Download Atomic Structure PowerPoint Presentation

Early Atomic Theory and Structure
Chapter 5—Early Theories
o What is stuff made
of?
o What makes
something move?
o How do we know
it’s alive?
o Is there a
fundamental
particle that
everything is made
up of?
o Is there a universal
constant to all
matter?
Chapter 5.1 Early Thoughts
o Roots of atomic
theory are as old as
440 B.C. with
Democritus’ idea of
the atom
o It took 2 000 years
for us to expand on
this idea. The new
theory was to be
done by an English
schoolmaster John
Dalton in the early
1800s.
Dalton’s Atomic Theory
o
o
o
o
His theory included 6 postulates
1. Elements are made up of atoms
2. Atoms of the same element are alike
3. Atoms of different elements are different by
virtue of their size and mass
o 4. Chemical compounds are formed by the
union of two or more atoms of different
elements
o 5. Atoms combine to form compounds in
whole number ratios (1:2 or 2:2, etc]
o 6. Atoms of two elements may combine in
different ratios to form more than one
compound
5.2 Dalton’s Atomic Theory
(cont.)
o Why is it a Theory?
o Which are still true?
o Which do we know
More info about now?
5.3 Composition of
Compounds
o The Law of Definite Composition
states that a compound always
contains two or more elements
combined in definite proportion
by mass
Law of Multiple Proportions
o The Law of Multiple Proportions
states that atoms of two or more
elements may combine in
different ratios to produce more
than one compound
5.4 - 5.8 Subatomic Particles
o Through the years of the late 1800s and
into the early 1900s it was determined
that there are three subatomic particles
o Electrons (discovered first)
o Protons (reasoned to exist if elements are
neutral)
o Neutrons (discovered last)
Electrons
o 1. Electron which occupies
the area outside the nucleus
and has a negative charge,
relative to the other
subatomic particles it has
negligible (so small that it can
be ignored) mass.
Protons
o 2. Proton which exists in the
nucleus, has a positive charge
and has mass roughly equal
to neutrons
Neutron
o 3. Neutrons (discovered last]
functions as the glue that
holds the nucleus together so
that the protons don’t repel
each other, it has no charge
and roughly the same mass as
the proton
Isotopes
oIsotopes have same
number of protons (so
they are the same
element) but different
number of neutrons
oSome isotopes are
radioactive
Atomic Number
oAtomic Number = the
number of protons; unique to
each element and the way
the periodic table is
arranged
Mass number
oMass Number = protons +
neutrons (whole number
oCannot be found on the
periodic table!
Check yourself
o The nucleus is made up of what
two types of subatomic
particles?
Formulas you should know
o Atomic number = # of protons
o In an atom (uncharged):
o # of protons = # of electrons
o Mass # = # protons + # neutrons or
o # neutrons = mass # - # protons
o Charge = # protons - # electrons (for
ions)
o Remember the atomic # and # of
protons give the element its identity
and does not change
Elements composed of atoms
Elements or atoms in
an unbonded state
have the same
number of electrons
as protons
(They are neutral)
Ions
Ions have an unequal
number of electrons and
protons. An atom loses or
gains electrons to take on
a charge (protons/neutrons
are not transferred)
Charge = #protons - #
electrons
Ionic Charge
o Charge is written in the upper
RIGHT corner of the element’s
symbol.
o It is written with the number first
and the sign second unless it is a
+ 1 or a -1 in which case it is just
written as + or -.
o Negative ions change their
names to end in –ide like fluorine
is fluoride
Ionic Notation
3X
This means that this
element has a -3
charge.
Self Check
o What is the charge of a
substance with 14 protons, 15
neutrons, and 14 electrons?
Self Checker
o If a substance has a charge of
+2, this means that the number of
protons is (circle one: LESS than
or GREATER than) the number of
electrons?
ISOTOPIC NOTATION
isotopes are atoms with the same number of protons
but different number of neutrons
A
Z
X
A = mass number
(the total number of protons +
neutrons)
Z = atomic number
(the total number of protons)
X = element symbol
READING ISOTOPIC NOTATION
46
21
Sc
46 = mass number
(the total number of protons (21) +
neutrons (25)
21 = atomic number
(the total number of protons (21))
Sc = element symbol
In a neutral atom, the number of electrons
(21) is equal to the number of protons.
PRACTICE PROBLEMS
15N
7
# protons = ____
___
35P
15
# p = ____
62Cu2+
29
# p = ____
76Se3-
34
# p = ____
8
# neutrons= ____
#electrons = 7
# n=20
____
15
#e- = ___
33
# n= ____
27
#e- = ___
42
# n= ____
#e- =37
___
Writing ISOTOPIC NOTATION
1. Write the symbol for the atom with an atomic
number of 21 and a mass number of 48.
48
Sc
2. Give the complete chemical notation for the
nuclide with 23 protons, 26 neutrons and 20
electrons.
