Atoms and Atomic Theory
 Atom - the smallest particle of a given type
of matter.
 Atomic Theory – the idea that matter is
made up of fundamental particles called
atoms.
What discoveries led to the
development of atomic structure?
 Democritus – Father of the Atom
 400 B.C.
 atomos – “indivisible”
John Dalton – Father of the Modern
Atom
 Dalton’s Atomic Theory.
1.) All matter is made of atoms.
2.) All atoms are indestructible and cannot
be divided into smaller particles.
3.) All atoms of one element are exactly
alike, but they are different from all
other elements.
Subatomic Particles
 Electrons
 Protons
 Neutrons
 Nucleus
Electrons
 Symbol: e Mass = 1/1837 amu
 1 amu ~ the mass of a hydrogen atom
 Charge = (-1)
 J.J. Thomson – Discovered the electron
using a cathode ray tube
Indirect Evidence
 Evidence providing only a basis for inference
about the fact in dispute.
 Indirect evidence provided information
about electrons.
 Diagram this experiment.
Protons
 Symbol: +p
 Mass = 1 amu
 Charge = (+1)
Neutrons
 Symbol: n0
 Mass = 1 amu and there is no charge
Thomson’s Model
 Thomson put both
the electron and
proton together and
proposed the “Plum
Pudding” model of
the atom.
Nucleus

Ernest Rutherford
 Using the “Gold Foil” experiment,
Rutherford determines:
1. An atom is mostly empty space.
2. There is a nucleus in the middle of the
atom containing the protons.
3. The electrons orbit a large distance away
from the nucleus.
 Proposes the nuclear model of the atom.
Let’s Review
 http://www.brainpop.com/science/matterandchemistr
y/atomicmodel/
How Big is an Atom?
 Scale of the Universe website:
http://www.youtube.com/watch?v=fUAFqkS7y9M&fe
ature=related
 Current Atomic Model:
http://www.youtube.com/watch?v=xqNSQ3OQMGI
Unit 1 • Investigation III
Atomic Number
 Designates the number of protons in a
nucleus of an atom.
 Each element has a characteristic atomic
number.
 The number of electrons equals the atomic
number in a neutral atom.
Mass Number
 Designates the total number of protons and
neutrons in an atom.
 Number of neutrons = mass number atomic number.
 Atoms of the same element can have
different mass numbers.
Objective 2
 Explain why isotopes differ.
Isotopes
 Atoms with the same number of protons, but
different numbers of neutrons.
 Isotopes of an element have the same atomic
number, but different mass numbers.
 Nuclear Symbol or isotopic
symbol – shows number of
protons, neutrons and electrons in
an atom.
 Isotopes may also be described by labeling
the name following by the mass number
(Example: carbon-12).
Atomic Mass
 Weighted average mass of the atoms in
naturally occurring sample of an element.
 Masses are based off of the atomic mass
unit (amu) defined as one twelfth the mass
of a carbon-12 atom.
Atomic Mass Example
 In nature carbon is composed of 98.89% 12C
atoms and 1.11% 13C atoms. 12C has a mass of
12 amu and 13C has a mass of 13.0034 amu.
What is the average atomic mass of carbon?
Review
 John Dalton – Dalton’s Atomic Theory
stating atoms are solid, indivisible spheres
 J. J. Thomson – Thomson Model or “Plum
Pudding Model” stating electrons are
dispersed through a sphere of positive
charge
 Ernest Rutherford – Rutherford Model
stating atoms have a small, dense, positively
charged nucleus surrounded by electrons
Objective 3
 Diagram the Bohr Model of the atom.
Bohr Model
 Niels Bohr (1913) – Bohr Model stating electrons
travel around nucleus in energy levels.
 Uses rings to show the energy levels.
 The number of rings should match the row the
element is in on the periodic table.
Bohr Model Diagram
 The protons in the nucleus are found by looking at
the atomic number.
 The neutrons in the nucleus are found by
subtracting the atomic number from the mass
number.
 Electrons fill the shells from low to high until the
correct number of electrons are added (number of
electrons = number of protons).
 This model works for the first 20 elements.
