Chemical Foundations:
Elements, Atoms, and Ions
Chapters 4&5
Pages 85-155
1
Early History of Chemistry500 BC
• Greeks (around 500-400 BC)
– divided matter into four elements: earth, air,
fire, and water.
• Aristotle- Continuous theory of matter
matter can be divided and subdivided into
smaller and smaller parts indefinitely- each
part, no matter how small would retain the
original properties of matter.
• Democritus – Discontinuous theory of
matterbelieved all matter is made up of
tiny partices, or ATOMOS (indivisible),
which could not be broken down into
smaller particles
2
History- 500 BC1600s
• Alchemy- the belief that cheaper metals could be
turned into more valuable metals such as gold.
• Alchemists discovered many elements in their
attempts to convert metals to gold. Many elements’
names have Latin origins
Element
Element
Symbol
Origin
Word
Meaning
Gold
Au
Latin
aurum
Shining
dawn
Lead
Pb
Latin
plumbum
heavy
Bromine
Br
Latin
stench
3
History- 1600s
Robert Boyle (1627-1691)– Science should be
grounded in experiments
– Termed ELEMENTS- a
substance was an element
unless it could be broken
down into two or more
simpler substances
4
Key Concepts and Vocabulary:
Atoms, Elements and Molecules
• Atom: smallest unit of an element
• Element: substance that is composed of one type
of atom. There are114 elements on the periodic
table- 88 are naturally occurring, the remainder
have been made in laboratories.
• Molecule: a collection of atoms chemically
bonded.
– Molecules can be made from only one element, such as
H2 or O2
– Molecules can be made from different elements, such
as H2O or CO2
5
The Complexities of Molecules
• English language- composed of alphabet
containing 26 letters.
• Elements are the fundamental building
blocks of matter in the universe.
Distribution of the
most abundant
elements in Earth’s
crust
6
The Complexities of Molecules
One segment of a hemoglobin
molecule
7
History- 1800’s
1. Most natural materials are mixtures of
pure substances.
2. Pure substances are either elements or
combinations of elements called
compounds.
3. Law of Constant Composition-
8
Dalton’s Atomic Theory
•
1.
2.
3.
4.
5.
1766-1844- English School Teacher
Elements are made up of tiny particles called atoms.
All atoms of a given element are identical
The atoms of a given element are different from those of
any other element.
Atoms of one element can combine with atoms of other
elements to form compounds. A given compound
always has the same relative numbers and types of
atoms.
Atoms are indivisible in chemical processes. That is,
atoms are not created or destroyed in chemical reactions.
A chemical reaction simply changes the way the atoms
are grouped together.
9
Impact of Dalton’s Theory
10
Formulas and Compounds
• Compound-
11
• Chemical Formula-
• Elements’ symbols are used to simplify the
formulas. Subscripts are used to indicate
the number of each type of element.
GlucoseC6H11O6
12
History- 1800s –1900s
The Structure of the Atom
• JJ Thomson (1897) Cathode Ray TubeCredited with discovering the electron.
13
History- 1800s –1900s
The Structure of the Atom
• If negative charges are present and the
overall charge on an atom is neutral,
Thomson theorized that a positively charged
particle must also exist.
William
Thomson (Lord
Kelvin)- Plum
Pudding Model
of the Atom
14
History- 1900s
The Structure of the Atom
• Ernest Rutherford (1911)- Gold Foil Experiment
• Alpha Particles- + charged particles 7500 times
more massive than electrons.
15
History- 1900s
Rutherford’s Gold Foil Experiment
16
History- 1900s
• The Results of Rutherford’s Gold Foil Experiment
led scientists to believe that the atom must have a
densely packed positive center (NUCLEUS)
around which tiny electrons moved in a space that
was otherwise empty.
• 1919- Rutherford concludes that the atom’s
nucleus must have particles positively charged
called protons and that the number of protons in
an atom equal the number of electrons.
17
History- 1900s
The Nuclear Atom
• Rutherford’s collegue- James Chadwick
determined that the mass of the atom did not equal
the mass of electrons + protons. Therefore, there
must be another particle without a charge
(NEUTRON) found in the nucleus.
18
Introduction to the Modern
Concept of Atomic Structure
• From a very simplistic perspective- the atom is composed of
a tiny nucleus (10-13 cm in diameter) and electrons that move
about the nucleus at a distance of about 10-8 cm from it.