49V3+
3. Write the isotopic notation for
110Pd
a. Z = 46
A = 110
b. An atom containing 24 protons, 28 neutrons,
and 21 electrons
52Cr3+
c. Titanium-50 50Ti
PRACTICE PROBLEMS
1. 196 Pt4+
118
78 # n = _____
# p = _____
196
mass number = ________
195.1 amu
atomic mass = ________
74
#e- = _____
78
atomic number = _______
platinum
name of element = _______
2. Indicate the appropriate atomic mass of an element with
30 protons, 30 neutrons, and 28 electrons.
65.39 amu
Atomic Mass
oAtomic Mass = number on
the periodic table reflecting
the mass all isotopes known
and their relative
percentages (on periodic
table below element’s
symbol--usually not a whole
number)
Atomic Mass
o The atomic mass of an element represents
the average mass of all the isotopes found in
nature. No element exists with only one
possible isotope. Hydrogen has the smallest
number of isotopes: 1H protium, 2H
deuterium, 3H tritium. Its atomic mass is
1.0079 amu (atomic mass units). The atomic
mass is calculated by adding the % of 1H
mass found in nature to the % of 2H mass
found in nature plus the % of 3H mass.
o % 1H + % 2H + % 3H = average mass (atomic
mass)
o Generally the formula used is:
% X + % Y + % Z… = atomic mass.
An instrument called the mass spectrometer is
generally used to determine the percentages
and individual masses of each isotope.
Atomic Mass
o Silver is found to have two stable isotopes, one has an
atomic mass of 106.904 amu and the other weighs
108.905 amu. The first isotope represents 51.82 % of the
mass of the element and the second represents 48.18
%. What is the atomic mass of the element silver?
The equation to use is %X + % Y = average
And remember to turn your percents into fractions before
multiplying.
(0.5182) 106.904 amu + (0.4818) 108.905 amu =?
55.398 amu + 52.470 amu =?
107.868 amu !!
Now look at the periodic table to verify the answer.
PRACTICE PROBLEMS # 8
1. A sample of neon contains three isotopes, neon-20 (with an
isotopic mass of 19.9924 amu), neon-21 (20.9939 amu) and
neon-22 (21.9914 amu). The natural abundances of these
isotopes are 90.92%, 0.257 %, and 8.82 %. Calculate the
atomic weight of neon.
20.17 amu
2. There are only two naturally occuring isotopes of copper,
63Cu and 65Cu. Copper has an atomic mass of 63.55 amu.
What is the natural abundance of each isotope?
65Cu
= 30% & 63Cu = 70%
3. There are only two naturally occuring isotopes of gallium,
69Ga and 71Ga. What is the natural abundance of each
isotope?
69Ga
= 60% and 71Ga = 40%
GROUP STUDY PROBLEM #8
_______1. The element with atomic number 53 contains
a) 53 neutrons b) 53 protons C) 26 neutrons & 27 protons d) 26 protons & 27
neutrons
_______2. The mass of one atom of an isotope is 9.746 x 10-23 g. One atomic mass unit has
the mass of 1.6606 x 10-24 g. The atomic mass of this isotope is
a) 5.870 amu
b) 16.18 amu
c) 58.69 amu
d) 1.627 amu
108
_______3. The number of neutrons in an atom of
a) 47
b) 108
c) 155
47
Ag is
d) 61
27
Al3+
_______4. The number of electrons in an ion of 13
is
a) 13
b) 10
c) 27
d) 14
_______5. What is the relative atomic mass of boron if two stable isotopes of boron have the
following mass and abundance:
10.0129 amu (19.91%) & 11.0129 (80.09%)
a) 10.81 amu
b) 10.21 amu
c) 10.62 amu
d) 10.51 amu
Test your Knowledge
Name
Symbol
Atomic
#
# of
protons
Neon
# of
neutrons
Mass #
11
Pb
207
74
110
88
226
8
Carbon
F
19
79
118
# of
electrons
Test your Knowledge
Name
Neo
n
Lead
# of
Symbol Atomic
protons
#
# of
neutron
s
Mass #
# of
electro
ns
Ne
10
10
11
21
10
Pb
W
82
74
82
74
125
110
207
184
82
74
radiu
m
Carb
on
Ra
88
88
138
226
88
C
6
6
8
14
6
fluorin
e
F
9
9
10
19
9
Gold
Au
79
79
118
197
79
Tungste
n
Table Information
Hydrogen
H
Atomic Number
1
Atomic Weight
1.00794
Oxidation States
+1, -1
Electronegativity, Pauling 2.2
State at RT
Gas, Non-metal
Melting Point, K
14.01
Boiling Point, K
20.28
Symbol
The Periodic Table
o Horizontal rows are called
periods
o Vertical columns are called
groups
o We will use 1- 18 as group
designations.
o Group 1 is Alkali Metals
o Group 2 is the Alkaline Earth
Metals
o Group 18 Inert or Noble Gases
o Group 17 Halogens
Larger Groups
o Groups 3 –12 are the heavy
metals or transition elements
o Two periods at the bottom are
called the rare earth elements or
the inner transition elements.