Chemical Foundations: Elements, Atoms, and Ions
Chapter 4
There are 117 known elements.
 Element 117, ununheptium, has not yet been
discovered, however, elements 115, 116, and
118 have been discovered.
 88 of the elements are naturally occurring.
 The elements were first defined by Robert
Boyle (1627-1691) after many experiments
 READ: “Chemistry in Focus- Trace Elements:
Small but Crucial” in book

Chapter 4
Element symbols are used to abbreviate the
name of the element.
 Some are one letter, two letters, or three
letters.
 The first letter is always a capital letter, the
other one (two) are lowercase.
 Example: Helium- He

Chapter 4
Dalton: History Report
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English Scientist and teacher
Son of a poor English weaver
Started a school in his village at only 12 years old
Consumed with the study of the atmosphere
Kept meticulous records of the weather for most of his life
Never married
Loved to bowl- did every Thursday afternoon
Poor public speaker
Unsuccessful as a lecturer, made his living as a private tutor
Physical Characteristics- tall, gaunt, rather unattractive
1.
2.
3.
4.
5.
Elements are made 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.

Law of Constant Composition means that a given compound
always has the same composition, regardless of where it comes
from.


Ex. Sodium chloride is always NaCl.
Dalton’s Model (1808)

Successfuly explained the law of consistent composition

Not accepted immediately

Dalton was sure he was right and used his model to predict how a given pair of
elements might combine to form more than one compound

Example: Nitrogen and Oxygen
NO
N2O
NO2
Read…

“Chemistry in Focus: No Laughing Matter” in
book on page 79
Chapter 4
A compound is a distinct substance that is
composed for the atoms of two or more
elements and always contains exactly the same
relative masses of those elements.
 The types of atoms and the number of each
type in each unit (molecule) of a given
compound are conveniently expressed by a
chemical formula.

1.
2.
3.
Each atom present is represented by its
element symbol.
The number of each type of atom is indicated
by a subscript written to the right of the
element symbol.
When only one atom of a given type is present,
the subscript 1 is not written.
Examples of Writing Formulas…
• Acid Rain SO3
– How many S’s?
– How many O’s?
• CO2
– How many C’s?
– How many O’s?
• H2O
• Sugar C6H12O6
– How many C’s?
– How many H’s?
– How many O’s?
– How many H’s?
– How many O’s?
• N2O5
– How many N’s?
– How many O’s?
Examples of Writing Formulas…

Writing formulas for compounds given in written
form:
A
molecule contains 4 phosphorous atoms and 10
oxygen atoms
A
molecule contains one uranium atom and six
fluorine atoms
A
molecule contains one aluminum atom and three
chlorine atoms
Chapter 4
Dalton’s Atomic Theory convinced scientists
that elements consisted of atoms.
 This led to many scientists pondering the
nature of the atom…
 A physicist in England named J.J. Thomson
showed in the late 1890s that the atoms of any
element can be made to emit tiny negative
particles.
 He concluded that all types of atoms must
contain these negative particles, which are now
called electrons.

Read: “Chemistry in Focus- Glowing Tubes for
Signs, TV Sets and Computers” (page 83)
 Lord Kelvin
suggested that the atom is like “plum
pudding”.
The atom is a uniform “pudding” of positive
charge with enough negative electrons
scattered within to counterbalance that
positive charge.


