The document discusses the atomic theory of matter and the development of atomic structure models. It describes John Dalton's atomic theory which stated that elements are composed of atoms that are unique and atoms are neither created nor destroyed in chemical reactions. The discovery of the electron by J.J. Thompson and experiments by Robert Millikan and Ernest Rutherford helped develop the modern atomic structure model of a small, dense nucleus surrounded by electrons. The document also discusses isotopes, atomic numbers, mass numbers, and how the periodic table is arranged based on atomic structure.

1. ATOMS, MOLECULES
AND IONS
QBA Miguel A. Castro Ramírez

2. THE ATOMIC THEORY OF MATTER
The theory that atoms are the fundamental building blocks
of matter reemerged in the early 19th century,
championed by John Dalton.

3. THE ATOMIC THEORY OF MATTER
Dalton atomic theory:
• Each element is composed of extremely small particles called
atoms
• All atoms of a given element are identical to one another in mass
and other properties, but the atoms of one element are different
from the atoms of all other elements
• Atoms of an element are not changed into atoms of a different
element by chemical reactions; atoms are neither created nor
destroyed in chemical reactions. E.g : electrolysis
• Compounds are formed when atoms of more than one element
combine; a given compound always has the same relative number
and kind of atoms.

4. THE ATOMIC THEORY OF MATTER
Law of Chemical Combination:
• Law of Constant Composition: In a given compound, the relative
numbers and kinds of atom are constant
- The basis of Dalton’s postulates 4.
• Law of the Conservation of Mass (Law of Conservation of Matter):
- The total mass of materials present after a chemical
reaction is the same as the total mass present before reaction.
- Atoms can neither be created nor destroyed in a chemical
rxn; thus the total mass remains unchanged
Therefore,
Total mass of reactants = Total mass of the products
E.g. the total mass of element A and B is equivalent to the
total mass of C and D in the reaction:
A+B C+D
- Basis for Dalton’s postulates 3.

5. THE ATOMIC THEORY OF MATTER
• Law of Multiple proportions : If two elements A & B
combine to form more than one compound, the masses of B
that can combine with a given mass A are in the ratio of small
whole numbers.
Example:
1) CO and CO2
For this example, the number of carbon atom is
constant (fixed mass of the both element). Therefore, the
ratio of oxygen atoms in CO and CO2 compound is always 1:2
2) H2O and H2O2
For this example, the number of hydrogen atom is
constant (fixed mass of both element). Therefore, the ratio of
oxygen atoms in H2O and H2O2 compound is always 1:2

6. THE DISCOVERY OF ATOMIC STRUCTURE
Behavior of charged particles: Particles with the same charge
repel one another, whereas particles with unlike charges
attract one another.
The electron
• Streams of negatively charged
particles were found to emanate from
cathode tubes.
• J. J. Thompson is credited with their
discovery (1897).
• Thompson measured the charge/mass
ratio of the electron to be 1.76 x 108
coulombs/g.

7. THE DISCOVERY OF ATOMIC STRUCTURE
Millikan Oil Drop Experiment
• Once the charge/mass ratio of the
electron was known (previous
finding), determination of either
the charge or the mass of an
electron would yield the other.
• Robert Millikan (University of
Chicago) determined the charge
on the electron (1.602 x 109 C) in
1909.
• So, he calculated the mass of the
electron:
1.602 x 109 C
Electron mass = = 9.10 x 10- 28 g
1.76 x 108 C/g

8. THE DISCOVERY OF ATOMIC STRUCTURE
Radioactivity
• Radioactivity: Spontaneous emission of radiation by an atom
• It was first observed by Henri Becquerel.
• Marie and Pierre Curie also studied it
• Three types of radiation were discovered by Ernest Rutherford:
 α , β and γ
• Each types differ in its response to an electric field.
• Rutherford showed, both α and β rays consist of fast moving
particles, which were called α and β particles.
• α- positive charge, β – negative charge and γ - no charge

9. THE DISCOVERY OF ATOMIC STRUCTURE
The Atom
Circa 1900
• Thompson: Atom consisted of a
uniform positive sphere of matter in
which the electrons were embedded.
• The prevailing theory was that of the
“plum pudding” model.
• It featured a positive sphere of
matter with negative electrons
embedded in it.
• Ernest Rutherford proved this model
is wrong

