Download Notes on Atomic Structure

Early models of the atom
 The
Greek philosopher Democritus (460
B.C. – 370 B.C.) was among the first to
suggest the existence of atoms.
 He believed that atoms were indivisible
and indestructible.
Early Models of the atom
 John
Dalton (1766-1844)
 Using experimental methods,
Dalton transformed the ideas of
Democritus into a scientific theory.
Early Models of the atom
 Dalton’s
Atomic Theory
1.) All elements are composed
of tiny indivisible particles called
atoms.
2.) Atoms of the same element
are identical. The atoms of any
one element are different from
those of any other element.
Early Models of the atom
 Dalton’s
 3.)
Atomic Theory
Atoms of different elements can physically
mix together or can chemically combine in
simple whole-number ratios to form
compounds.
 4.) Chemical reactions occur when atoms are
separated, joined or rearranged. Atoms of
one element, however, are never changed
into atoms of another element as a result of a
chemical reaction.
Modern Atomic Theory
 The
essence of Dalton’s theory
remains valid
 Today we know that atoms can be
destroyed via nuclear reactions but
not by chemical reactions
 There are also different kinds of
atoms within an element known as
isotopes
ATOMS
 Atoms.
An atom is the smallest particle of an
element that retains its identity in a chemical
reaction.
ATOMS
 Atoms
are very small
 An atom has 3 kinds of particles:
 Protons (positive charge), neutrons
(neutral), electrons (negative charge)
 Protons and neutrons are located in the
atomic nucleus (center of atom)
 Electrons are located around the
nucleus in an electron cloud and
occupies almost all of the volume of the
atom.
What forces keep an atom
together?
4




Fundamental Forces
Gravitational force of
nature
Electromagnetic force
of nature
Strong nuclear force of
nature
Weak nuclear force of
nature
 Stable
nucleus stays
together
Electromagnetic Forces
 Oppositely
charged
 Electrons are kept in the
particles attract each
orbit around the nucleus
other, while like particles by the electromagnetic
repel one another.
force, because the
nucleus in the center of
the atom is positively
charged and attracts the
negatively charged
electrons.
Strong Force
 The
strong forces oppose the
electromagnetic force of repulsion
between protons. Like ”glue” the strong
force keeps the protons together to form
the nucleus.
 The strong forces and electromagnetic
forces both hold the atom together.
Weak forces

Weak forces are
important because they
are responsible for
stabilizing particles
through the process of
radioactive decay.
Gravitational Force
 Gravity
is the force of
attraction exerted
between all objects in
nature.
 Gravity is most easily
observed in the behavior
of large objects. Inside
the tiny nucleus of an
atom, the effect of
gravity is small compared
to the effects of the other
three forces.
Subatomic Particles
 Properties
of Subatomic Particles
Particle
Symbol
Relative
Charge
Relative Mass
Actual Mass
Electron
e-
1-
1/1840
9.11 x 10-28
Proton
p+
1+
1
1.67x 10-24
Neutron
n0
0
1
1.67x 10-24
(mass of proton =
1)
Elements
 Elements
are the simplest form of
matter that has a unique set of
properties.
 Cannot be broken down into simpler
substances by chemical means.
How many known elements
are there?
Distinguishing Among Atoms
 Elements
are different because they
contain different numbers of protons.
 The atomic number of an element is the
number of protons in the nucleus of that
element.


Atomic number
Identifies an element.
Atomic number
Element
Atomic Number
Hydrogen (H)
1
Beryllium (Be)
4
Carbon (C)
6
Nitrogen (N)
7
Neon (Ne)
10
Number of Protons
Atomic number
 Atoms
are electrically neutral therefore
number of electrons= number of protons.
 Ions- have unequal number of protons
and electrons (+/-)
Element
Atomic
Number
Hydrogen
(H)
1
Beryllium
(Be)
4
Carbon (C)
6
Nitrogen
(N)
7
Neon (Ne)
10
Number of
Protons
Number of
Electrons
Mass number
 Remember
that most of the mass of an
atom is found in the nucleus (protons +
neutrons).
Mass number= # of protons + # of neutrons
 How
can we determine the number of
neutrons?



Protons= atomic number
Electrons= atomic number (if atom has no charge)
Neutrons= atomic mass – atomic number
Element
Atomic
Number
Mass
number
Hydrogen
(H)
1
1
Beryllium
(Be)
4
9
Carbon
(C)
6
12
Nitrogen
(N)
7
14
Neon
(Ne)
10
20
Number of
Protons/
Electrons
Number of
Neutrons
Ions and
Isotopes
Ions
 An


atom or molecule that has a charge.
Loss of electrons  Positive charge (cation)
Gain of electrons  Negative charge
(anion)
Ions
Ion
Atomic
#
Hydrogen
ion
1
Fluoride
ion
Symbol
Aluminum
ion
Number of Number of
Protons
Electrons
0
F-1
Oxide
ion
Calcium
ion
Cation or
anion?
10
Ca+2
10
What is different between the
three isotopes of hydrogen?
Isotopes

Isotopes= atoms of the same element that
have the same atomic number but different
atomic masses due to a different number of
neutrons.
Atomic Mass Units
 Today
it is possible to determine the mass
of an element using a mass spectrometer.
 However, numbers are small and
impractical to work with. (i.e. Mass of F=
3.155 x 10 -23 g)
 It is easier to use relative masses of atoms
using a reference isotope as a standard.
Atomic Mass Units
 The
reference isotope= Carbon 12. This
isotope is assigned a mass of exactly 12
amu.
 Atomic mass units= 1/12 of the mass of a
Carbon-12 atom.
Atomic Mass
 In
nature, most elements occur as a
mixture of two or more isotopes.
 Atomic
mass= the weighted
average mass of the atoms in a
naturally occurring samples of the
element.
 Takes into account the relative
abundance
Atomic Mass
 To
determine atomic mass, you must
know the number of isotopes, the mass of
each isotope and the percent
abundance of each isotope.
 Atomic
mass= multiply the mass of each
isotopes by natural abundance and add
the products.
Atomic Mass
Atomic mass= multiply the mass of each
isotopes by natural abundance and add
the products.
Atomic Mass of Carbon
 Example:
Carbon has 2 isotopes: Carbon12 which has a natural abundance of
98.89% (0.9889) and Carbon-13 which has
a natural abundance of 1.11% (0.0111).
 Atomic
Mass of carbon=
Atomic Mass of Magnesium
 Magnesium
has three naturally occurring
isotopes. 78.70% of Magnesium atoms
exist as Magnesium-24 (23.9850 g/mol),
10.03% exist as Magnesium-25 (24.9858
g/mol) and 11.17% exist as Magnesium-26
(25.9826 g/mol). What is the average
atomic mass of Magnesium?
 Atomic Mass of Magnesium=
Atomic Mass of Lithium
 What
is average atomic mass of Lithium if
7.42% exists as Li-6 (6.015 g/mol) and
92.58% exists as Li-7 (7.016 g/mol)?
 Atomic Mass of Lithium=