Ernest Rutherford- Gold Foil Experiment
Concluded that Lord Kelvin was incorrect in his
“plum pudding” model of the atom.
Explained the nuclear atom – an atom with a
dense center of positive charge (the nucleus).
Discovered that the nucleus contained protons –
the same size of charge, but opposite sign.
Rutherford and a coworker, James Chadwick,
were able to show that most nuclei also contains
a neutral particle that they named the neutron.
A neutron is slightly more massive than a
proton but has no charge.
Rutherford’s Gold Foil Experiment
Main area of interest- alpha particles- positively
charged particles with a mass approx. 7500
times larger than the electron.
 Studied the flight of alpha particles through air
 Realized that something was deflecting the
particles in the air
 To figure this out- created an experiment which
involved directing the alpha particles toward a
thin metal foil (look at figure 4.5 on page 82)

Rutherford’s Findings
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Most of the particles passed straight through
Some were deflected at large angles and some backwards
Plum pudding model incorrect- if it was correct and there was
just electrons floating around there would be minor deflections
if any
Large deflections could only be caused by a center of
concentrated positive charge that would deflect the positively
charged particles
Most particles went right through because the atom is mostly
open space with a concentrated center
Must be a nucleus containing protons of positive charge to
balance out the negative electron
Most nuclei also contain a neutral particle that is slightly bigger
than a proton but having no charge
Chapter 4
More discoveries since Thompson and
Rutherford…
 In this model, the atom is called a nuclear atom
because the positive charge is localized in a
small, compact structure (the nucleus) and not
spread out uniformly, as in the plum pudding
view.
 The chemistry of an atom arises from its
electrons (electronic structure).

Particle
Relative Mass
Relative Charge
Location
Electron
1
1-
Outside the nucleus
Proton
1836
1+
Nucleus
Neutron
1839
0
Nucleus
Chapter 4
All atoms of the same element have the same
number of protons (the element’s atomic
number) and the same number of electrons.
 In a free atom, the positive and negative
charges always balance to yield a net zero
charge.

Isotopes are atoms with the same number of
protons but different numbers of neutrons.
 Atomic Number is the number of protons in an
atom. It is also the number of electrons in an
neutral atom.
 Atomic Mass is the number of protons and the
number of neutrons in an atom.
 To get the number of neutrons in an atom,
subtract the Atomic Number from the Atomic
Mass.

X = Symbol of the Element
 A = Atomic Mass
 Z = Atomic Number

Calculating Protons, Neutrons and Electrons…
Book problems on pages 86-88
 Calculating protons, neutrons and electrons
worksheet

Chapter 4
The Periodic Table
Dmitri Mendeleev actually arranged the
elements in order of increasing atomic mass
and increasing atomic number.
 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.

These families of elements with similar
chemical properties that lie in the same vertical
column on the periodic table are called groups.
 Several groups have special names:

 Group
1 – Alkali Metals
 Groups 2 – Alkaline Earth Metals
 Group 7 – Halogens
 Group 8 – Noble Gases
 Groups 3 – 12 – Transition Metals
Most elements are metals.
 Metals are located to the left of the “staircase”
 Properties of Metals

 Efficient
conduction of heat and electricity.
 Malleability (they can he hammered into thin
sheets)
 Ductility (the can be pulled into wires)
 A lustrous (shiny) appearance
Nonmetals are those elements to the right of
the “staircase”.
 Nonmetals sometimes have one or more
metallic properties.
 Metalloids, or semimetals, lie on the
“staircase”.
 Metalloids include: Si, Ge, As, Sb, Te, Po, At, Al,
B

Periodic Trends

Atomic Radius - Atomic radius is simply the radius of the atom, an indication of the
atom's volume.

Period - atomic radius decreases as you go from left to right across a period.


Group - atomic radius increases as you go down a group.


Why? Stronger attractive forces in atoms (as you go from left to right) between the opposite charges in
the nucleus and electron cloud cause the atom to be 'sucked' together a little tighter.
Why? There is a significant jump in the size of the nucleus (protons + neutrons) each time you move
from period to period down a group. Additionally, new energy levels of elections clouds are added to the
atom as you move from period to period down a group, making the each atom significantly more
massive, both is mass and volume.
Electronegativity - Electronegativity is an atom's 'desire' to grab another atom's electrons.

Period - electronegativity increases you go from left to right across a period.