• If the nucleus was a golf ball, the electrons would be about 2
Km away.
• The nucleus contains protons and neutrons.
• The mass of the protons is about equal to the mass of
neutrons.
• Neutrons have zero charge, electrons –1, protons +1.
• The number of protons = the number of electrons in a neutral
atom.
19
Particle Relative Relative
Mass
Charge
Electron
1
-1
Proton
1836
+1
Neutron
1839
none
20
• X = the symbol of
the element
• A = mass number
(number of protons
+ neutrons)
• Z = atomic number
(number of protons)
21
The Hydrogen Atom
• One electron orbiting a
nucleus
• Z = atomic number
p
• N
e
1H
• Total mass =
• Singly ionized Hydrogen is
missing one electron = 1H+
22
The Helium Atom
• Two electrons orbiting a nucleus
e
p
n
n
p
e
4He
• Z = atomic number
• N
• Total mass =
• Singly ionized Helium is missing
one electron = 4He+
• Doubly ionized Helium is missing
both electrons = a particle = 4He++
23
Isotopes
• Dalton’s theory is modified as a result of Chadwick’s
studies of the nucleus and the discovery of the neutron.
• All atoms of the same element contain the same number of
protons and electrons, but atoms of a given element may
have different numbers of neutrons.
Isotopes-
24
Isotopes of Hydrogen
25
Isotopes of Carbon
26
Introduction to the Periodic Table
27
Periodic Table
• Arranged according to increasing atomic
number (number of protons)
• Horizontal Rows –
• Vertical Columns –
• This arrangement is based on chemical similarities
that exist in the vertical columns (groups). These
groups are referred to as
• This system of arrangement was 1st proposed by
Dmitri Mendeleev in 1869. His first table
consisted of 62 known elements. He was able to
predict the presence of several elements that had
28
not yet been discovered based on his table.
Periodic Table
• The name periodic
table refers to the fact
that as we increase the
atomic numbers, every
so often an element
occurs with properties
similar to those of an
earlier (lower atomic
number) element.
29
Chemical Families of the
Periodic Table
•
•
•
•
•
•
•
•
•
Metals
Nonmetals
Metalloids
The Alkali Metals
The Alaline Earth Metals
Transition Metals
The Oxygen Family
Halogens
Nobel Gases
http://www.papernapkin.com/
30
History –1920s
• Following Rutherford’s planetary model of the
atom, it was realized that the attraction between
the electrons and the protons should make the
atom unstable
• Neils Bohr (1922) proposed a model in which the
electrons would stably occupy fixed orbits, as
long as these orbits had special quantized
locations
31
History- 1920s
The Bohr Model of the
Atom
• Neils Bohr expands Rutherford’s
model of the atom allowing the
electrons to travel in successively
larger distinct orbits around the
nucleus. The outer orbits hold
more electrons than the inner
orbits and the outer orbits
determine the atom’s chemical
properties.
32
History –The Bohr Model
• In the Bohr model, the electron can
change orbits, accompanied by the
absorption or emission of a photon of a
specific color of light.
33
History- The Bohr Model
• Bohr’s Model was used to explain why the
negatively charged electrons did not fall into the
positively charged nucleus of the atom.
• The electrons were only able to occupy distinct
energy levels or orbits. “Quantized” energy levels.
34
History- The Bohr Model
• If an atom absorb a specific amount of
energy (quantum), the outer shell electrons
(valence electrons) could be excited into
higher energy states. This excited state is
unstable, so the electron releases a photon
of energy (quantized) as light. Light of
specific wavelength correspond to the
energy emitted by the electron dropping
back to the ground state.
35
History- The Bohr Model
http://www.mhhe.com/physsci/che
mistry/essentialchemistry/flash/lin
esp16.swf
36
Ions
• Ions –
• Neutral atoms become ions through the addition or
removal of electrons. Atoms form ions to gain
stability in their valence electron levels.
37
Ions
• Cations =
– Metals want to lose electrons to gain a stable octet
• Anions =
– Nonmetals want to gain electrons to gain a stable octet
38
Ions
• The resulting charge on the ion is
determined based on the number of
electrons gained or lost.
39
Ion Charges and the Periodic Table
• The position of an element in the periodic
table can help determine the resulting
charge on its ion.
