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Chapter 2
ATOMS, MOLECULES,
AND IONS
2.1 THE ATOMIC THEORY OF MATTER
Material world made up of tiny indivisible particles they
called atomos, meaning “indivisible or uncuttable.”
Dalton’s Atomic Theory
Dalton’s atomic theory was based on the four postulates
1. Each element is composed of extremely small particles
called atoms.
An atom of the element oxygen
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An atom of the element nitrogen
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2. All atoms of a given element are identical, but the
atoms of one element are different from the atoms of all
other elements.
3. Atoms of one element cannot be changed into atoms of
a different element by chemical reactions; atoms are
neither created nor destroyed in chemical reactions.
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4. Compounds are formed when atoms of more than one
element combine; a given compound always has the
same relative number and kind of atoms.
Dalton’s theory explains several laws of chemical
combination that were known during his time
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Law of conservation of mass
The total mass of materials present after a chemical
reaction is the same as the total mass present before the
reaction.
based on postulate 3
Law of constant composition
In a given compound, the relative numbers and kinds of
atoms are constant.
based on postulate 4
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A good theory explains known facts
and predicts new ones
Dalton used his theory to deduce the
law of multiple proportions:
If two elements A and B combine to form more than one
compound, the masses of B that can combine with a given
mass of A are in the ratio of small whole numbers.
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considering water and hydrogen peroxide
8.0 g Oxygen
water
1.0 g Hydrogen
16.0 g Oxygen
hydrogen peroxide
1.0 g Hydrogen
Thus, the ratio of the mass of oxygen per gram of
hydrogen in the two compounds is 2:1.
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Using Dalton’s atomic theory, we conclude that hydrogen
peroxide contains twice as many atoms of oxygen per
hydrogen atom as does water.
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
Today we know that the atom is composed of subatomic
particles.
fact: Particles with the same charge repel one another,
whereas particles with unlike charges attract one another.
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Cathode Rays and Electrons
 Electrons move from the cathode (negative electrode) to
the anode (positive electrode).
 The tube contains a glass screen (set diagonally to the
electron beam) that fluoresces, showing the path of the
cathode rays.
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The rays are deflected by a magnet.
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 J. J. Thomson observed several many properties of the
cathod rays.
 They are the same regardless of the identity of the
cathode material
 Cathode rays are streams of negatively charged
particles ´electrons´
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Thomson constructed a cathode-ray tube
From these measurements
1.76×108 C/g of electron
Charge-to-mass ratio
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R. Millikans oil-drop experiment
The value of Charge = 1.602×108 C
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The value of Charge of electron = 1.602×108 C
Thomsons Charge-to-mass ratio = 1.76×108 C/g of electron
1.602  1019 C
28
Electron mass 

9.10

10
g
8
1.76  10 C/g
Radioactivity
spontaneous emission of radiation.
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Rutherford
three types of radiation:
alpha (), beta (), and gamma ().
The paths of () and () radiation are bent by an electric field,
although in opposite directions; () radiation is unaffected by
the field
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Gamma radiation is high-energy radiation similar to X-rays; it
does not consist of particles and carries no charge.
The Nuclear Model of the Atom
Rutherford
He proposed that the atom
consisted of a uniform positive
sphere of matter in which the
electrons were embedded like
raisins in a pudding or seeds in
a watermelon
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Rutherford’s -scattering experiment.
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Rutherford explained the
results by postulating the
nuclear model of the atom,
 most of the mass of each gold atom and all of its positive
charge reside in a very small, extremely dense region that
he called the nucleus.
 most of the volume of an atom is empty space in which
electrons move around the nucleus.
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Subsequent experiments led to the discovery of
positive particles (protons)
and neutral particles (neutrons) in the nucleus.
Protons were discovered in 1919 by Rutherford
and neutrons in 1932 by British scientist James Chadwick
(1891–1972).
Thus, the atom is composed of electrons,
protons, and neutrons.
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2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
The charge of an electron is 1.602 × 1019 .
That of a proton is equal in magnitude, +1.602 × 10-19 C.
The quantity 1.602 × 1019 C is called the electronic charge
Every atom has an equal number of electrons and
protons, so atoms have no net electrical charge.
Atoms have extremely small masses.
The mass of the heaviest known atom, is approximately
4 × 1022 g
we use the atomic mass unit (amu)
1 amu = 1.66054 × 1024 g
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Because it takes 1836 electrons to equal the mass of one
proton or one neutron, the nucleus contains most of the
mass of an atom.
