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Transcript
Chapter 4
Elements, Atoms, and Ions
Ch#4 Contents
4.1
4.2
4.3
4.4
4.5
4.6
4.7
4.8
4.9
4.10
4.11
4.9
4.10
4.11
The Elements
Symbols for the Elements
Dalton’s Atomic Theory
Formulas of Compounds
The Structure of the Atom
Modern Concept of Atomic Structure
Isotopes
Introduction to the Periodic Table
Natural States of the Elements
Ions
Compounds that Contain Ions
Natural State of the Elements
Ions
Compounds that Contain Ions
Elements
A substance that cannot be broken down by chemical
means. Elements are defined by the number of protons
they possess.
• 115 known: 88 found in nature, others are man
made.
• Most of the know elements are found on the
Periodic Chart handing on the wall.
• One or two letters are used to represent an element
and they are called the symbol
• Some elements are solids (black), some are liquids
(blue) and some are gases (red)
Element Distribution in Earth’s Surface
•
•
•
Symbols
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 O
Krypton Kr
Sometimes the symbol is taken from the element’s
original Latin, Greek, English, or German name.

Examples:
Gold
Au aurum
Lead
Pb plumbum
Wolfram W tungsten
Symbols
Located on the front page of your textbook
Dalton’s Atomic Theory
In 1808 Dalton published A New System of Chemical
Philosophy, where he presented his theory of atoms:
•Each element is made up of tiny particles called atoms.
•The atoms of a given element are identical; the atoms of
different elements are different in some fundamental way
or ways
•Chemical compounds are formed when atoms of different
elements combine with each other. A given compound
always has the same relative numbers and types of atoms.
•Chemical changes involve reorganization of the atoms, to
different ratios.
Dalton’s Atomic Theory
Which of the following statements regarding Dalton’s
atomic theory are still believed to be true?
•Elements are made of tiny particles called atoms.
• All atoms of a given element are identical.
• A given compound always has the same relative
numbers and types of atoms.
• Atoms are indestructible.
Dalton’s Atomic Theory
Which of the following statements regarding Dalton’s
atomic theory are still believed to be true?
• Elements are made of tiny particles called atoms.
• All atoms of a given element are identical.
• A given compound always has the same relative
numbers and types of atoms.
• Atoms are indestructible.
Mendeleev’s Periodic Table
• Dmitrii I. Mendeleev
arranged elements in the
periodic table by their
chemical and physical
properties.
• He left open spaces in his
periodic table to account
for elements not yet
discovered.
The Modern Periodic Table
• The modern periodic table is also based on a
classification of elements in terms of their
physical and chemical properties.
• The horizontal rows are called periods.
• Columns contain elements of the same family or
group.
• Transition metals are the elements in group 3
through 12 in the periodic table.
Formulas
• Compound – distinct substance that is composed of
the atoms of two or more elements and always
contains exactly the same whole number ratio of
those elements.
• Chemical Formulas – expresses the types of atoms
and the number of each type in each formula unit,
or molecule of a given compound.
• Formula Examples
– H2O (molecule)
– C6H12O6 (molecule)
– NaCl (Formula Unit)
Metals
Atomic Structure
J. J. Thomson (1898—1903)
• Postulated the existence of electrons using cathoderay tubes.
• The atom must also contain positive particles that
balance exactly the negative charge carried by
particles that we now call electrons.
Cathode-ray Tube
Atomic Structure
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.
Atomic Structure
Ernest Rutherford (1911)
• Explained the nuclear atom.
• Atom has a dense center of positive charge called the
nucleus.
• Electrons travel around the nucleus at a relatively
large distance.
• A proton has the same magnitude of charge as the
electron, but its charge is positive.
Atomic Structure
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.
Atomic Structure
• 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
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.
Atomic Structure
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.
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.
Symbols of Isotopes
• Each element consists of atoms with the same
number of protons in the nucleus. This number is
called atomic number (Z).
• Protons and neutrons in atomic nuclei are called
nucleons.
• The mass number (A) is the total number
nucleons in one atom of an element.
