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Transcript
Chapter 4
Chemical Foundations:
Elements, Atoms, and Ions
Copyright © Cengage Learning. All rights reserved
1
Section 4.1
The Elements
•
•
•
118 known: 88 found in nature, others are
made in laboratories.
Abundance is the percentage found in nature.
Oxygen most abundant element (by mass) on
earth and in the human body.
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2
Section 4.2
Symbols for the Elements
•
•
•
Each element has a unique one- or two-letter symbol.
First letter is always capitalized and the second is not.
The symbol usually consists of the first one or two
letters of the element’s name.

•
Examples:
Oxygen
Krypton
O
Kr
Sometimes the symbol is taken from the element’s
original Latin or Greek name.

Examples:
Gold Au aurum
Lead Pb plumbum
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3
Section 4.3
Dalton’s Atomic Theory
Dalton’s Atomic Theory (1808)
1. Elements are made of tiny particles called
atoms.
- tiny, hard, unbreakable, spheres
2. All atoms of a given element are identical.
- all carbon atoms have the same chemical and
physical properties.
3. Atoms of a given element are different from
those of any other element.
- carbon atoms have different chemical and
physical properties than sulfur atoms.
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4
Section 4.3
Dalton’s Atomic Theory
Law of Constant Composition
4. 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.
• A given compound always has the same
composition, regardless of where it comes
from.


Water always contains 8 g of oxygen for every 1 g of
hydrogen.
Carbon dioxide always contains 2.7 g of oxygen for
every 1 g of carbon.
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5
Section 4.3
Dalton’s Atomic Theory
Dalton’s Atomic Theory (continued)
4. 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.
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6
Section 4.3
Dalton’s Atomic Theory
Dalton’s Atomic Theory (continued)
5. Atoms are indivisible in a chemical
process.
- 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 (Lead cannot be
turned into Gold by a chemical reaction)
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7
Section 4.3
Dalton’s Atomic Theory
Concept Check
Which of the following statements regarding
Dalton’s atomic theory are still believed to be
true?
I. Elements are made of tiny particles called atoms.
II. All atoms of a given element are identical.
III. A given compound always has the same relative
numbers and types of atoms.
IV. Atoms are indestructible.
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8
Section 4.4
Formulas of Compounds
Chemical Formulas Describe Compounds
•
•
Compound – distinct substance that is
composed of the atoms of two or more
elements and always contains exactly the
same relative masses of those elements.
Chemical Formulas – expresses the types of
atoms and the number of each type in each
unit (molecule) of a given compound.
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9
Section 4.4
Formulas of Compounds
Rules for Writing Formulas
1. Each atom present is represented by its element
symbol.
2. The number of each type of atom is indicated by a
subscript written to the right of the element symbol.
3. When only one atom of a given type is present, the
subscript 1 is not written.
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10
Section 4.4
Formulas of Compounds
Exercise
The pesticide known as DDT paralyzes insects by
binding to their nerve cells, leading to uncontrolled firing
of the nerves. Before most uses of DDT were banned in
the U.S., many insects had developed a resistance to it.
Write out the formula for DDT. It contains 14 carbon
atoms, 9 hydrogen atoms, and 5 atoms of chlorine.
C14H9Cl5
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11
Section 4.5
The Structure of the Atom
J. J. Thomson (1898—1903)
•
•
•
•
•
Investigated a beam called a cathode ray
He determined that the ray was made of tiny
negatively charge particles we call electrons.
His measurements led him to conclude that
these electrons were smaller than a hydrogen
atom.
If electrons are smaller than atoms, they must
be pieces of the atom. The atom must be
breakable.
He found that atoms of different elements all
produce these electrons.
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12
Section 4.5
The Structure of the Atom
The ELECTRON
- Tiny, negatively charge particle
- Very light compared to the mass of the atom. (
1/1836th the mass of a Hydrogen atom.
- Move very rapidly within the atom
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13
Section 4.5
The Structure of the Atom
William Thomson (Plum Pudding Model)
•
Reasoned that the atom might be thought of as a
uniform “pudding” of positive charge with enough
negative electrons scattered within to counterbalance
that positive charge.
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14
Section 4.5
The Structure of the Atom
Thomson’s Plum Pudding Model
• The atom IS breakable!!
• The atom has structure.
• Electrons suspended in a positively charge
electric field. – must have positive charge to
balance negative charge of electrons and make
the atom neutral.
