Transcript Elements

Chapter 4
Atoms and
Elements
Homework
 Assigned
Problems (odd numbers only)
 Recommended:
Exercises, 1-27
 Required:
Problems, 29-101
 Required:
Cumulative Problems, 103-117
 Optional:
Highlight Problems, 119
Indivisible: The Atomic Theory
Matter is anything with a mass and occupies space
 Matter (in our world) is composed of combinations
of about 100 basic substances called elements
 109 elements have been discovered and isolated
 88 are found in nature
 21 are (synthetic) man-made
 Oxygen most abundant element (by mass) on earth

The Atomic Theory

An element is a pure substance that cannot be broken
down into simpler substances by a chemical means
 Single atom of that element
 Sample of the element large enough to weigh on a
scale
 Generally referring to the presence of that element
(compound), not necessarily in its free form
The Atomic Theory
 An
atom is the smallest particle of
an element that can exist and still
have properties of that element
 All atoms of a certain type are
similar to one another and different
from all other types
 109 different types are known and
each “type” is a different element
The Atomic Theory (1808)

Dalton proposed a set of five
statements that summarize the
modern scientific concept about
atoms:
1)
All matter is made from small particles
called atoms (109 different types)
All atoms of a given type are similar to
one another and significantly different
from all other types
2)
`
Dalton’s Atomic Theory
The number and arrangement of
different types of atoms in a pure
substance (a compound) determine
its identity
4) A chemical change is a combination,
separation, or rearrangement of
atoms (forms new substances)
5) Only whole atoms take part or result
from any chemical reaction
3)
`
Dalton’s Atomic Theory
 Atoms
are indivisible in a chemical
process (indestructible)
All atoms present at beginning are
present at the end
Atoms are not created or
destroyed, just rearranged
Atoms of one element cannot
change into atoms of another
element
Cannot
turn Lead into Gold by a
chemical reaction
Thomson’s Model of the Atom

J.J. Thomson (1897) used a gas discharge tube to
examine a glowing stream of light called a cathode ray
 Determined that the ray was made of tiny negatively
charged particles we call electrons
 He determined the electrons were smaller than a
hydrogen atom
 Since electrons are smaller than atoms they must be
parts of an atom
 Atoms must be divisible
 Atoms of different elements all produced these same
electrons
Thomsons’s Model of the Atom
 Defined
by Thomson
 Tiny, negatively charged particle
 Electrical charge = 1 ¯
 Very light compared to mass of atom
 1/2000th the mass of a H atom
 Moves very rapidly within the atom
Thomson’s Model of the Atom
Thomson’s experiment showed that atoms
have a structure and are divisible
 Thomson reasoned that electrons must be a
fraction of the entire size of the atom since
their mass is much smaller that the whole
atom
 Thomson also reasoned since atoms are
neutral, the electrons were embedded in a
sphere of uniform positive charge
 Thomson (1898) proposed the “Plum
Pudding” model or “Raisin Muffin” model of
the atom

Thomson’s Model of the Atom

Thomson Atomic
Model (early 1900’s):
Proposed a uniform,
positive sphere of
matter with small
negative electrons
attached to the
surface of the sphere
 This became known
as the plum-pudding
model
The Nuclear Atom:
Rutherford’s Experiment
 1911
Rutherford designed an experiment
to test the Thomson model (“plumpudding”) of the atom
 Rutherford
directed positively charged
particles (alpha particles) towards a thin
gold foil sheet
 Rutherford
expected the particles to pass
straight through a uniform area of mass
and positive charge
The Nuclear Atom:
Rutherford’s Experiment
Some 
particles are
scattered
Most particles
pass straight
through foil
Source of
 particles
Beam of
 particles
Fig4_5
Screen to detect
scattered  particles
Thin
metal foil
The Nuclear Atom:
Rutherford’s Experiment
Results:
 Most (alpha) particles
mostly went straight
through
 A few particles were
unexpectedly deflected
from their expected
(straight) path
 A few deflected nearly
back towards alpha
particle source
The Nuclear Atom
Rutherford proposed:
 A very small, dense core at the center of
the atom
 This dense core was called the “nucleus”
 It contains most of the mass of the atom
and it has a positive charge (protons)
 Most of an atom is empty space filled with
electrons
The Nuclear Atom (Rutherford Model)

The nucleus is the
center (core) of the
atom
Extranuclear
 The nucleus
region
has most of the
mass of the atom
 protons
 neutrons

The extranuclear region
 It contains all the
electrons
 Electrons occupy
most of the total
volume of the atom
nucleus
The Nuclear Atom (Rutherford Model)
Rutherford’s Nuclear Theory of the Atom

A very dense, small center exists in the
center of the atom called the nucleus

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Volume of nucleus is about 1/10 trillionth the volume of
the entire atom
Nucleus is basically the entire mass of the
atom

The protons and neutrons are located in the nucleus

Most of the atom is empty space with fast-moving
electrons
The Nuclear Atom (Rutherford Model)
The nucleus
The core of the atom
 Positively
charged
 Contains most of the mass of the
atom
 Within a neutral atom, there are
equal numbers of protons and
electrons, so atom has a net charge
of zero
The Properties of Protons, Neutrons, and
Electrons

