Download Unit 5: Atoms and the Periodic Table

Unit 5: Atoms and the
Periodic Table
Chapter 4
Part 1: Atomic Structure
Objectives
Explain Dalton’s atomic theory, and
describe why it was more successful
than Democritus’s theory.
 State the charge, mass, and location
of each part of an atom according to
the modern model of the atom.
 Compare and contrast Bohr’s model
with the modern model of the atom.

Atomic Models
Atom comes from the Greek word that
means “unable to be divided.”
 Democritus (4 b.c.) - came up with the
first theory of atomic structure; said
that the universe was made of invisible
units called atoms, but was unable to
provide evidence.
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John Dalton (1808) - used scientific
research to claim that atoms could not
be divided, that all atoms of an
element were exactly alike, and that
atoms of different elements could join
to form compounds.
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Neils Bohr (1913) - said that the
electrons in an atom orbit around the
nucleus like planets around the sun.
The path of an electron is determined
by how much energy it has, which puts
it in a specific energy level.
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Modern wave model (1925) - says that
electrons behave more like waves on a
vibrating string, and move back and
forth between energy levels. Thus, an
electron’s exact location at any given
moment cannot be determined.
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An Atom’s Contents

Atoms contain smaller pieces called
subatomic particles.
PARTICLE
CHARGE
LOCATION
MASS
proton
positive
in nucleus
1 amu
electron
negative
around
nucleus
almost 0 amu
neutron
neutral
in nucleus
1 amu
Unreacted atoms (atoms that are not
part of a chemical compound) have no
overall charge.
 That means the number of positive
charges (protons) must equal the
number of negative charges
(electrons).
 Neutrons do not affect the overall
charge.

Electrons and Energy Levels
If an electron does not have much
energy, it will be closer to the nucleus
in the first or second energy level.
 The first energy level will hold only 2
electrons.
 If an atom has more than 2 electrons,
the first two will fill the first level, and
the rest will begin filling the second
energy level.
 The second level will hold 8 electrons.

If the first and second levels are full
and there are still leftover electrons,
they will go to the third level, which
holds 18 electrons.
 Once the first three levels are full
(that’s 28 electrons!), electrons must
go to the fourth level, which holds 32
electrons.
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Orbitals

There are four different types of orbitals that
can be found within the energy levels. They
are s, p, d, and f.
 An s orbital has the lowest amount of
energy and can hold 2 electrons.
 A p orbital has more energy than s orbitals,
and there are 3 of them. Each one can hold
2 electrons, for a total of 6.
 There are 5 d orbitals and 7 f orbitals. F
orbitals have the most energy, and each one
holds 2 electrons.
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Valence Electrons
Electrons found in the outermost
energy level of an atom.
 Each atom contains between 1 and 8
valence electrons.
 These electrons are the ones that are
used to form chemical bonds with
other atoms to form molecules or
compounds.

Part 2: The Periodic Table
Objectives

Relate the organization of the periodic table
to the arrangement of electrons within an
atom.
 Explain why some atoms gain or lost
electrons to form ions.
 Determine how many protons, neutrons,
and electrons at atoms has given its symbol,
atomic number, and mass number.
 Describe how the abundance of isotopes
affects an element’s average atomic mass.
Locate alkali metals, alkaline-earth
metals, and transition metals in the
periodic table.
 Locate semiconductors, halogens, and
noble gases in the periodic table.
 Relate an element’s chemical
properties to the electron arrangement
of its atoms.

The Periodic Table

The periodic law says that if the elements
are arranged in a specific order, similarities
in their properties will occur in a regular
pattern.
 Period - horizontal row; the number of
protons (and therefore, the number of
electrons) increases by one as you move
from left to right.
 Group (or family) - vertical column; elements
in the same group have the same number of
valence electrons, so they have similar
characteristics.
The Periodic Table
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Groups
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Group 1 - Alkali metals (1 v.e.)
Group 2 - Alkaline-earth metals (2 v.e.)
Groups 3 -12 - Transition metals (number of
v.e. varies)
Group 13 - Boron family (3 v.e.)
Group 14 - Carbon family (4 v.e.)
Group 15 - Nitrogen family (5 v.e.)
Group 16 - Oxygen family (6 v.e.)
Group 17 - Halogens (7 v.e.)
Group 18 - Noble (or inert) gases (8 v.e.)
If an atom has only 1 v.e., it will be
very reactive because it will want to
stabilize itself by giving away its v.e.
The goal of an atom is to become
stable by having a totally full or totally
empty outer energy level. This will
cause it to be an ion with a +1 charge.
 Likewise, an atom with 7 v.e. will be
very reactive because it only needs 1
v.e. to be stable. Where can it find 1
v.e.?

Calculating P, N and E

Atomic number - the number of protons in
an atom (and thus, the number of electrons)
 Mass number - the number of protons plus
neutrons in an atom
 Isotope - atoms of an element that have the
same number of protons, but different
number of neutrons. This will not affect the
number of electrons, but will affect the
mass.
 Average atomic mass - the average mass of
all the isotopes of an element

To calculate the number of neutrons in
an atom:
Mass number - atomic number = neutrons

Mass is measured in atomic mass units
(amu).
 1 amu = 1/12 the mass of a standard C12 atom
Part 3: Using Moles to Count
Atoms Objectives
Explain the relationship between a
mole of a substance and Avogadro’s
constant.
 Find the molar mass of an element by
using the periodic table.
 Solve problems converting the amount
of an element in moles to its mass in
grams, and vice versa.

Using Moles to Count Atoms
We use moles to count atoms because
they are so small and so numerous. If
we know the mass of the atom, we can
estimate how many atoms are in a
sample of a substance by counting
them in groups.
 The mole has a value of 6.022 x 1023
particles in exactly 1 mole of substance.
This value is called Avogadro’s
constant.

Molar Mass

The mass in grams of 1 mole of a
substance is its molar mass, which is
the same as its average atomic mass
on the periodic table.
Using Conversion Factors
A fraction that equals 1.
 For example: 12 in/1 ft = 1;
5280 ft/ 1 mi = 1; 365.25 d/1 yr = 1


Example: If you have 5.50 mol of Fe,
and Fe has a molar mass of 55.85
g/mol, what is its mass in g?
5.50 mol Fe x 55.85 g Fe= 307 g Fe
1 mol Fe
Examples

If you have 2.50 mol of S, and S has a
molar mass of 32.07 g/mol, what is its
mass in g?
2.50 mol S x 32.07 g S = 80.18 g S
1 mol S
A) 1.80 mol Ca
 B) 0.50 mol C
 C) 3.20 mol Cu

A) 72.14 g Ca
 B) 6.01 g C
 C) 203.36 g Cu


How many moles are in 620 g of Hg?
620 g Hg x 1 mol Hg = 3.09 mol Hg
200.59 g Hg
A) 352 g Fe
 B) 11 g Si
 C) 205 g He

A) 6.30 mol He
 B) 0.39 mol Si
 C) 51.25 mol He
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