• Mass of the atom is due to electrons.
• Atom is mostly “empty” space.
- compared size of the electrons to size of the
atom
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15
Chapter 4
Table of Contents
Rutherford’s Gold Foil Experment
• How can you prove something is empty?
• put something through it
– use large target atoms
• use very thin sheets of target so do not absorb “bullet”
– use very small particle as bullet with very high energy
• but not so small that electrons will affect it
• bullet = alpha particles, target atoms = gold foil
–  particles have a mass of 4 amu & charge of +2
c.u.
– gold has a mass of 197 amu & is very malleable
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16
Chapter 4
Table of Contents
Rutherford’s experiment on
α-particle bombardment of metal foil.
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17
Chapter 4
Table of Contents
Rutherford’s Results
• Over 98% of the 
particles went straight
through
• About 2% of the 
particles went through
but were deflected by
large angles
• About 0.01% of the 
particles bounced off
the gold foil
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18
Chapter 4
Table of Contents
(a) The results that the metal foil experiment
would have yielded if the plum pudding model
had been correct. (b) Actual results.
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19
Section 4.5
The Structure of the Atom
Ernest Rutherford (1911)
•
•
•
•
Explained the nuclear atom.
Atom has a dense center of positive charge
called the nucleus.
Electrons travel around the nucleus in the
empty space at a relatively large distance.
A proton has the same magnitude of charge as
the electron, but its charge is positive.
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20
Section 4.5
The Structure of the Atom
Rutherford and Chadwick (1932)
•
•
Most nuclei also contain a neutral particle
called the neutron.
A neutron is slightly more massive than a
proton but has no charge.
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21
Section 4.5
The Structure of the Atom
Structure of the Nucleus
• The nucleus was found to be composed of two kinds of
particles
• Some of these particles are called protons
– charge = +1
– mass is about the same as a hydrogen atom
• Since protons and electrons have the same amount of
charge, for the atom to be neutral there must be equal
numbers of protons and electrons
• The other particle is called a neutron
– has no charge
– has a mass slightly more than a proton
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22
Section 4.6
Introduction to the Modern Concept of Atomic Structure
The atom contains:
•
•
•
Electrons – found
outside the nucleus;
negatively charged
Protons – found in the
nucleus; positive charge
equal in magnitude to
the electron’s negative
charge
Neutrons – found in the
nucleus; no charge;
virtually same mass as a
proton
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23
Section 4.6
Introduction to the Modern Concept of Atomic Structure
•
The nucleus is:
 Small compared with the overall size of the
atom.
 Extremely dense; accounts for almost all of
the atom’s mass.
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24
Section 4.6
Introduction to the Modern Concept of Atomic Structure
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25
Section 4.6
Introduction to the Modern Concept of Atomic Structure
A nuclear atom viewed in cross section.
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26
Section 4.6
Introduction to the Modern Concept of Atomic Structure
Why do different atoms have different chemical
properties?
•
•
•
The chemistry of an atom arises from its
electrons.
Electrons are the parts of atoms that
“intermingle” when atoms combine to form
molecules.
It is the number of electrons that really
determines chemical behavior.
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27
Section 4.7
Isotopes
Isotopes
•
•
•
Atoms with the same number of protons but
different numbers of neutrons.
Show almost identical chemical properties;
chemistry of atom is due to its electrons.
In nature most elements contain mixtures of
isotopes.
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28
Section 4.7
Isotopes
Two Isotopes of Sodium
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29
Section 4.7
Isotopes
Isotopes
A
Z
•
•
•
X
X = the symbol of the element
A = the mass number (# of protons and
neutrons)
Z = the atomic number (# of protons)
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30
Section 4.7
Isotopes
Isotopes – An Example
14
6
•
•
•
•
•
•
12
6
C
Name of element
carbon
Number of protons
(6 protons)
Number of neutrons
(8 neutrons)
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•
•
•
•
•
•
C
Name of element
carbon
Number of protons
(6 protons)
Number of neutrons
(6 neutrons)
31
Section 4.7
Isotopes
Exercise
A certain isotope X contains 23 protons and 28
neutrons.
• What is the mass number of this isotope?
• Identify the element.
