Windsor University
School of Medicine
ATOMS,
MOLECULES & IONS
TO KNOW WHAT IS RIGHT AND NOT TO DO IT IS THE WORST COWARDICE.
CONFUCIUS
Ch 2.b
J.C. Rowe
Atoms & Elements
An element is a substance which
cannot be broken down into simpler
substances by chemical means.
 An atom is the smallest particle of an
element which has the properties of
that element.

MARIE CURIE (Radioactivity)
One of the pieces evidence for the
fact that atoms are made of smaller
particles came from the work of
MARIE CURIE (1876-1934).
 She discovered RADIOACTIVITY, the
spontaneous disintegration of some
elements into smaller pieces.

Electrons (discovery)

J.J. Thompson determined charge : mass ratio of the
electron, e-, in 1897
 the
charge is -1
 the mass is 5.486 x 10 -4 amu or 9.10939 x 10 -31 kg


Robert Millikan measured the charge of e-, in 1909
Thompson developed the "plum pudding" model of
the atom
 this
model had all atomic particles distributed evenly
throughout the volume of the atom
The plum pudding model
the plum pudding model of the atom — negative charges (electrons)
embedded in a larger structure of positive charge — disproved by Ernest
Rutherford's gold foil experiment in 1911
Electron discovery Cont’d.

Rutherford developed the "nuclear" model of the atom


based upon his experiment which showed that atoms
contains regions of highly dense, positive material, called
the nucleus.
Rutherford discovered through his famous experiment
with gold foil in which he shot alpha particles (fairly
massive particles with a positive charge) through thin
gold foil and found that many particles were strongly
deflected and some bounced back at him!

This could only happen if the gold foil atoms contained
massive centers that had a positive charge, as exhibited in
the figure on the next slide.
Rutherford Experiment
shot of alpha particles (fairly massive particles with a positive
charge) through thin gold foil .
Rutherford atomic planetary model
the plum pudding model of the atom — negative charges (electrons) embedded in a
larger structure of positive charge — disproved by Ernest Rutherford's gold foil
experiment in 1911
Plum pudding model (Thompson)
Nuclear model (Rutherford)
Atomic Structure
Atomic Composition

PROTONS
Positive electrical charge
 Mass =1.67262 x 10 -27
 Relative mass = 1.007 atomic mass units (amu)


ELECTRONS
Negative electrical charge
 Relative mass = 0.0005 amu


NEUTRONS
No electrical charge
 Mass = 1.009 amu

ATOMIC PARTICLES (table)
Particle
Mass (kg)
Mass (amu)#
Charge*
#1
amu (atomic mass
unit) = 1.66054 x
10 -27 kg
* unit charge =
1.602 x 10 -19 C
(coulomb)
Electron
9.10939 x 10 -31
0.00055
-1
Proton
1.67262 x 10 -27
1.00728
+1
Neutron
1.67493 x 10 -27
1.00866
0
Atom composition
The atom is mostly empty space
 Protons & Neutrons are in the
nucleus
The
number of electrons = the
number of protons

Electrons are in space around the
nucleus
Electrons are in space around the nucleus
ATOMIC NUMBER



the atomic number has the symbol, Z, and is shown
as a subscript to the element symbol
the atomic number gives the number of protons in
the nucleus (and the number of electrons if the
species is neutral) of a particular atom
the atomic number defines a specific type of atom
since each different type of atom (representing
each element) will have a different number of
protons in the nucleus
MASS NUMBER

the mass number has the symbol, A, and is shown as
a superscript to the element symbol
 the
mass number gives the mass of atom in amu, atomic mass
number, and is approximately equal to the number of
protons plus the number of neutrons

Mass number = # protons + # of neutrons
ELEMENTAL SYMBOLS

Elemental symbols
are typically written
as below where X is
the element symbol
with the mass number
as a superscript and
the atomic number as
a subscript
A
 ZX
Quick Hit (to solve problems)



the atomic number always gives the number of protons
if the element is neutral (no charge), the number of
electrons will equal the number of protons
if the element is charged the number of electrons and
protons will be different
the numerical value of the charge is the difference between
the number of electrons and the number of protons
 the sign of the charge is positive if the number of protons is
greater than the number of electrons and negative if the
number of electrons is greater than the number of protons

Sample Problems Cont’d
SYMBOL
CHARGE
# of Protons
# of Neutrons
# of Electrons
0
15
16
15
31 P
15
(31-15)=
0
35
79 Br
35
35
(79-35)=
+2
55 Mn
25
44
25
30
23
(55-25)=
(25-2)=
ISOTOPES


Almost all atoms have "Isotopes“
Elements with the same number of protons (atomic
number) but differing number of neutrons – isotopes
are the same elements (atoms) with different masses
 isotopes
will have slightly different chemical and
physical properties due to the difference in mass, which
can be very helpful in characterizing substances
Example of three isotope of carbon, C
12 C
6
13 C
6
14 C
6
#p
6
6
6
#e
6
6
6
#n
6
7
8
Isotopes Cont.’d


because nearly all elements have one or more
isotopes, the mass of a naturally occurring element
will be a "weighted average" of all the isotopes
which occur naturally, for example:
Carbon has two prominent isotopes 12C6 &13C6
which occur naturally . their relative percent
abundance:
 C-12
12.0000 amu 98.9% abundant
 C-13 13.0000 amu 1.1 % abundant
Isotopes Cont.’d


C-12 12.0000 amu 98.9% abundant
C-13 13.0000 amu 1.1 % abundant
To determine the mass of naturally occurring
carbon, we calculate the weighted average of the
two isotopes by
summing (fractional abundance) x (mass of isotope) for
each isotope:
 mass C = (0.989)(12.0000) + (0.011)(13.0000) =
12.011 amu