Experimentation in the early 20th century (Thomson
and Rutherford) proved atoms were not indivisible
spheres

Atoms are comprised of smaller particles:
Subatomic particles

More experiments led to the discovery of two more
fundamental subatomic particles: Protons and
neutrons
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Electron: Negatively charged (1897)
Proton: Positively charged (1919)
Neutron: No electrical charge (1932)
The Properties of Protons, Neutrons, and Electrons:
Mass of Subatomic Particles

The usual standards of mass
such as the gram or pound
are not practical for use with
atoms
 Chemists base the mass of
atoms on the atomic mass
scale
 A relative scale based on the
mass of one carbon atom:
12.00 amu (atomic mass unit)
 One amu is 1/12 the mass of
one carbon atom, so the
approximate mass of one
proton or neutron is 1.00 amu
The Properties of Protons, Neutrons and Electrons:
The Electron
 The
electron is a subatomic particle that
has one unit of negative charge (-1)
 Its relative mass is about 1/2000 times a
proton
 It is found in the region outside of the
nucleus
The Properties of Protons, Neutrons and Electrons:
The Proton
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The proton has the same magnitude charge as the
electron, but oppositely charged
The proton has one unit of positive charge (+1)
Its relative mass is about 2000 times an electron
It is located in the nucleus
In a neutral atom:
Number of protons = number of electrons

An element is a pure substance in which all atoms
have the same atomic number
 No. of protons is a characteristic of an element
The Properties of Protons, Neutrons, and Electrons:
The Neutron
The neutron is the last of the three subatomic
particles to be discovered
 The neutron is neutral: It has no unit charge
 Its relative mass is also about 2000 times an
electron
 It is located in the nucleus
 Variable amounts are possible in atoms of the
same element: This is the basis for isotopes

Elements:
Defined by Their Numbers of Protons

The atomic number is equal to the number
of protons in the nucleus of an atom
Atomic Number = number of protons in an atom
 Determines
the identity of the atom
 It is also equal to the number of electrons in the
neutral atom
 The top number in each square in the periodic
table
 All elements in periodic table arranged
according to the atomic number
Elements:
Chemical Symbol


Each element has a unique chemical symbol
One or two letter abbreviations
 If two letters, the second is lower case
 The letter symbol often corresponds to the name of
the element
 F = Fluorine
 P = Phosphorous
 Some symbols derived from the Latin or Greek
names
 Lead – Pb (plumbum)
 Gold – Au (aurum)
 Sodium – Na (natrium)
Elements:
Names of Symbols and Some Common Elements
 Required:
Know the name and symbol
of some of the most common elements
(from list given)
 See handout given in class
 A periodic table will be given on each
test or quiz
 Required: Know how to use a periodic
table to find needed information
The Periodic Law and the Periodic Table
In early part of 19th century many chemical
facts were being obtained for the elements
known (at that time)
 As the number of known elements
increased, scientists began to attempt useful
classifications
 Scientists had discovered that certain
groups of elements had similar chemical
properties
 Many attempts were made to arrange or
classify the elements in order to explain the
similarities

The Periodic Law and the Periodic Table

In 1869, Mendeleev attempted to arrange the elements by
their relative masses

Mendeleev (Meyer) noted that similar chemical properties
recur periodically when the elements are arranged by
increasing atomic mass
Mendeleev then arranged the elements with similar
chemical properties in the same vertical columns
The periodic law (Mendeleev): When elements are
arranged in order of increasing atomic mass (number),
elements with similar properties occur at periodic intervals
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The Periodic Law and the Periodic Table
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
The periodic table is a chart of the elements with similar chemical
properties arranged into vertical columns (groups)
Horizontal rows are called periods
The periodic table has two categories of elements:
 Main group elements are those in the columns labeled with
numbers (1A-8A)
 Transition elements are those in the columns labeled with the letter
“B” (1B-8B)
Group number
The Periodic Law and the Periodic Table



Each box in the periodic table designates an element, with that element’s symbol in the
center of the box
The number above the symbol is the element’s atomic number
The elements are arranged in the order of increasing atomic number (across the period)
The Periodic Table
1A
1
2
3
4
5
6
7
2A
8A
3A 4A 5A 6A7A
The Periodic Law and the Periodic Table
nonmetalss
metals

Metals
 Everything to the left of the metal/nonmetal barrier
 Shiny solid, good conductor of electricity, ductile and
malleable
 Nonmetals
 Everything to the right of the metal/nonmetal barrier
 Dull appearance, not ductile or malleable not good
conductors of electricity
The Periodic Law and the Periodic Table
Metalloids
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Elements with properties intermediate between those metals and
nonmetals
On the metal/nonmetal barrier
Have some physical properties of metals but some chemical properties
of nonmetals
Semiconductors
 Si, Ge, As, Sb, Te
The Periodic Law and the Periodic Table
Families of Elements
Certain groups of elements have their own
special names due to the chemical similarity
of the elements in them
 Alkali Metals
 Group 1A
 Alkaline Earth Metals
 Group 2A
 Halogens
 Group 7A
 Noble Gases
 Group 8A
 Become familiar with these group names
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The Periodic Law and the Periodic Table:
Main Group Elements
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Also called representative elements
The elements in the A-groups
 First two columns (1A and 2A)
 The last 6 columns (3A to 8A)
Easy to predict ionic structure
Some groups have names
The Periodic Law and the Periodic Table:
Transition Elements
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