•
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Mass Number = 51 Vanadium
32
Section 4.7
Isotopes
Elements
• Arranged in a pattern called the Periodic Table
• Position on the table allows us to predict
properties of the element
• Metals
– about 75% of all the elements
– lustrous, malleable, ductile, conduct heat and
electricity
• Nonmetals
– dull, brittle, insulators
• Metalloids
– also know as semi-metals
– some properties of both metals & nonmetals
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33
Section 4.8
Introduction to the Periodic Table
The Periodic Table
•
The periodic table shows all of the known elements in
order of increasing atomic number.
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34
Section 4.8
Introduction to the Periodic Table
The Periodic Table
•
•
•
Metals vs. Nonmetals
Groups or Families – elements in the same
vertical columns; have similar chemical
properties
Periods – horizontal rows of elements
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35
Section 4.8
Introduction to the Periodic Table
The Periodic Table
•
•
•
Most elements are metals and occur on the left side.
The nonmetals appear on the right side.
Metalloids are elements that have some metallic and
some nonmetallic properties.
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36
Section 4.8
Introduction to the Periodic Table
Physical Properties of Metals
1. Efficient conduction of heat and electricity
2. Malleability (they can be hammered into thin
sheets)
3. Ductility (they can be pulled into wires)
4. A lustrous (shiny) appearance
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37
Section 4.8
Introduction to the Periodic Table
Physical Properties of Nonmetals
1. Lack properties of metals
2. Exhibit much variation in properties
3. Can be gases, liquids, or solids
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38
Section 4.8
Introduction to the Periodic Table
Physical Properties of Metalloids
1. Exhibit a mixture of metallic and non-metallic
properties
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39
Section 4.8
Introduction to the Periodic Table
•
•
Most elements are very reactive.
Elements are not generally found in
uncombined form.
 Exceptions are:
• Noble metals – gold, platinum and silver
• Noble gases – Group 8
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40
Section 4.8
Introduction to the Periodic Table
Diatomic Molecules
•
Nitrogen gas contains N2
molecules.
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• Oxygen gas contains O2
molecules.
41
Section 4.8
Introduction to the Periodic Table
Diatomic Molecules
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42
Section 4.8
Introduction to the Periodic Table
Important Group - Halogens
• Group 7A = Halogens
• very reactive
nonmetals
• react with metals to
form ionic compounds
• HX all acids
• Fluorine = F2
– pale yellow gas
• Chlorine = Cl2
– pale green gas
• Bromine = Br2
– brown liquid that has lots of
brown vapor over it
– Only other liquid element at
room conditions is the
metal Hg
• Iodine = I2
– lustrous, purple solid
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43
Section 4.8
Introduction to the Periodic Table
Allotropes
•
Many solid nonmetallic elements can exist in
different forms with different physical
properties, these are called allotopes.
• Different forms of a given element.
• Example:
 Solid carbon occurs in three forms.
• Diamond
• Graphite
• Buckminsterfullerene
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44
Section 4.8
Introduction to the Periodic Table
Carbon Allotropes
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45
Section 4.8
Introduction to the Periodic Table
•
Atoms can form ions by gaining or losing electrons.
 Metals tend to lose one or more electrons to form
positive ions called cations.

Cations are generally named by using the name of
the parent atom.
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46
Section 4.8
Introduction to the Periodic Table
•
Nonmetals tend to gain one or more electrons to form
negative ions called anions.
•
Anions are named by using the root of the atom name
followed by the suffix –ide.
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47
Section 4.8
Introduction to the Periodic Table
Ion Charges and the Periodic Table
•
The ion that a particular atom will form can
be predicted from the periodic table.
Group or Family
Alkali Metals (1A)
Alkaline Earth Metals (2A)
Halogens (7A)
Noble Gases (8A)
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Charge
1+
2+
1–
0
48
Section 4.8
Introduction to the Periodic Table
Ion Charges and the Periodic Table
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49
Section 4.8
Introduction to the Periodic Table
Electrical Nature of Matter
• Some substances
dissolve in water to form
a solution that conducts
well - these are called
electrolytes
• When dissolved in water,
electrolyte compounds
break up into component
ions
– ions are atoms or
groups of atoms that
have an electrical
charge
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50
Section 4.8
Introduction to the Periodic Table
Exercise
An ion with a 3+ charge contains 23 electrons.
Which ion is it?
a)
b)
c)
d)
Fe3+
V3+
Ca3+
Sc3+
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51
Section 4.8
Introduction to the Periodic Table
Exercise
A certain ion X+ contains 54 electrons and 78
neutrons.
What is the mass number of this ion?
•
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133
52