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CHAPTER 2

Atoms, Molecules
and Ions
1
Atoms
Each element composed of atoms
 All atoms of a given element are identical
 Atoms of an element are not changed during
a chemical reaction (the atoms just move
from one chemical species to another)
 Compounds are formed when atoms of more
than 1 element combine

2
Atoms
 John
Dalton
• The Atomic Theory of Matter
• Credited with developing Natural Laws
3
Atoms
 Law
of conservation of mass
 Law of constant composition
 Law
of multiple proportions: When two
elements form different compounds,
the mass ratio of the elements in one
compound is related to the mass ratio
in the other by a small whole number.
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Discovery and Properties of
Electrons

Humphrey Davy (early 1800’s) - passed
electricity through compounds
• compounds decomposed into elements
• compounds are held together by electrical forces

Michael Faraday - (1832-1833) - amount of
reaction that occurs during electrolysis is
proportional to current passed through
compounds
• Matter (atoms) is electrical in nature.
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Discovery and Properties of
Electrons

Cathode Ray Tubes - (late 1800’s & early
1900’s)
• 2 electrodes in a glass tube with a gas at low
pressure
• voltage applied to tube causing a glow discharge
• “rays” emitted from cathode (- end) to anode (+
end)

Cathode Rays must be negatively charged!
6
Discovery and Properties of
Electrons
 J.J.
Thomson - (1897) - changed
cathode ray tube experiments by adding
two adjustable voltage electrodes into
the experiment
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Discovery and Properties of
Electrons
8
Discovery and Properties of
Electrons
 measured
charge to mass ratio of
electrons
• e/m = -1.75881 x 108 coulomb/g of e named
cathode rays electrons
 Thomson is the “discoverer of electrons”
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Discovery and Properties of
Electrons
 Robert
A. Millikan - 1st American Nobel
Laureate
• determined the charge and mass of the
electron (1909)
• oil drop experiment
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Discovery and Properties of
Electrons
on 1 electron = -1.60219 x 10-19
coulomb
 using Thomson’s charge to mass ratio we
get that the mass of 1 electron is
9.11 x 10-28 g
 charge
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Canal Rays and Protons
 Goldstein
(1886) - “Canal Rays”
• streams of positively charged particles in
cathode rays
• flow in opposite direction of cathode rays
 must
be positive
• postulated existence of “proton”
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Radioactivity
Spontaneous emission of high energy radiation
A radioactive substance is placed in a shield
containing a small hole so that a beam of radiation
is emitted from the hole.
The radiation is passed between two electrically
charged plates and detected.
Three spots are noted on the detector:
•a spot in the direction of the positive plate,
•a spot which is not affected by the electric field,
•a spot in the direction of the negative plate.
13
Radioactivity
14
Radioactivity
A high deflection towards the positive plate
corresponds to radiation which is negatively
charged and of low mass. This is called bradiation (consists of electrons).
No deflection corresponds to neutral radiation.
This is called g-radiation.
Small deflection towards the negatively charged
plate corresponds to high mass, positively charged
radiation. This is called a-radiation.
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Discovery of the Nuclear Atom
Plum-Pudding?
Thomson assumed all
these charged species
were found in a sphere.
16
Rutherford and the Nuclear Atom
 Ernest
Rutherford - 1910 - basic picture
of atom
Geiger & Marsden’s experiment on a- particle
scattering from thin Au foils
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Rutherford and the Nuclear Atom
In order to get the majority of a-particles
through a piece of foil to be undeflected, the
majority of the atom must consist of a low mass,
diffuse negative charge - the electron.
To account for the small number of high
deflections of the a-particles, the center or
nucleus of the atom must consist of a dense
positive charge.
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Rutherford and the Nuclear Atom
 Rutherford
decoded the scattering
information
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Rutherford and the Nuclear Atom
 atom
is mostly empty space
 very small, dense center called nucleus
 nearly all of atom’s mass in nucleus
 nuclear diameter is 1/10,000 to 1/100,000
times less than atom’s radius
 nuclear density is 1015g/mL
 equivalent to 3.72 x 109 tons/in3
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Neutrons
 James
Chadwick - 1932
 analyzed evidence from a-particle
scattering off Be
 recognized existence of massive
neutral particles - “neutrons”
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The Modern View of Atomic Structure
The atom consists of positive, negative, and
neutral entities (protons, electrons, and neutrons).
Protons and neutrons are located in the nucleus
of the atom, which is small. Most of the mass of
the atom is due to the nucleus.
There can be a variable number of neutrons for the
same number of protons. Isotopes have the same
number of protons but different numbers of neutrons.
Electrons are located outside of the nucleus.
Most of the volume of the atom is due to electrons.
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The Modern View of Atomic Structure
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Fundamental Particles