Why? Elements on the left of the period table have 1 - 2 valence electrons and would rather give those
few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's
electrons. As a result, they have low electronegativity. Elements on the right side of the period table only
need a few electrons to complete the octet, so they have strong desire to grab another atom's electrons.
Group – electronegativity decreases as you go down a group.

Why? Elements near the top of the period table have few electrons to begin with; every electron is a big
deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have
so many electrons that loosing or acquiring an electron is not as big a deal. This is due to the shielding
affect where electrons in lower energy levels shield the positive charge of the nucleus from outer
electrons resulting in those outer electrons not being as tightly bound to the atom.
Periodic Trends

Ionization Energy - Ionization energy is the amount of energy required to remove the
outmost electron. It is closely related to electronegativity.

Period - ionization energy increases as you go from left to right across a period.


Group - ionization energy decreases as you go down a group.

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Why? Elements on the right of the chart want to take others atom's electron (not given them up) because
they are close to achieving the octet. The means it will require more energy to remove the outer most
electron. Elements on the left of the chart would prefer to give up their electrons so it is easy to remove
them, requiring less energy (low ionization energy).
Why? The shielding affect makes it easier to remove the outer most electrons from those atoms that
have many electrons (those near the bottom of the chart).
Reactivity - refers to how likely or vigorously an atom is to react with other substances.
This is usually determined by how easily electrons can be removed (ionization energy)
and how badly they want to take other atom's electrons (electronegativity) because it is
the transfer/interaction of electrons that is the basis of chemical reactions.

Metals



Period - reactivity decreases as you go from left to right across a period.
Group - reactivity increases as you go down a group
Why? The farther to the left and down the periodic chart you go, the easier it is for electrons to be given
or taken away, resulting in higher reactivity.
Nonmetals


Period - reactivity increases as you go from the left to the right across a period.
Group - reactivity decreases as you go down the group.
Why? The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a
more vigorous exchange of electron
Chapter 4
Gold, silver, and platinum are called noble metals,
because they are relatively unreactive.
 Helium, neon, argon, krypton, xenon, and radon
are called noble gases because they are relatively
unreactive.
 Seven elements are called diatomic because they
are found as a molecule in nature: H2, N2, O2, F2,
Cl2, Br2, I2.
 Sulfur is found in nature as S8.
 Phosphorus is found in nature as P4.
 Two elements are liquids, Br and Hg, at room
temperature.

Chapter 4
Ions are charged atoms created by the gain or
loss of electrons.
 If an atom loses electrons, the atom takes on a
positive charge because there are more
protons than electrons and forms a positive ion
called an cation.
 If an atom gains electrons, the atom takes on a
negative charge because there are more
electrons than protons and forms a negative
ion called an anion.

The size decreases dramatically when an atom
loses one or more electrons to form a positive
ion.
 The size increases dramatically when an atom
gains one or more electrons to form a negative
ion.
 The name of an anion is obtained by adding –
ide to the root of the name. (ex. Chlorine
becomes chloride)

Chapter 4
Melting means that the solid, where the ions
are locked into place, is changed to a liquid,
where the ions can move.
 A substance containing ions that can move can
conduct an electric current.
 An ionic compound cannot contain only anions
or only cations, because the net charge of a
compound must be zero.
 The net charge of a compound (zero) is the sum
of positive and negative charges.

Naming Ions

Any anion (negatively charged ion) gets is name
from the root of the neutral atom plus the suffix
–ide
 Ex:
Cl- is the chloride ion
 Ex: F- is the fluoride ion
 Ex: O2- is the oxide ion

Cations keep their original names and have the
word ion added
 Ex:
Mg2+ is the magnesium ion
Ionic Charges and the Periodic Table