10. THE MODERN VIEW OF ATOMIC STRUCTURE
• An atom comprises of protons, neutrons and electrons. (also called
subatomic particles)
• Charge of electron: -1.602 x 10-19 C
• Charge of proton: + 1.602 x 10-19 C
• 1.602 x 10-19 = Electric charge
• For convenience, the charges of atomic and subatomic particles are
expressed as multiple of this charge rather than in coulomb.
• Thus, electron = 1-, proton= 1+
• Neutron has no charge
• Atoms: Equal number of electrons n protons, so atoms have no
electrical charge
• Protons and neutrons are packed into a tiny positively charged core
of the atom known as the nucleus.

11. THE MODERN VIEW OF ATOMIC STRUCTURE
• Atomic mass unit (AMU) : 1 amu = 1.66054 x 10 -24 g
• Atoms are also small
• Convenient way to express atomic dimensions is agstrom (Å)
- 1 angstrom = 10-10 m

12. THE MODERN VIEW OF ATOMIC STRUCTURE
Atomic Numbers, Mass Numbers and Isotopes
• What makes an atom of one element different from another element?
• Bcoz of the subatomic compositions
• Atomic Number: Number of protons in the nucleus of an atom of
any particular element
• No of protons = no. of electrons
• Mass number: Total number of protons + electrons in the atom
X = Atomic symbol of the element
A = mass number; Proton + Neutron
Z = atomic number
(the number of protons or electrons)
THE SYMBOL OF THE ATOM OR ISOTOPE

13. THE MODERN VIEW OF ATOMIC STRUCTURE
Isotopes
• Atoms with identical atomic numbers but different mass
numbers
• Isotopes have different numbers of neutrons.
11 12 13 14
6 C 6 C 6 C 6 C

14. ATOMIC WEIGHTS
The Atomic Mass Scale
• 1 amu = 1.66054 x 10-24 and 1g= 6.02214 x 1023
• The atomic mass unit is presently defined by assigning
a mass of exactly 12 amu to an atom of the 12C isotope
of the carbon.
• In these units, an 1H atom has a mass of 1.0078 amu
and an 16O atom has a mass of 15.9949 amu.

15. ATOMIC WEIGHTS
Average Atomic Masses
• We can determine the average atomic mass of an element by using the
masses of its various isotopes and their relatives abundance
• Because in the real world we use large amounts of atoms and molecules, we
use average masses in calculations
How to Calculate:
The 3 naturally occurring isotopes of neon are; neon-20, 90.51%,
19.99244 amu; neon-21, 0.27%, 20.99395 amu; neon-22, 9.22%, 21.99138
amu. Calculate the weighted average atomic mass of neon.
Solution:
(0.9051 x 19.99244 amu) + (0.0027 x 20.99395 amu) + (0.00922 x 21.99138
amu) = 20.1794 amu
• Average atomic mass (expressed in atomic mass units) is also known as its
atomic weight.

16. ATOMIC WEIGHTS
Average Atomic Masses
EXERCISE
The 2 naturally occurring isotopes of copper are copper-63,
mass 62.9298u, and copper-65, mass 64.9278u. What must
be the percent natural abundances of the 2 isotopes if the
atomic mass of copper listed in a table of atomic masses is
63.546u?

17. THE PERIODIC TABLE
1. The modern P.T. is based on Mendeleev’s earlier work.
2. Elements are organized and categorized according to
their physical and chemical properties and in the order of
increasing atomic number.
3. Each vertical column is called a group. The main group is
Group 1 to 8. Elements from the same group have similar
physical and chemical properties.
4. Periods are the horizontal rows of elements, labeled 1 to
7.

18. THE PERIODIC TABLE

20. THE PERIODIC TABLE
• For each element in the table, the atomic number and atomic
symbol are given.
• May notice slightly difference and variations in PTs from 1
book to another. These are simply matters of style, or they
might concern the particular information included. NO
FUNDAMENTAL DIFFERENCES.
• Elements that belong to the same group often exhibit
similarities in physical and chemical properties.
Example: ‘Coinage metal’- Cu, Ag and Au- belong to
group 1B.
Used throughout the world to make coins.