The elements in the B-groups
Middle block of elements (3B through 2B)
Includes the two groups at the bottom
 Lanthanides and Actinides
 Difficult to predict ionic structure
Ions: Losing and Gaining Electrons
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
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Electron transfer may occur between two or more
atoms
It produces charged particles called ions
An ion is an atom that is electrically charged from
the loss or gain of electrons
Atoms are neutral only when the number of
protons equals the number of electrons
When atoms are lost or gained, this protonelectron balance (neutrality) falls apart
This leaves a net charge on the atom
Ion charge = number of protons - number of electrons
Ions: Losing and Gaining Electrons
Cations

Positive ions form when an electron or electrons are
lost from a metal
 If the number of protons is the same and the number
of electrons decreases, excess positive charge results
 The number of protons never changes when ions form
Sodium
Magnesium
Aluminum
Calcium


Sodium Ion
Na  Na  e
2

Mg  Mg  2e Magnesium Ion
 Aluminum Ion
 3e
2
 Calcium Ion
3
Al  Al
Ca  Ca
 2e
Ions: Losing and Gaining Electrons
Anions

Negative ions form when an electron or electrons are
gained from a nonmetal
 If the number of protons is the same and the number
of electrons increases, excess negative charge results
 The number of protons never changes when ions form
Fluorine
Bromine
Oxygen
Sulfur


Fe F


Br  e  Br

2
O  2e  O

2
S  2e  S
Fluoride
Bromide
Oxide
Sulfide
Ions and the Periodic Table
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Valence electrons are important in determining the
bonding characteristic of an atom
The outermost electrons in any atom (furthest from
the nucleus)
The maximum number of valence electrons for
any element is eight which is the most stable
valence electron configuration
Only noble gases have the maximum number
5 of the 6 noble gases have eight valence
electrons
Helium (the exception) has only two valence
electrons
Ions and the Periodic Table

The number above each main-group gives the number of
valence electrons for element in that group
 The charge of an ion is directly correlated with the number
of electrons lost or gained when compounds are formed
 Atoms of elements tend to lose, gain, or share electrons as
to produce a noble gas electron configuration for each of
atoms involved
 The charge on the ion is directly related to the number of
electrons lost or gained
Ions and the Periodic Table
Positive Ions
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Group IA, IIA, and IIIA metal
atoms contain one, two, or
three more electrons
(respectively) than the
preceding noble gas
Metal atoms (IA, IIA, IIIA)
easily lose electrons to acquire
the noble gas electron
configuration
Produces a (+) charged atom:
cation
The charge on the cation is
directly correlated with the
number of electrons lost
Ions and the Periodic Table
Negative Ions
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Group VA, VIA, and VIIA
nonmetal atoms contain three,
four, or five less electrons
(respectively) than the nearest
noble gas
Nonmetal atoms (VA, VIA,
VIIA) will gain the necessary
number of electrons to acquire
the noble gas electron
configuration
produces a (-) charged atom:
anion
The charge on the anion is
directly related to the number
of electrons gained
Isotopes: When the Number of Neutrons Varies

All atoms of the same element have the same
atomic number (Z)

The same element can differ in the mass number
(A) due to a different number of neutrons

All Mg atoms have 12 protons, but may have 12,
13, or 14 neutrons
Isotopes: When the Number of Neutrons Varies
 Atoms
that have the same number of protons
and electrons but different numbers of neutrons
are called isotopes
 Since
isotopes are atoms of the same element,
 They have the same atomic number
 They display the same chemical properties
 All elements have their own exclusive percent
natural abundances of the isotopes
Isotopes: Mass Number (A)

The total number of protons and neutrons in
an atom
Mass Number = number of protons + number of neutrons

Mass number is always a whole number
(no decimals)

An oxygen atom has a mass number of 16
(8 protons and 8 neutrons)
Nuclear (Isotopic) Symbols
A
notation used when necessary to
differentiate between isotopes
A
X
Z
A is the mass number
Z is the atomic number
X is the chemical symbol
Atomic Mass
A
specific element can have several
mass values if it exists in isotopic forms
 For
example, oxygen atoms can have
any one of three masses but often
treated as if it has one mass
 The
atomic mass of an element is the
mass of the “average atom” of that
element
Atomic Mass

1)
2)
3)
Atomic mass is a
“weighted average
mass” based on:
The number of
isotopes that exist
for the element
The relative mass of
each isotope
The percent
abundance of each
isotope
Example:
Isotopes and Atomic Mass
Complete
Symbol
Name
1
Hydrogen
H
1
19
9 F Fluorine
64
Copper
Cu
29
2
Hydrogen
H
1
the following table:
# Protons #Neutrons
#Electrons
1
0
1
9
10
9
29
35
29
1
1
1
end