Writing the Symbol of an Isotope
A
Z
X
• A is the mass number
• Z is the atomic number
• X is the atomic symbol
Isotopes
14
6
•
•
•
C
C = the symbol for
carbon
6 = the atomic number
(6 protons)
14 = the mass number
(6 protons and 8
neutrons)
12
6
C
• C = the symbol for carbon
• 6 = the atomic number (6
protons)
• 12 = the mass number (6
protons and 6 neutrons)
Two Isotopes of Sodium
BOHR MODELS OF HYDROGEN
(The three isotopes of hydrogen)
E
E
P
Hydrogen-1
P N
Hydrogen-2
E
NP
N
NP
N
Hydrogen-3
BOHR MODELS OF HYDROGEN
E
P
Hydrogen-1
1
H
E
E
P
N
Hydrogen-2
2
H
P
NN N
N
P
N
Hydrogen-3
3
H
BOHR MODELS OF HYDROGEN
E
P
P
1
H
2
1
H
P
NN N
N
P
N
N
Hydrogen-2
Hydrogen-1
1
E
E
Hydrogen-3
3
1
H
BOHR MODELS OF HYDROGEN
E
P
Hydrogen-1
1
1
E
E
P
N
Hydrogen-2
P
NN N
N
P
N
Hydrogen-3
2
3
1
1
Atomic Number (number of protons); found on the periodic chart
H
H
H
BOHR MODELS OF HYDROGEN
E
P
Hydrogen-1
1
1
P
Hydrogen-2
H
P
NN N
N
P
N
N
2
1
Atomic Number (number of protons)
H
E
E
Hydrogen-3
3
1
H
Mass number (sum of protons and neutrons); Not found on periodic chart
BOHR MODELS OF HYDROGEN
E
P
Hydrogen-1
1
1
E
E
P
N
Hydrogen-2
P
NN N
N
P
N
Hydrogen-3
2
3
1
1
Atomic Number (number of protons)
Mass number (sum of protons and neutrons); Not found on periodic chart
H
H
Oxidation number (Protons – electrons)
H
Relative Atomic Mass
Some atoms naturally occur as a mixture of isotopes.
e.g. 11H, 12H and 13H
The atomic weights given in the periodic table take this into
account.
They are the average atomic weight taking into account the
amounts of each isotope present.
Relative Atomic Mass of Hydrogen
When computing the average atomic mass of any
element the radioactive (unstable) isotopes are
excluded since there relative abundances are slowly
decreasing.
Hydrogen has three isotopes. The first two protium
and deuterium are stable isotopes and the third
tritium is unstable, thus excluded in the calculation.
The relative abundance of isotopes of a particular
element is constant here on our planet Earth.
Mass spectroscopy gives information of relative
abundances and relative masses of isotopes.
Relative Atomic Mass of Hydrogen
Mass spectroscopy gives the following information for
the two stable isotopes of hydrogen.
isotope name
Hydrogen-1
(protium)
Hydrogen-2
(deuterium)
relative abundance
99.985 %
0.015 %
Mass (amu)
1.007825
2.01355
To calculate a weighted average, convert the
percent to a decimal by moving the decimal two
places to the left. Then multiply this number by
the mass and add the two masses.
Relative Atomic Mass of Hydrogen
0.99985 X 1.007825 = 1.007673
0.0015 X 2.01355 = 0.003021
1.010694 amu
1.0107 amu
Average Atomic Masses
• A mass spectrometer is an instrument that
measures precise masses and relative amounts of
ions of atoms and molecule.
• The natural abundance of an isotope is its relative
proportion among all the isotopes found a natural
sample.
• The average atomic mass of an element is
calculated by multiplying the natural abundance
of each isotope by its exact mass in atomic mass
units and then summing these products.
Example
Neon is found in three isotopes in nature.
Isotope
Mass (amu)
Natural abundance (%)
Neon-20
19.9924
90.4838
Neon-21
20.99395
0.2696
Neon-22
21.9914
9.2465
19.9924 x 0.904838 =
20.99395 x 0.002696 =
21.9914 x 0.092465 =
18.0898323 amu
0.056599689 amu
2.033434801 amu
20.1797 amu
Relative Atomic Mass
If we assign a mass of 12 atomic mass units (u or sometimes
amu) to a carbon-12 atom then we can compute the relative
atomic weight for any other atom.
Atomic weights or atomic masses are given in your periodic
table.
Natural State of Elements
•
•
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
Diatomic Molecules
Allotropes
•
•
Different forms of a given element.
Example:
 Solid carbon occurs in three forms.