21. THE PERIODIC TABLE
• Elements on the left and the middle side of the PT are metallic
elements or metal
• Share characteristic properties such as luster and high
electrical and heat conductivity.
• All metals, with the exception of mercury (Hg), are solids at
room T.
• The metals are separated from nonmetallic elements or
nonmetals by a diagonal step-like line.
• The elements that lie along the line that separates metals
from non-metals, such as (Sb), have properties that fall
between those of metals and those of non-metals.
• Referred as metalloids.
- Tend to be semiconductors – Conduct electricity, but not
nearly so well as metals.

22. THE PERIODIC TABLE
Names of Some Groups in the Periodic Table

23. MOLECULES AND MOLECULAR COMPOUNDS
Molecules and Chemical Formulas
• Many elements are found in nature in molecular form
Examples: Oxygen – O2 Ozone- O3
• Compounds that are composed of molecules contain more
than one type of atom and are called molecular compounds.
Example: Water –H2O Hydrogen peroxide- H2O2
• Even though composed of same elements, both water and
hydrogen peroxide have different properties.
• Most molecular substances that we will encounter contains
only non-metals.

24. MOLECULES AND MOLECULAR COMPOUNDS
Molecular and Empirical Formula
• Chemical formulas that indicate the actual
numbers and types of atoms in a molecule
are called molecular formula.
• Empirical formulas give the lowest whole-
number ratio of atoms of each element in a
compound
Example: Molecular Formula: H2O2
Empirical Formula: HO
• Molecular formula provide more
information about molecules than do
empirical formula

25. MOLECULES AND MOLECULAR COMPOUNDS
Picturing Molecules
• The structural formula of a substance
shows which atoms are attached to which
within the molecule.
• Perspective drawings also show the three-
dimensional array of atoms in a compound.
• Examples: Water and hydrogen peroxide:
O
HO OH
H H
• Atoms are represented by their chemical
symbols and lines are used to represent the
bonds.

26. MOLECULES AND MOLECULAR COMPOUNDS
• Structure formula does not depict the actual geometry of the
molecule.
• Perspective drawings can be written to show the three-
dimensional array of atoms in a compound
• Ball-and-stick model show atoms as spheres and bonds as
sticks.
- The advantage of accurately representing the angles at
which the atoms are attached to one another within the
molecule.
• Space-filling model: Molecule would look like if the atoms were
scaled up in size.
- These models show the relative sizes of the atoms, but
the angles between atoms, which help define their
molecular geometry, are often more difficult to see than
in ball-and-stick models.

27. IONS AND IONIC COMPOUNDS
Predicting Ionic Charges
• The nucleus of an atom is unchanged by chemical processes, but some
atoms can readily gain or lose electrons.
• When atoms lose or gain electrons, they become ions.
– Cations are (+) and are formed by elements on the left side of the
periodic chart. e.g: Na+, Ca2+
– Anions are (-) and are formed by elements on the right side of the
periodic chart. e.g: Cl- , Br-

28. IONS AND IONIC COMPOUNDS
Ionic Compounds
• Sodium Chloride (NaCl): Example of an ionic compound.
- Contains both + and – charges
• Ionic Compounds: Generally combinations of metals
and nonmetals (NaCl).
•In contrast, molecular compounds are generally
composed of non-metals only (H2O)

29. IONS AND IONIC COMPOUNDS
The Formation of Ionic Compounds
• The atoms transfer electrons between them when they react
e-
• Na + Cl  Na+ + Cl-

30. IONS AND IONIC COMPOUNDS
The Formation of Molecular Compound
• Molecular compounds: Formed when non-metallic elements
combine
• The attraction that hold atoms to each other in molecules,
called chemical bonds, are electrical in nature.
•They arise from sharing of electron between one atom to
another.

