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Definition of Atoms
• greek word
• Atoms are the
atomos which
fundamental building
means
blocks of all matter, not
“indivisible”. This
able to be split by ordinary
is based upon
chemical reactions
the discontinuous
theory of matter.
(meaning matter
can not
continually be
split and still
remain the same)
HISTORY OF THE ATOM
460 BC Democritus develops the idea of atoms
he pounded up materials in his pestle and
mortar until he had reduced them to smaller
and smaller particles which he called
ATOMA
(greek for indivisible)
HISTORY OF THE ATOM
350 BC
• Aristotle modified an earlier theory
that matter was made of four
“elements”: earth, fire, water, air.
• Aristotle was wrong. However, his
theory persisted for 2000 years.
Aristotle
fire
earth air
water
HISTORY OF THE ATOM
suggested that all matter was made up of
1808 John Dalton tiny spheres that were able to bounce around
with perfect elasticity and called them
ATOMIC THEORY
• All matter is made of atoms.
• Atoms of an element are identical.
• Each element has different atoms.
• Atoms of different elements combine in
constant ratios to form compounds.
• Atoms are rearranged in reactions
HISTORY OF THE ATOM
1898 Joseph John Thompson
found that atoms could sometimes eject a far
smaller negative particle which he called an
ELECTRON
HISTORY OF THE ATOM
1904
Thompson develops the idea that an atom was made up of
electrons scattered unevenly within an elastic sphere surrounded
by a soup of positive charge to balance the electron's charge
like plums surrounded by pudding.
PLUM PUDDING
MODEL
HISTORY OF THE ATOM
1910 Ernest Rutherford
oversaw Geiger and Marsden carrying out his
famous experiment.
they fired Helium nuclei at a piece of gold foil
which was only a few atoms thick.
they found that although most of them
passed through. About 1 in 10,000 hit
HISTORY OF THE ATOM
gold foil
helium nuclei
helium nuclei
They found that while most of the helium nuclei passed
through the foil, a small number were deflected and, to their
surprise, some helium nuclei bounced straight back.
HISTORY OF THE ATOM
Rutherford’s new evidence allowed him to
propose a more detailed model with a central
nucleus.
He suggested that the positive charge was all
in a central nucleus. With this holding the
electrons in place by electrical attraction
However, this was not the end of
the story.
HISTORY OF THE ATOM
1913 Niels Bohr
studied under Rutherford at the Victoria
University in Manchester.
Bohr refined Rutherford's idea by adding
that the electrons were in orbits. Rather
like planets orbiting the sun. With each
orbit only able to contain a set number of
electrons.
Bohr’s Atom
electrons in orbits
nucleus
Bohr’s Planetary Model of Atom
another look
HELIUM ATOM
Shell
proton
+
N
-
+
- N
electron neutron
What do these particles consist of?
ATOMIC STRUCTURE
Particle Charge Mass
proton + 1 charge 1amu
neutron 0 1amu
electron -1 charge 1/1836
• Amu = atomic mass unit
• 1 amu = 1.66 x 10-24g
Relative Sizes
• Thus 99.99% of the mass of an atom
comes from the nucleus (protons and
neutrons) and essentially nothing from e-
• the major volume of atoms coming from
the electron cloud
• Remember that most of the atom is empty
space like spinning blades of a fan take up
more space than any blade would if not
moving
ATOMIC STRUCTURE
4
He
Atomic mass
the number of protons and
neutrons in an atom
2 Atomic number
the number of protons in
an atom
number of electrons = number of protons
Calculating subatomic particles
• p+ = atomic number
• n0 = mass # - p+
because mass of atom is p+ + n0
get the mass # by rounding average atomic mass on
periodic table to nearest whole number if needed
• e- = p+ - charge
because charge of atom/ion is p+ - e-
Protons, Electrons, Neutrons
Determine the number of protons, electrons and
neutrons for each of the following elements;
40 23 16
a) Ca b) Na c) O
11 8
20
P=20 e=20 p=11 e=11 p=8 e=8
n=20 n=12 n=8
35 28 11
d) Cl e) Si f)
5
B
17 14
p=17 e=17 p=14 e=14 p=5 e=5
n=18 n=14 n=6
Equations:
Protons = atomic #,
neutrons = mass # - protons,
electrons = protons with a neutral atom (charge = 0)
electrons = protons – charge with a ion (charged atom)
Symbol protons Mass # neutrons electrons
7Li atom 3 p+ 7 amu 7-3 = 4 n0 3 e-
the atomic # (same as p+)
Cs atom
Rb atom
32Cl-1 ion
Al-3 ion
42Ca+2 ion
Equations: mass # = atomic mass rounded to nearest whole
Protons = atomic #
Neutrons = mass # - protons
with a neutral atom (charge = 0) the electrons = protons,
with a ion (charged atom) the electrons = protons – net charge
Symbol protons Mass # neutrons electrons
7Li atom 3 p+ 7 amu 7-3 = 4 n0 3 e-
the atomic # (same as p+)
Cs atom 55 133 78 55
Rb atom 37 85 48 37
32Cl-1 ion 17 32 15 18
Al-3 ion 13 27 14 16
42Ca+2 20 42 22 18
ion
Isotopes
• Atoms with same atomic number (number of
protons), but with different masses (due to different
number of neutrons)
Protons Neutrons Mass #
Carbon-12 6 6 12
Carbon-13 6 7 13
Carbon-14 6 8 14
12-6 = 6 13-6 = 7 14-6 = 8
Lithium Isotopes
• Li-6 is the chemical
symbol for Lithium with a
mass number of 6 amu.