Most atoms have diameters between 1 × 1010 m to 5 × 1010m
1 Å = 1  1010 m
The diameter of a chlorine atom, for example, is 200 pm, or
0.20 Å
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Atomic Numbers, Mass Numbers, and Isotopes
The atoms of each element have a characteristic
number of protons.
the number of protons in an atom of any particular
element is called that element’s atomic number
the mass number, is the number of protons plus neutrons
in the atom
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Atoms with identical atomic numbers but different mass
numbers (that is, same number of protons but different
numbers of neutrons) are called isotopes of one another.
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a) The superscript 197 is the mass number (protons + neutrons)
gold has atomic number 79
197Au
has 79 protons, 79 electrons, and 197 - 79 = 118
neutrons
b) The atomic number of strontium (listed on inside front
cover) is 38. Thus, all atoms of this element have 38 protons
and 38 electrons. The strontium-90 isotope has 90 - 38 = 52
neutrons
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(a) Magnesium has atomic number 12, so all atoms of
magnesium contain 12 protons and 12 electrons.
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12
Mg
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12
Mg
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12
Mg
The numbers of neutrons in an atom of each isotope are
therefore 12, 13, and 14, respectively
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2.4 ATOMIC WEIGHTS
The Atomic Mass Scale
The atomic mass unit is presently defined by assigning a
mass of exactly 12 amu to an atom of the 12C isotope of
carbon.
In these units, an 1H atom has a mass of 1.0078 amu and
an 16O atom has a mass of 15.9949 amu.
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Atomic Weight
Most elements occur in nature as mixtures of isotopes.
We can determine the average atomic mass of an
element, usually called the element’s atomic weight, by
using the masses of its isotopes and their relative
abundances:
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2.5 THE PERIODIC TABLE
patterns in chemical behavior
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2.6 MOLECULES AND MOLECULAR COMPOUNDS
Only the noble-gas elements are normally found in
nature as isolated atoms.
Most matter is composed of molecules or ions.
Molecules and Chemical Formulas
Several elements are found in nature in molecular form—two or
more of the same type of atom bound together.
chemical formula
O2
A molecule made up of two atoms is
called a diatomic molecule.
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molecular compounds: Compounds composed of molecules
contain more than one type of atom.
Most molecular substances we will
encounter contain only nonmetals.
Molecular and Empirical Formulas
Molecular Formulas
Chemical formulas that indicate the actual numbers of atoms in
a molecule.
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Empirical Formulas
Chemical formulas that give only the relative number of
atoms of each type in a molecule.
 The subscripts in an empirical formula are always the
smallest possible whole-number ratios.
 For many substances, the molecular formula and the
empirical formula are identical, as in the case of water, H2O.
 Whenever we know the molecular formula of a
compound, we can determine its empirical formula.
The converse is not true, however.
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empirical formula for glucose is CH2O.
the empirical formula for nitrous oxide is the same as its
molecular formula, N2O.
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Picturing Molecules
A structural formula
shows which atoms are attached to which
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2.7 IONS AND IONIC COMPOUNDS
If electrons are removed from or added to an atom, a charged
particle called an ion is formed. An ion with a positive charge
is a cation (pronounced CAT-ion); a negatively charged ion is
an anion (AN-ion).
consider the sodium atom, which has 11 protons and 11 electrons.
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Chlorine, with 17 protons and 17 electrons
 In general, metal atoms tend to lose electrons to form cations and
nonmetal atoms tend togain electrons to form anions.
 Thus, ionic compounds tend to be composed of metals bonded with
nonmetals, as in NaCl.
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a) 48Ti3+
b) 32S2-
Predicting Ionic Charges
Many atoms gain or lose electrons to end up with the same
number of electrons as the noble gas closest to them in the
periodic table.
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Predictable charges of some common ions.
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Ionic Compounds
A compound made up of cations and anions.
The ions in ionic compounds are arranged in three-dimensional
structures, as for NaCl. Because there is no discrete “molecule” of NaCl,
we are able to write only an empirical formula for this substance.
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Ionic compounds are generally combinations of metals
and nonmetals, as in NaCl.
Molecular compounds are generally composed of
nonmetals only, as in H2O.
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We can write the empirical formula for an ionic compound if
we know the charges of the ions.
Chemical compounds are always electrically neutral.
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2.8 NAMING INORGANIC COMPOUNDS
The system used in naming substances is called chemical nomenclature,
from the Latin words nomen (name) and calare (to call).
water (H2O) and ammonia (NH3), do have traditional names (called
common names).
Names and Formulas of Ionic Compounds
ionic compounds usually consist of metal ions combined with nonmetal
ions. The metals form the cations, and the nonmetals form the anions.
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Prefixes are used when the series of oxyanions of an element extends to four
members,
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K2S
Ca(HCO3)2
Ni(ClO4)2
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