Trends:
 Group
1 tend to form +1 ions
 Group 2 tend to form +2 ions
 Group 3 tend to form +3 ions
 Group 4 tend to form +/-4 ions
 Group 5 tend to form -3 ions
 Group 6 tend to form -2 ions
 Group 7 tend to form -1 ions
 Group 8 are stable and do not form ions
Ionic Charges and the Periodic Table cont..
Transition metals- most form cations with
various (+) charges
 Note:

 metals
always form positive ions (have a tendency
to lose electrons)
 nonmetals on the other hand tend to gain electrons
and form anions (negatively charged ions)
Assigning Oxidation Numbers…

Practice on the board with Mrs. Harlan
Electron
Distribution in
Molecules
G. N. Lewis
1875 - 1946
 Electron
distribution is
depicted with Lewis
(electron dot)
structures
 This is how you decide
how many atoms will
bond covalently!
(In ionic bonds, it was
decided with charges)
Bond and Lone Pairs
 Valence
electrons are distributed as
shared or BOND PAIRS and unshared
or LONE PAIRS.
••
H
•
•
Cl
••
lone pair (LP)
shared or
bond pair
This is called a LEWIS
structure.
Bond Formation
A bond can result from an overlap of
atomic orbitals on neighboring atoms.
••
H
+
••
•
•
Cl
••
H
•
•
Cl
••
Overlap of H (1s) and Cl (2p)
Note that each atom has a single,
unpaired electron.
Valence Electrons
 Valence
electrons are the
electrons in the
OUTERMOST energy
level…
 They are the electrons
used for bonding
 Find them by looking at the
periodic table
Valence Electrons
Number of valence electrons of a main (A) group
atom = Group number
Steps for Building a Dot Structure
Ammonia, NH3
1. Decide on the central atom; never H. Why?
If there is a choice, the central atom is atom of lowest affinity
for electrons. (Most of the time, this is the least
electronegative atom…in advanced chemistry we use a thing
called formal charge to determine the central atom. But that’s
another story!)
Therefore, N is central on this one
2. Add up the number of valence electrons that can
be used.
H = 1 and N = 5
Total = (3 x 1) + 5
= 8 electrons / 4 pairs
Building a Dot Structure
3.Form a single bond
between the central atom
and each surrounding
atom (each bond takes 2
electrons!)
4.Remaining electrons form
LONE PAIRS to complete the
octet as needed (or duet in the
case of H).
H N H
H
••
H N H
3 BOND PAIRS and 1 LONE
PAIR.
Note that N has a share in 4 pairs (8
electrons), while H shares 1 pair.
H
Building a Dot Structure
5.
Check to make sure there
are 8 electrons around each
atom except H. H should
only have 2 electrons. This
includes SHARED pairs.
••
H N H
H
6. Also, check the number of electrons in
your drawing with the number of
electrons from step 2. If you have more
electrons in the drawing than in step 2,
you must make double or triple bonds. If
you have less electrons in the drawing
than in step 2, you made a mistake!
Carbon Dioxide, CO2
1. Central atom =
2. Valence electrons =
3. Form bonds.
C 4 eO 6 e- X 2 O’s = 12 eTotal: 16 valence electrons
This leaves 12 electrons
(6 pair).
4. Place lone pairs on outer atoms.
5. Check to see that all atoms have 8 electrons
around it except for H, which can have 2.
Carbon Dioxide, CO2
C 4 eO 6 e- X 2 O’s = 12 eTotal: 16 valence electrons
How many are in the drawing?
6. There are too many electrons in our drawing. We must form
DOUBLE BONDS between C and O. Instead of sharing only 1 pair,
a double bond shares 2 pairs. So one pair is taken away from
each atom and replaced with another bond.
Double and even
triple bonds are
commonly
observed for C, N,
P, O, and S
H2CO
SO3
C2F4
Now You Try One!
Draw Sulfur Dioxide, SO2
Violations of the Octet Rule
(Honors only)
Usually occurs with B and elements
of higher periods. Common
exceptions are: Be, B, P, S, and
Xe.
Be: 4
B: 6
P: 8 OR 10
S: 8, 10, OR 12
BF3
Xe: 8, 10, OR
12
SF4