Three fundamental particles make up atoms. The
following table lists these particles together with
their masses and their charges.
Particle
Mass (amu)
Charge
-
0.00054858
-1
+
1.0073
+1
0
1.0087
0
Electron (e )
Proton (p,p )
Neutron(n,n )
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Mass Number & Isotopes
 H.G.
J. Moseley (1912-1914) - recognized
that atomic number is the defining
difference between elements
 new understanding of Mendeleev’s
periodic law
25
Atomic Number
 Sometimes
given the symbol Z
 number of protons in the nucleus
 determines the element
• also determines number of electrons in a
neutral atom
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Isotopes, Atomic Numbers & Mass
Numbers
 All
atoms of an element have the same
number of protons in the nucleus
 Isotopes of an atom have a different
number of neutrons in the nucleus
 Atomic number = # of protons
 Mass number = # protons + # neutrons
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Isotopes, Atomic Numbers & Mass
Numbers
By convention, for element X, we write
A
Z
X
Isotopes have the same Z but different A.
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Isotopes, Atomic Numbers & Mass
Numbers
29
Isotopes
 Give
the number of protons, neutrons
and electrons in each of the following
species:
56Fe
56Fe3+
31P
31P3-
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Atomic Weights
weighted average of the masses of the constituent
isotopes
 lower number on periodic chart

• How do we know what the values of these numbers are?
31
The Periodic Table
The Periodic Table is used to organize the 114
elements in a meaningful way.
As a consequence of this organization, there are
periodic properties associated with the periodic
table.
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The Periodic Table
Columns in the periodic table are called groups
(numbered from 1A to 8A or 1 to 18).
Rows in the periodic table are called periods.
.
33
The Periodic Table
Metals are located on the left hand side of the
periodic table (most of the elements are metals).
Non-metals are located in the top right hand side
of the periodic table.
Elements with properties similar to both metals
and non-metals are called metalloids and are
located at the interface between the metals and
non-metals.
34
The Periodic Table
35
The Periodic Table
Some of the groups in the periodic table are
given special names.
These names indicate the similarities between
group members:
Group 1A: Alkali metals.
Group 2A: Alkaline earth metals.
Group 6A: Chalcogens.
Group 7A: Halogens.
Group 8A: Noble gases.
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The Periodic Table
 Name
the following elements. Indicate if
each is a metal, nonmetal, or metalloid.
P
Sn
Mn
K
Cu
Hg
F
As
N
 Si
Na
Ca
 Fe
Ag
Mg
37
Molecules and Molecular Compounds
Molecules are assemblies of two or more atoms
bonded together.
Each molecule has a chemical formula.
The chemical formula indicates
•which atoms are found in the molecule
•in what proportion they are found.
Compounds formed from molecules are
molecular compounds.
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Molecules and Molecular Compounds
empirical formula - simplest molecular formula,
shows ratios of elements but not actual numbers
of elements
 molecular formula - actual numbers of atoms of
each element in the compound

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Molecules and Molecular Compounds
Molecules occupy three dimensional space.
However, we often represent them in two
dimensions.
The structural formula gives the connectivity
between individual atoms in the molecule.
The structural formula may or may not be used
to show the three dimensional shape of the
molecule.
40
Molecules and Molecular Compounds
41
Molecules and Molecular Compounds
If the structural formula does show the
shape of the molecule, then either a
perspective drawing, ball-and-stick model,
or space-filling model is used.
42
Chemical Formulas
show the ratio of the elements present in the
molecule or compound
 He, Au, Na - monatomic
 O2, H2, Cl2 - diatomic
 O3, S4, P8 - more complex elements
 H2O, C12H22O11 - compounds