40
Electrostatics
• Like charges repel
• Unlike charges attract
41
Compounds That Contain Ions
• Substances that want to lose electrons (metals) will react
with substances that want to gain electrons (nonmetals)
to forms ions. The resulting ions are attracted to each
other through electrostatic forces.
2Na + Cl2  2Na+ + 2Cl-  2NaCl
2 Sodium atoms transfers an electron each to 2 chlorine
atoms. The resulting ions (sodium and chloride) are
attracted to each other through electrostatic forces.
42
Ionic Bond = Electrostatic Force
43
Properties of Ionically Bonded
Substances
• 1. Composed of metals + nonmetals
• 2. High melting point. NaCl mpt = 8000C
• 3. Good conductors in liquid or aqueous
state. Do not conduct as solids.
44
Ionic Compounds
• A chemical compound must have a net
charge of zero if it is neutral!
• 1. There must be both positive ions (cations)
and negative ions (anions) present.
• The numbers of cations and anions must be
such that the net charge is zero.
NaCl
Na+
Cl45
Always check the substances location in the Periodic Table!
MgCl2
Mg2+
Cl-
Cl-
Li3N
Li+
Li+
Li+
N346
Nomenclature = Naming Compounds
• Binary Compounds- compounds that
contain two elements
47
Binary Ionic Compounds
• 1. The cation is always named first and the anion
second.
• 2. A simple cation (obtained from a single atom)
takes its name from the name of the element. For
example, Na+ is called sodium in the names of
compounds containing this ion.
• 3. A simple anion (obtained from a single atom) is
named by taking the first part of the element name
(the root) and adding –ide. Thus the Cl- ion is
called chloride.
48
Eamples- Binary Ionic Compounds
•
•
•
•
•
NaCl 
KI 
CaS 
CsBr
MgO 
49
Binary Ionic Compounds
Containing Transition Metals
• Many metals can form more than one type
of cation.
• Lead (Pb) can form Pb2+ or Pb4+
• Gold (Au) can form Au+ or Au3+
• Iron (Fe) can form Fe2+ or Fe3+
Chemists use Roman numerals to specify the
charge on the cation.
• Fe2+ = Iron II
Fe3+ = Iron III
50
Examples – Binary Ionic
Compounds with Transition Metals
•
•
•
•
•
CuCl 
HgO 
Fe2O3 
MnO2 
PbCl4 
51
Naming Compounds That
Contain Polyatomic Ions
• Polyatomic ions –
• In order to name these compounds, you
must memorize the names of the polyatomic
ions. Using Table F in your reference
tables will help you to do this.
52
Examples
•
•
•
•
•
Na2SO4 
KH2PO4 
Fe(NO3)3 
Na2SO3 
Mn(OH)2 
53
Naming Binary Compounds
That Contain Only Nonmetals
• 1. The first element in the formula is named first,
and the full element name is used.
• 2. The second element is named as through it were
an anion.
• 3. Prefixes are used to denote the numbers of
atoms present. (see next slide)
• 4. The prefix mono- is never used for naming the
first element. CO is carbon monoxide NOT
monocarbon monoxide.
54
Prefixes Used to Indicate
Numbers in Chemical Names
Prefix
Number Indicated
1
2
3
4
5
6
7
8
55
Examples
•
•
•
•
•
•
BF3 
NO 
N2O5 
CCl4 
NO2 
IF3 
56
Naming Acids
• Acids- a classification of molecules that
when dissolved in water produce H+ ions
(protons).
• 1. If the anion does not contain oxygen, the
acid is named with the prefix hydro- and the
suffix- ic attached to the root name for the
element.
• HCl (aq) 
• H2S (aq) 
57
Naming Acids Cont.
• 2. When the anion contains oxygen, the acid
name is formed from the root name of the
central element of the anions or the anion
name, with a suffix of -ic, or –ous. When
the anion name ends in –ate, the suffix ic is
used. When the anion name is -ite, the
suffix –ous is used.
• H2SO4 
• H2SO3 
58
59
Writing Formulas from Names
• Empirical Formulas- uses element symbols
to indicate the atoms or ions in a compound,
with subscripts to indicate their smallest
whole number ratio.
Calcium fluoride 
Potassium hydroxide 
Cobalt (III) nitrate 
60