ATOMIC MASSES

Because atoms are too small to weigh individually,
a relative mass scale has been developed for
elements on the periodic table.
 amu
– atomic mass unit – is one such relative mass
scale—
 one amu equals exactly 1/12 the mass of an atom of
carbon-12 isotope

rounded masses: C has 12 amu; Mg has 24 amu
Al has 27 amu; F has 19 amu
Molecules & Compounds
Molecules
 Ionic compound
Molecular compound

1. Elements vs. Compounds
Compounds and elements are pure substances which are the basic
building blocks of all matter. Notice that both elements and compounds
are pure substances.
2. Atoms/element vs. Molecules/compounds



Atoms are the smallest particles that can be
identified as a particular element,
and molecules are the smallest particles that can be
identified as a particular compound.
Elements can also occur in a molecular form in which
the same type of elements (atoms) are chemically
combined, such as two oxygen atoms, O, which form
molecular oxygen, O2, when chemically combined.
3. Diatomic molecules

There are some elements, which you should know, that
only occur naturally (under normal conditions) in their
molecular forms.



They are called diatomic molecules or, sometimes, molecular
elements and they are: hydrogen, H2; nitrogen, N2; oxygen,
O2; fluorine, F2; chlorine, Cl2; bromine, Br2; iodine, I2.
Compounds are chemical combinations of elements
(atoms) of different types, such as water, H2O or
carbon dioxide, CO2.
Let us look further at the formation of compounds,
both molecular compounds and ionic compounds.
4. Molecular Compounds

molecules are formed from chemical combinations
of atoms; atoms are combined in specific ratios to
one another:
eg. water is H2O with a ratio of 2 : 1 in H : O

H2O2 is not water -- it is hydrogen peroxide with a
ratio of 2 : 2 in H : O
5. Ionic Compounds

molecular substances are compounds formed
between different non-metal elements
 the
molecules that make up molecular substances are
individual units which act independently but are
identical to one another.

ionic substances are formed between metals and
non-metals and are quite different from molecular
substances
Masses of Molecules

masses of molecules or compounds are simply the
summed masses of all atoms or elements which combine
to form a molecule or compound and are represented
by the chemical formula
eg. alcohol
vitamin C
sugar
saccharin
aspirin
cocaine
C2H6O
C6H8O6
C12H22O11
C7H5NO3S
C9H8O4
C17H21NO4
ether C2H6O
Cont’d.

The mass of a compound is the sum of the masses of
all the atoms that combine to form the compound:
Mass of vitamin C is 176.08 amu
6C =
72.00 amu
8H =
8.08 amu
6O =
96.00 amu
C6H8O6 176.08 amu
Chemical Formula

The formulas give the type, ratio and number of
atoms in the chemical combination but they say
nothing of the actual structure,

for example, alcohol and ether are very different
although they have the same chemical formula
 alcohol
 ether
C 2 H6 O
C2H6O
Structural vs. chemical Formulas
A structural formula will show not only type, ratio and number atoms in the
chemical formula, but also which atoms are attached to which atomsof
Chemical Equations
Depict the kind of reactants & products
and their relative amounts in a reaction.
 4 Al (s) + 3 O2 (g)
2Al2O3 (s)
 The numbers in the front are called
stoichiometric coefficients.
 the letters (s), (g), and (l) are the physical
states of compounds.

Chemical equations cont’d.

4 Al(s)+3O2(g)

This equation means

2 Al2O3 (s)
4 Al atoms + 3 O2 molecules
2 molecules of Al2O3
or
4
moles of Al + 3 moles of O2
2 moles of Al2O3
Law of the Conservation of Matter



Because of the same atoms are present in a
reaction @ the beginning & @ the end, the amount
of matter in a system doesn’t change.
Because of the principle of the conservation of
matter, an equation must be balanced.
It must have the same number of atoms of the same
kind on both sides.
Part2. Properties of Ionic compounds
A metal atom can transfer an electron to a
nonmetal
 The resulting cation & anion are attracted to
each other by electrostatic forces.

 Cation
has a positive (+) sign
 Anion has a negative (-) sign

The oppositely charged ions in ionic compounds
are attracted to one another by electrostatic
forces.
ions are substances that have either a positive
or negative charge



cations
have a positive charge of
one or greater
are generally derived from
either metal elements or
groups of elements from
which one or more electrons
have been removed

cations (monotomic) are
always smaller that the
element from which they
are derived



anions
have a negative charge of
one or greater
are generally derived from
either non-metal elements or
groups of elements to which
one or more electrons have
been added

anions (monotomic) are
always larger than the
element from which they are
derived
Monotomic Ions.

monotomic ions are derived from single elements
 examples:
Na
Ca
Na+
®
Ca2+ +
®
Cl + 1e- ®
O + 2e-
+
Cl®
1e-
(e- is lost)
2e-
(e- 's are lost)
(e- is gained)
O2- (e- 's are gained)
Polyatomic ions


polyatomic ions are derived from groups of elements which
are generally non-metals
 examples: CO32- ; PO43- ; SO42- ; NH4+
Ionic substances
 are formed between oppositely charged ions : cations and
anions
 are held together by ionic bonds which are due to the
electrostatic attractions between the opposite charges
 ionic compounds are always neutral species formed by
combining the same number of positive and negative
charges:
for example: Mg2+ and Cl- produce MgCl2 not MgCl or
MgCl3 or any other combination
Electrostatic forces
Electrostatic forces are governed by Coulomb’s
law
 Coulomb’s law :


Force of attraction = (charge on +) x (charge on -)
divided by (distance between ions)^2.
As ions charge increases, the attractive force
increases.
 As the distance between ions increases, the
attractive force decreases.

Upton Sinclair
It is difficult to get a man to understand
something when his job depends on not
understanding it.