• Diamond
• Graphite
• Buckminsterfullerene
Carbon Allotropes
Ions
• Atoms can form ions by gaining or losing
electrons.
 Metals tend to lose one or more electrons to
form positive ions called cations and are named
by using the name of the parent atom.
 Nonmetals tend to gain electrons to form
negative ions called anions and are named by
using the root of the atom name followed by the
suffix –ide.
• Atoms become more stable by gaining or losing
electrons.
Common Oxidation States
Common Oxidation States
1+
2+
1-
Common Oxidation States
By Group Number
Sample Problem
An ion with a 3+ charge contains 23 electrons.
Which ion is it?
Sample Problem
An ion with a 3+ charge contains 23 electrons.
Which ion is it? Fe3+
A certain ion X+ contains 54 electrons and 78
neutrons. What is the mass number of this
ion?
Sample Problem
An ion with a 3+ charge contains 23 electrons.
Which ion is it? Fe3+
A certain ion X+ contains 54 electrons and 78
neutrons. What is the mass number of this
ion?
Sample Problem
An ion with a 3+ charge contains 23 electrons.
Which ion is it? Fe3+
A certain ion X+ contains 54 electrons and 78
neutrons. What is the mass number of this
ion? 133
Ionic Compounds
•
•
Ions combine to form ionic compounds.
The simplest ratio between the ions is called the
empirical formula or a formula unit.
• Properties of ionic compounds
– High melting points
– Conduct electricity
• If melted
• If dissolved in water
• Ionic compounds are electrically neutral.
• The charges on the anions and cations in the
compound must sum to zero.
Ionic Compounds
Ionic compounds do not exist as discrete molecules. Instead they
exist as crystals where ions of opposite charges occupy
positions known as lattice sites.
Ions combine in the ratio
that results in zero charge
to form ionic compounds.
Which ions are the smaller
ones?
Crystal Lattice of NaCl
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers (O.N. = P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
Sodium and nitrogen
Calcium and oxygen
Calcium and phosphorus
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
Calcium and oxygen
Na O
Na N
Ca O
Calcium and phosphorus
Ca P
Sodium and nitrogen
2-
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
2-
Calcium and oxygen
Na O
1+ 3Na N
Ca O
Calcium and phosphorus
Ca P
Sodium and nitrogen
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
2-
Calcium and oxygen
Na O
1+ 3Na N
2+ 2Ca O
Calcium and phosphorus
Ca P
Sodium and nitrogen
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
2-
Calcium and oxygen
Na O
1+ 3Na N
2+ 2Ca O
Calcium and phosphorus
Ca2+P 3-
Sodium and nitrogen
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
2-
Calcium and oxygen
Na O
1+ 3Na N
2+ 2Ca O
Calcium and phosphorus
Ca2+P 3-
Sodium and nitrogen
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
2-
Calcium and oxygen
Na O
1+ 3Na N
2+ 2Ca O
Calcium and phosphorus
Ca2+P 3-
Sodium and nitrogen
≡
Na 2O
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
2-
Calcium and oxygen
Na O
1+ 3Na N
2+ 2Ca O
Calcium and phosphorus
Ca2+P 3-
Sodium and nitrogen
≡
≡
Na 2O
Na 3N
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
2-
Na O ≡ Na 2O
1+ 3Sodium and nitrogen
Na N ≡ Na 3N
2+ 2Calcium and oxygen
Ca O ≡ Ca 2O 2 CaO
reduced
2+ 3Calcium and phosphorus Ca P
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
2-
Na O ≡ Na 2O
1+ 3Sodium and nitrogen
Na N ≡ Na 3N
2+ 2Calcium and oxygen
Ca O ≡ Ca 2O 2 CaO
reduced
2+ 3Calcium and phosphorus Ca P
Formula Writing Rules
Step 1
Write the symbols of the elements.
Step 2
Assign oxidation numbers O.N. = (P – e)
Step 3
Slide with Clyde! (number only)
Step 4
Reduce if the compound is ionic
Examples
Write formulas for the following two elements
Sodium and oxygen
1+
2-
Na O ≡ Na 2O
1+ 3Sodium and nitrogen
Na N ≡ Na 3N
2+ 2Calcium and oxygen
Ca O ≡ Ca 2O 2 CaO
reduced
2+ 3Calcium and phosphorus Ca P ≡ Ca P
3 2
The End