 ounces?
 None!
 What does 9.2 oz mean?
 Fictitious non-existent piece of fruit
Atomic Mass
 If you have a recipe, you could count items to put
in, say 200 chocolate chips, 3 eggs, etc.
 But suppose I have a recipe to make a compound.
 I need 100 hydrogen atoms and 50 oxygen atoms-you
cannot count atoms or pluck them out with atomic
tweezers!
 So instead they mass them (weigh them)
 Careful here, the mass of an object is completley
different from the weight of an object.
Question
 What accounts for the mass of the atom?
 # protons & # neutrons in the nucleus
Atomic Mass
 Know that 1.0 amu is defined as exactly 1/12
the mass of a 126C atom.
 Carbon-12 has 6 protons and 6 neutrons,
therefore 1 proton or 1 neutron = ~1 amu
 1 amu = 1.6606 x 10 -24 grams
 Since the mass mostly depends on # protons
and # neutrons, you’d think atomic mass
would be a whole number, but it isn’t. How
come?
Atomic Mass
 In nature, most elements exist as a mixture of 2 or more isotopes.
 Each isotope of an element has a fixed, constant mass and fixed constant
relative abundance.
 Relative abundance(amount)
X 100=
(how much of each isotope is present)
%
 Sample of carbon from anywhere in the world; coal from S. Africa, W.
Virginia or Pennsylvania
→ 99% C-12 and 1% C-13
 Atomic Mass of periodic table takes into account the larger and smaller
masses of the isotopes, just like the average piece of fruit
accommodated the larger and smaller masses.
 → Idea of weighted average
Calculating Atomic Mass
 To calculate atomic mass you need to know 3 things:
 # of stable isotopes
 Mass of each isotope
 % abundance of each isotope
 Each isotope is a piece of fruit and the isotope’s mass is
the weight of each piece of fruit.
Example: Chlorine Calculation
 mass of isotope X relative abundance
+ mass of isotope X relative abundance
=_______amu
Isotope
Mass of Isotope
Relative Abundance
Cl-35
34.969
75.77%
Cl-37
36.935
24.23%
Atomic Mass
 (34.969)(.7553) + (36.935)(.2447) =
35.4500amu
 That’s the same value on the periodic
table!
Question
 How many chlorine atoms actually have a mass of
35.45 amu?
 NONE
 So the atomic mass, in amu, is the average of a
fictitious non-existent atom of an element.
Example: Copper Calculation
Isotope
Mass of Isotope
Relative Abundance
Cu-63
62.9298amu
69.09%
Cu-65
64.9278
30.91%
Atomic Mass
(62.9298)(.6909)+(64.9278)(.3091)= 63.5464 amu
Calculating Relative Abundance
 To Calculate % Abundance:
 Make a Chart
 Isotopic Mass X %Abundance of each isotope
 Set-up equation
 Solve for “x”
 Plug in “x” value to solve for “y”
Example
Isotope
Mass of Isotope
B-10
10.013
B-11
11.009
10.103 (x) + 11.009 (1 –x) = 10.811
10.103x + 11.009 -11.009x = 10.811
-0.996x = -0.198
x = .1987
y= 1-.1987
y= .8013
B-10 = 19.87%
B-11 = 80.13%
Relative Abundance
x
1- x
1.00
x + y = 1.00
y=1–x
Atomic Mass
The End