Diagram Lithium-7
Diagram Nitrogen-14
Diagram Aluminum-27
Diagram Calcium-40
Objective 4
 Identify the position of groups, periods, and
different chemical families on the periodic
table.
Periodic Table
 Dmitri Mendeleev – mid 1800’s
 Proposed a table for 70 elements based on
mass and properties
 Henry Moseley – 1913
 Determined the atomic number of
elements and arranged the table in order
of atomic number
How is the periodic table of
elements arranged?
 Periods: Rows of the periodic table
 Groups: Columns of the periodic table
 Periodic Law: When elements are arranged
in order of increasing atomic number, there is
a periodic repetition of their chemical and
physical properties.
Groupings to know on the
Periodic Table
 Representative Elements
 Metals
 Non-metals
 Metalloids
 Transition Metals
 Inner-transition Metals
 Alkali Metals
 Alkaline Earth Metals
 Halogens
 Noble (inert) gases
Properties of Metals
 Metals:
 Bright metallic
luster
 Solids are easily
deformed
 Good conductors
of electricity and
heat
Properties of Nonmetals
 Nonmetals:
 Non-lustrous,
various colors
 Solids may be
hard or soft,
usually brittle
 Poor conductors
of electricity
Properties of Metalloids
 Metalloids have
properties
intermediate
between metals and
nonmetals.
Objective 5
 Identify forces between atoms.
Lewis Dot Structures
 Show valence electrons.
 Look at the group number. This is how many dots you
draw for each atom.
 Ex:
Potassium
Fluorine
Sulfur
Phosporous
Aluminum Beryllium
Silicon
Xenon
Bonding & Reactions
 The main goal of chemical boding and
reacting is to fill the outermost energy level.
 This is called the octet rule. (8 electrons in
the outermost shell, EXCEPT H, He, Li, Be,
which need 2.)
Two Types of Bonds
1.) Ionic
2.) Covalent
Ionic Bonds
 When a metal ion with a positive charge
sticks to a nonmetal ion with a negative
charge.
 Also called salts.
Ions: an atom that has extra or
missing electrons
 Metals lose their outer electrons, becoming positive.
Ex: Na
Be
Al
 Nonmetals gain electrons in their outer shell, becoming
negative.
Ex: N
O
Cl
Where do these gained or lost
electrons go?
 Ex: Sodium chloride
 Metal = sodium
 Ex: Aluminum oxide
 Metal = aluminum
nonmetal = chlorine
nonmetal = oxygen
Use the valence electrons!
 To find the charge of an ion, look at the valence
electrons and count how many they need to give away
or add to reach a full energy level (8 electrons for
most).
 Ex: What charge will the following ions have?
 sodium
magnesium
 phosphorous
carbon
 Bromine
sulfur
Forming the compound
 The compound is then balanced by crossing charges,
reducing if possible.
Metal
Charge
Nonmetal
Balanced Compound Formula
Charge
Li
+
C
=
Mg
+
F
=
K
+
O
=
Al
+
P
=
Brain Pop: Ions
Covalent Bonds
 Non-metals share electrons to fill outer shells. (1 bond
= 2 electrons)
 Remember, nonmetals like to gain electrons. Since
they both want to gain, they must share.
 Follow these rules:
1.) Look at the formula. Then add up total valence
electrons needed for your drawing.
2.) Single bond all of the atoms, picking a center atom
when possible.
3.) Fill every atoms energy level. Remember, most need 8
electrons, but H only needs 2.
4.) Count the electrons & erase some if you have too many.
Move the electrons into double or triple bonds if
needed.
Practice
Cl2
H2O
H2O2
CCl4
CH4
CO2
Ionic or Covalent?
(Metal + Nonmetal = Ionic
Compound
NaCl
H2O2
C6H12O6
MgF2
Al2O3
CO2
O2
Li3P
Nonmetal + Nonmetal = Covalent)
Metal + Nonmetal or
Nonmetal + Nonmetal?
Ionic or Covalent?
Lewis Dot Structures Practice
N2
F2 O
Cl2O2
CI4
CBr4
CO2