(this is 3 protons plus 3
neutrons in nucleus.
• Li-7 has 3 protons and 4
neutrons in nucleus and
a mass of 7 amu.
• The atomic mass (or
average weight) of
Lithium is 6.941.
Therefore which isotope
is most abundant in
nature?
Isotopes
Protons Neutrons Mass #
Lithium-8 3 5 8
Lithium-9 3 6 9
Lithium-11 3 8 11
8-3=5 9-3=6 11 – 3 = 8
IONS
• IONS are atoms or groups of atoms with a positive or negative charge.
• Taking away an electron from an atom gives a CATION with
a positive charge
• Adding an electron to an atom gives an ANION with a negative
charge.
• To tell the difference between an atom and an ion, look to see if there is a
charge in the superscript! Examples: Na+ Ca+2 I- O-2
Na Ca I O
Forming Cations & Anions
• A CATION forms when an atom loses
one or more electrons.
Mg --> Mg2+ + 2 e-
An ANION forms when an atom gains one
or more electrons
F + e- --> F-
PREDICTING ION CHARGES
In general
• metals (Mg) lose electrons --->cations
• Nonmetals (F) gain electrons anions
Ion Practice
State the number of protons, neutrons, and
electrons in each of these ions.
39 K+ 16O -2 41Ca +2
19 8 20
#p+ 19 8 20
#no 20 8 21
#e- 18 10 18
Charges on Common Ions
-3 -2 -1
+1
+2
By losing or gaining e-, atom has same
number of e-’s as nearest Group 8A atom.
Special Family Names on the Periodic Table
Alkali Metals = hot orange Semimetals/Metalloids = pink
Alkaline Earth Metals = Faded Blue Halogens = Yellow
Transition Metals = Indigo Noble Gases = Pumpkin Orange
Other Metals = Baby Blue Other Nonmetals = Green
Bohr - Rutherford diagrams
• Putting all this together, we get B-R diagrams
• To draw them you must know the # of protons,
neutrons, and electrons (2,8,8,2 filling order)
• Draw protons (p+), (n0) in circle (i.e. “nucleus”)
• Draw electrons around in shells
He Li
Li shorthand
2 p+ 3 p+ 3 p+ 2e– 1e–
2 n0 4 n0 4 n0
Be B
Al
4 p+ 5 p+ 13 p+
5 n° 6 n° 14 n°
O Na
8 p+ 2e– 6e– 11 p+ 2e– 8e– 1e–
8 n° 12 n°
ATOMIC STRUCTURE
Electrons are arranged in Energy Levels or
Shells around the nucleus of an atom.
• first shell a maximum of 2 electrons
• second shell a maximum of 8 electrons
• third shell a maximum of 8 electrons
ATOMIC STRUCTURE
There are two ways to represent the atomic
structure of an element or compound;
1. Electronic Configuration
2. Dot & Cross Diagrams
ELECTRONIC CONFIGURATION
With electronic configuration elements are represented
numerically by the number of electrons in their shells
and number of shells. For example;
Nitrogen configuration = 2 , 5
14
N
2 in 1st shell
2 + 5 = 7
5 in 2nd shell
7
ELECTRONIC CONFIGURATION
Write the electronic configuration for the following
elements;
40 23 16
a) Ca b) Na c) O
11 8
20
2,8,8,2 2,8,1 2,6
35 28 11
d) Cl e) Si f) B 5
17 14
2,8,7 2,8,4 2,3
DOT & CROSS DIAGRAMS
With Dot & Cross diagrams elements and compounds
are represented by Dots or Crosses to show electrons,
and circles to show the shells. For example;
X
14
N
Nitrogen X X N X X
7
XX
DOT & CROSS DIAGRAMS
Draw the Dot & Cross diagrams for the following
elements;
16 35
X
a) O b) Cl 17 X X
8 X
X X X
X
X X X X Cl X X X
X O X
X X
X X X
X X X
X
SUMMARY
1. The Atomic Number of an atom = number of
protons in the nucleus.
2. The Atomic Mass of an atom = number of
Protons + Neutrons in the nucleus.
3. The number of Protons = Number of Electrons.
4. Electrons orbit the nucleus in shells.
5. Each shell can only carry a set number of electrons.