43
Ions & Ionic Compounds
 ions
are atoms or groups of atoms that are
charged
 two basic types of ions
• positive ions or cations
– one or more electrons less than neutral
• negative ions or anions
– one or more electrons more than neutral
44
Ions and Ionic Compounds
The number of electrons an atom loses is
related to its position on the periodic table.
Metals tend to form cations whereas non-metals
tend to form anions.
45
Ions and Ionic Compounds
The majority of chemistry involves the transfer
of electrons between species. (Ionic Bonding)
Example:
To form NaCl, the neutral sodium atom, Na, must lose
an electron to become a cation: Na+.
The electron cannot be lost entirely, so it is transferred
to a chlorine atom, Cl, which then becomes an anion:
Cl-.
The Na+ and Cl- ions are attracted to form an ionic
NaCl lattice which crystallizes.
46
Ions & Ionic Compounds
 Sodium
chloride - table salt is an ionic
compound
47
Ion Names and Formulas
 Common
Polyatomic Ions
 Table 2.4 & 2.5
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Ion Names and Formulas
49
Ion Names and Formulas
50
Nomenclature
Naming of compounds, nomenclature, is
divided into organic compounds (those
containing C) and inorganic compounds (the rest
of the periodic table).
51
Nomenclature
Cations formed from a metal have the same
name as the metal.
Example: Na+ = sodium ion.
If the metal can form more than one cation,
then the charge is indicated in parentheses in the
name.
Examples: Cu+ = copper(I); Cu2+ = copper(II).
Cations formed from non-metals end in -ium.
Example: NH4+ ammonium ion.
52
Nomenclature
Monatomic anions (with only one atom) are
called
-ide.
Example: Cl- is chloride ion.
Exceptions: hydroxide (OH-), cyanide (CN-),
peroxide (O22-).
Polyatomic anions (with many atoms)
containing oxygen end in -ate or -ite. (The one
with more oxygen is called -ate.)
Examples: NO3- is nitrate, NO2- is nitrite.
53
Nomenclature
cations + anions must give neutral charge
 NaCl
sodium chloride
(+1 & -1)
 KOH
potassium hydroxide (+1 & -1)
 CaSO4
calcium sulfate
(+2 & -2)
 Al(OH)3 aluminum hydroxide (+3 & -3)

54
Nomenclature
Consider the formation of Mg3N2:
Mg loses two electrons to become Mg2+
Nitrogen gains three electrons to become N3-.
For a neutral species, the number of electrons
lost and gained must be equal.
However, Mg can only lose electrons in twos and
N can only accept electrons in threes.
55
Nomenclature
Therefore, Mg needs to lose 6 electrons (2
N gain those 6 electrons (3  2).
 3) and
 3Mg atoms need to form 3Mg2+ ions (total 3x2+
charges)
 2 N atoms need to form 2N3- ions (total 2x3charges).
Therefore, the formula is Mg3N2.
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Nomenclature
 Predict
the chemical formulas when the
following species combine:
 Na with O
 Ca with Cl
 Fe(III) with Cl
 Ca with O
 Cu(II) with OH57
Nomenclature
 name
of K2SO3
 formula
of ammonium sulfide
 formula
of aluminim sulfate
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Nomenclature
Polyatomic anions containing oxygen with
additional hydrogens are named by adding
hydrogen or bi- (one H), dihydrogen (two H), etc.,
to the name as follows:
CO32- is the carbonate anion
HCO3- is the hydrogen carbonate (or bicarbonate)
anion.
H2PO4- is the dihydrogen phosphate anion.
59
Nomenclature
 Polyatomic anions containing oxygen with
more than two members in the series are named
as follows (in order of decreasing oxygen):
per-….-ate
-ate
-ite
hypo-….-ite
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Nomenclature
61
Naming Molecular Compounds
Binary molecular compounds have two
elements.
The most metallic element is usually written
first (i.e., the one to the farthest left on the
periodic table). Exception: NH3.
If both elements are in the same group, the
lower one is written first.
Greek prefixes are used to indicate the number
of atoms.
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Naming Molecular Compounds
63
Naming Molecular Compounds
 Name
the following molecular
compounds
 CO2
 PF5
 Cl2O7
 H 2S
 SiH4
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Naming Acids
The names of acids (molecules that lose H+) are
related to the names of anions:
Molecular acids
-ide becomes hydro-…..-ic acid
If acid has oxygen (oxoacids)
-ate becomes -ic acid;
-ite becomes -ous acid.
(keep the “per” and “hypo” if necessary)
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Naming Acids
66
Name the following Acids
 HCl
 H2SO4
 HNO3
 H3PO4
 Hydrobrombic
acid
 Perchloric acid
 Hypochlorous acid
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