Chapter 5 Atomic Structure
Objective (1) Summarize Dalton’s atomic theory
Objective (2) Describe the size of an atom
 Early Models of the Atom
1. Democritus
First suggested the existence of small
particles, called atoms
2. John Dalton
a. Performed experiments to test his
atomic theory
b. Studied the ratios in which elements
combine in chemical reactions
c. Formulated theories from the
findings
d. Atomic Theory
1. All matter is composed of tiny
indivisible particles called atoms.
2. All atoms of a given element
are identical. Atoms of a specific
element are different from any other
element.
3. Elements can be chemically combined to
form compounds.
• Just how small is an atom?
1. A Cu penny composed of pure copper
2. Grind the penny into fine powder, each
speck of powder is made smaller
3. The smallest piece left is an atom
4. The Cu penny contains 2.4 x1022 atoms
5. However, if you lined up 100,000,000
copper atom side by side it would only
be 1 cm long
6. Atoms are very small
• Objectives:
3. Distinguish among protons, electrons, and
neutrons in terms of relative mass and charge
4. Describe the structure of an atom, including
the location of the protons, electrons, and
neutrons with respect to the nucleus
 Electrons
1. negatively charged subatomic particles
2. J.J. Thomson
- Discovered the electrons
- His experiment: passed an electrical
current through gases at low pressure
• Cathode ray tubes
Results:
1. Regardless of the types of electrode,
the same particles appeared
2. The particles were repelled by the
negative plate of a magnet.
3. The particles caused a paddle-wheel to
spin
Conclusions:
1. The particles were in all kinds of matter.
2. The particles were negatively charged.
3. The particles had mass.
4. He proposed the plum pudding model of
matter.
• Protons and neutrons
1. Chadwick: discovered the neutron
neutron: no charge subatomic particle
2. Rutherford: discovered the nucleus, and
that it was positive
: performed the Gold-Foil
Experiment
What he did:
1. Shot alpha particles (+ charge) at a sheet
of gold
2. Picture:
Results:
1. Most of the alpha particles passed
straight through the gold
2. A few of the alpha particles were slightly
deflected
3. A very small number of alpha particles
hit and bounced back off the foil
 Conclusion
1. He proposed the nuclear model of the
atom.
2. Matter is mostly empty space (with light
electrons)
3. Scattered throughout are small areas of
positive
4. Very dense matter called the nucleus.
Particle
Symbol
Electron
Proton
Neutron
ep+
N0
Relative
Electrical
charge
11+
0
mass of atoms comes from :
volume of atoms comes from:
Actual mass
(g)
9.11 x 10-28
1.67 x 10-24
1.67 x 10-24
• Objectives:
5. Explain how the atomic number identifies
an element.
6. Use the atomic number and mass number of
an element to find the number of protons,
electrons, and neutrons.
Atomic number
1. the number of protons in the nucleus of an
atom
2. In a neutral atom it also equals the
electrons
3. Ex. K atomic # = 19
p+ = 19
e_ = 19
Mass number
1. the total number of protons and neutrons in
the nucleus
2. mass number = p+ + n0
3. neutrons = mass number – p+
• How to write an atom:
1.
2.
A
Z
X
Where A = mass #
Z = atomic #
X = any element
3. Hydrogen – 1 ( mass number)
4. Fill in the table:
Element Atomic Mass
Protons Electron Neutron
number number
15
16
89
K
30
39
19
• What is a neutral atom?
- It has the same number of p+ as e−
-
Ex. Ca atomic # = 20
p+ = 20
e- = 20
so you have +20
so you have -20
What is the total charge? 0
Objectives:
7. Explain how isotopes differ and why the
atomic masses of elements are not whole
numbers.
8. Calculate the average atomic mass of an
element from isotope data.
Isotopes
1. atoms with the same number of
protons but different number of
neutrons
2. Identified by the mass
3. Example: Chlorine-35
Chlorine-37
mass number
Atomic mass unit– weighted average mass
of the isotopes of that element
1. Ex. Cl 35.453 amu
75% is chlorine-35
25% is chlorine-37
2. (mass)(% of element) + (mass)(% of
element) = (total mass)
* change % to a decimal
 Examples:
3. Ex. Element X has 2 natural isotopes. The
isotope with a mass of 10.012 amu (X-10) has
a relative abundance of 19.91%. The isotope
with a mass of 11.009 amu (X-11) has a
relative abundance of 80.09%. Calculate the
atomic mass of this element.
4. An element, E, has an atomic mass of 18.40 and
consists of two isotopes:
E-17 with a mass of 16.95 and E-20 with a mass of
19.35. How much E-20 does this element contain?
E-20 = x
E-17 = (100-x)
(19.35)(x) + (16.95)(100-x) = (100)(18.40)
 Ion
1. atoms lose or gains one or more electron
2. Mg+2
number of protons = 12
number of electrons = 10
charge of ion = +2
3. + ions : loses electrons
4. - ions : gains electrons
3. Examples
Atomic # Mass #
27
Al3+
13
70
Zn
30
p
n
e
• Periodic Table
1. Mendeleev
- created the first periodic table
- used atomic mass to place elements
- left blanks for undiscovered elements
2. Henry Moseley
- determined the atomic number
• The modern periodic table
• Characteristics
1. Rows or periods : there are seven
2. Groups, families or columns: there are 18
3. Periodic law: there is a repeating pattern
of physical and chemical properties
4. Metals
- have high electrical conductivity
- high luster (shiny)
-ductile
- malleable
- all solids at room temperature except:
mercury (Hg)
5. Nonmetals
- nonlustrous
- poor conductors
- brittle
- can be solids, liquids, or gases
6. Metalloids
- have both metal and nonmetal properties
- touch the zigzag line