31. NAMING INORGANIC COMPOUNDS
• System used in naming the substances: Chemical Nomenclature
• Organic Compounds: Contain C, usually in combination with
H,N,S or O. (CH4)
• Inorganic Compounds: All others (Al2O3)
Names and Formulas of Ionic Compounds
1) Positive Ions (cation)
• Cation formed from metal atoms have the same name as the
metal.
• Na+ sodium ion
•If a metal can formed different cations, the + charge is
indicated by a Roman numeral in parentheses following
the name of the metal
• Fe2+ iron(II) ion Fe3+ iron(III) ion

32. NAMING INORGANIC COMPOUNDS

33. NAMING INORGANIC COMPOUNDS
- Polyatomic or cations with variable - Monoatomic ions
charges - Do not have variable charges
Formula of the most common ions

34. NAMING INORGANIC COMPOUNDS

35. NAMING INORGANIC COMPOUNDS
• Anions derived by adding H+ to an oxyanion are named
by adding as a prefix the word hydrogen or dihydrogen,
as appropriate:
•CO32- carbonate ion
• HCO3- hydrogen carbonate ion (bicarbonate)
• H+ reduced the – charge of parent anion

36. NAMING INORGANIC COMPOUNDS
3) Ionic Compound
• Cation followed by anion
• CaCl calcium chloride
• Al(NO3)3 aluminium nitrate
Exercise
a. Give the names of the following ions: (i) O22- and (ii) SO42-
b. Write down the formulas of the following ions: (i)
ammonium and (ii) nitrate

37. NAMING INORGANIC COMPOUNDS
PROBLEM: Give the systematic names or the formula or the formulas
for the names of the following compounds:
(a) Fe(ClO4)2 (b) sodium (c) Ba(OH)2 8H2O
sulfite
PLAN: Note that polyatomic ions have an overall charge so when
writing a formula with more than one polyatomic unit, place
the ion in a set of parentheses.
SOLUTION: (a) ClO4- is perchlorate; iron must have a 2+ charge. This
is iron(II) perchlorate.
(b) The anion sulfite is SO32-; therefore you need 2 sodiums
per sulfite. The formula is Na2SO3.
(c) Hydroxide is OH- and barium is a 2+ ion. When water is
included in the formula, we use the term “hydrate” and a
prefix which indicates the number of waters. So it is barium
hydroxide octahydrate.

38. NAMING INORGANIC COMPOUNDS
Names and Formulas of Acids
• If the anion in the acid
ends in -ide, change the
ending to -ic acid and
add the prefix hydro- .
– HCl: hydrochloric
acid
– HBr: hydrobromic
acid
– HI: hydroiodic acid

39. NAMING INORGANIC COMPOUNDS
• If the anion in the acid ends
in -ite, change the ending
to -ous acid.
– HClO: hypochlorous acid
– HClO2: chlorous acid

40. NAMING INORGANIC COMPOUNDS
• If the anion in the acid ends
in -ate, change the ending to
-ic acid.
– HClO3: chloric acid
– HClO4: perchloric acid

41. NAMING INORGANIC COMPOUNDS
Names and Formulas of Binary Molecular Compounds
1. The name of the less electro(-) - written 1st
- exception for compound contain oxygen.
- O always written last except when combined with fluorine
2. If both elements are in the group in PT, higher atomic no
comes 1st.
3. Name of 2nd element (more electro(-))is given –ide ending
4. Greek prefixes are used to indicate the no of each element
CO2: carbon dioxide
CCl4: carbon tetrachloride
5) If the prefix ends with a or o and the name of the element
begins with a vowel, the two successive vowels are often
elided into one.
N2O5: dinitrogen pentoxide

42. SOME SIMPLE ORGANIC COMPOUNDS
• Organic chemistry is the study of carbon.
• Organic chemistry has its own system of nomenclature.
• The simplest hydrocarbons (compounds containing only
carbon and hydrogen) are alkanes.
• The first part of the names above correspond to the
number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).

43. SOME SIMPLE ORGANIC COMPOUNDS
• When a hydrogen in an alkane is replaced with something
else (a functional group, like -OH in the compounds above),
the name is derived from the name of the alkane.
• The ending denotes the type of compound.
– An alcohol ends in -ol.

45. •Alkene
•alkuna