Download Chapter 2

Chapter 2
Atoms, Molecules, and
Ions
Preview
The chapter covers the following:
• Fundamental Chemical Laws and Atom.
• Modern View of Atomic Structure,
Molecules, and Ions.
• Periodic Table.
• Naming Simple compounds, Ionic
compounds, Formula from names.
2
–
2.1 The early history of
chemistry
Greeks proposed: Matter was composed of four fundamental
substances: fire, earth, water, and air.
–
–
–
Greeks also considered: Substances are composed
of small indivisible particles. Abdera & Leucippus used
the term atomos (became atoms) to describe these
ultimate particles.
Foundation of Modern Chemistry: In 1494 – 1555
the development of systematic metallurgy (extraction of
metals from ores) of chemistry science was established
by George Barer (German).
Truly Quantitative Chemical Experiment: In 1627 –
1691 the first truly quantitative experiments were
performed by Robert Boyle who measured the
relationship between pressure and volume of air.
» Homework Read Assignment to answer the following:
1.What was Boyle’s experiment?
2.Who did discover the oxygen and its effect on combustion experiment?
3
2.2 Fundamental Chemical
Laws
• 1743 – 1794, Antoine Lavoisier explained the
nature of combustion and like Boyle, he
carefully measured the weight of reactant and
product. He suggested that "Mass is neither
created nor destroyed".
• Lavoisier states that combustion involved
Oxygen and life process used oxygen – similar
to combustion.
• Lavoisier published the first modern chemistry
textbook "Elementary Treatise on Chemistry"
4
2.2 Fundamental Chemical
Laws
• 1754 – 1826, Joseph Proust determined the
composition of various chemical compounds.
Proust's Law or “Law of definite proportion”
“a given compound always contains exactly the same proportion
of elements by mass"
• 1766 – 1844, John Dalton discovered that
"different relative amounts of elements may also found for similar
elements.”
{Law of multiple proportions}
Compound I
Compound II
Per 1g of carbon
1.33g of oxygen
2.66g of oxygen
•i.e. cpd. I is CO and cpd. II is CO2
5
2.2 Fundamental Chemical
Laws
Example 2.1:
The hollowing data were collected for several
compounds of nitrogen and oxygen:
Compound A
Compound B
Compound C
Per 1g of oxygen
1.75g
of Nitrogen
0.875g of Nitrogen
0.4375g of Nitrogen
Note: In fact, an infinite number of combination exists:
• N2O , NO
, NO2
• NO
, NO2 , NO4
• N4O2 , N2O2 , N2O4
6
2.3 Dalton's Atomic Theory
The modern phases of Dalton's ideas are:
1.
2.
3.
4.
Each element is made up of tiny particles called atoms.
The atoms of a given element are identical.
chemical compounds are formed when atoms combine with
each other.
Chemical reactions involve reorganization of the atoms changes
in the way they are bound together.
Dalton's theory lead to:
1gm hydrogen + 8gm of oxygen  water
he assume that water formula is "OH" and the mass
of hydrogen is "1" and of oxygen is "8".
Using the same concepts, Dalton's proposed the first
table of atomic masses.
It has been proved later that Dalton's table contain
incorrect masses.
7
.3 Dalton's Atomic Theory
Gay-lussac (1778-1850)
found experimentally
that:
Avogadro proposed that "at the same temperature and pressure, equal
volumes of different gases contain the same number of particles". If
Avogadro's hypothesis is correct, Gay-Lussaci result (figure 2.5):
Note: this suggests that the
formula for water is H2O,
not OH as Dalton believed.
8
2.3 Dalton's Atomic Theory
• Jakob Berzelius (1779 – 1848) use the same
Avogadro's concept.
• He discovered the elements cerium, selenium, silicon,
and thorium.
• He developed the modern symbols for the elements.
e.g.
Old
new symbols
1) Silver
﴿
Ag
2) Lead
5
Pb
3) Tin
4
Sn
4) Sea salt
Θ
NaCl
9
2.4 Early Experiments to Characterize
the Atom
The concept of atoms was a good idea, Then scientists began to wonder about the nature of
the atom.
• What is an atom made of?
• How do the atoms of various elements differ?
EXP # 1:
J. Thomson (1898 – 1903) use the cathode-ray-tube:
•
He observed an existence of –ve rays produced at the negative
electrode and repelled by –ve pole
•
He determine the charge –to- mass ratio of this particles:
e/m = -1.76 x 108 c/g
This particle, now called “Electrons”
10
2.4 Early Experiments to Characterize
the Atom
EXP # 2:
Millikan does another experiment to determine the mass of this
–ve particles (electrons). The experiment used mainly to
determine the magnitude of the electron charge and using e/m to
get m- value:
Mass of electron (me) = 9.11 x 10-13 kg
11
2.4 Early Experiments to Characterize
the Atom
Radioactivity
Henri Becquerel (1896) found accidentally that piece of
mineral containing Uranium could produce image on
photographic plates.
• Uranium emitted high radiation and this is called Radioactivity.
• Studies on radioactive material showed three types of
emission are expected:
g rays : high energy light
β particles : -ve particles = electrons
a particles : +ve particles has two +ve,
and mass 7300 times mass of electron
12
The nuclear Atom
Ernest Rutherford (1911) did many
experiment to explore radioactivity
found the a-particles reflected and
not completely pass through a metal
sheet as expected by Thomson's
model of the atom is correct [plum
pudding made].
Moved electron at a distance = 10-8 cm
Nucleus Diameter = 10-13 cm
Piece of nucleus material with pea size has mass equal 250 million tons!!.
13
2.5 The Modern View of Atomic
Structure
Rutherford’s Model of
the Atom:
atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
• More studies on nucleus showed that nucleus contain protons
(+ve particles) and neutrons (neutral particles) and electrons
moving around.
• Mass and charges of all of these are:
Electron
Proton
Neutron
Mass
- 31
9.11x10
kg
27
kg
1.67x10
27
1.67x10
kg
Charge
11+
None
14
2.5 The Modern View of Atomic
Structure
• The atomic components are related and con be derived from
the element symbols notation:
Mass Number
Atomic Number
A
ZX
Element Symbol
"Z" = number of protons = number of electrons for neutral atoms.
"A" = total number of protons and neutrons.
Isotopes
• Different mass number between two elements have the same
atomic number are called "isotopes".
• Isotopes: have similar chemical properties.
1
1H
2
1H
(D)
3
1H
(T)
15
2.6 Molecules and Ions



Atoms combined with each others to form
compounds.
Forces that hold atoms together in compounds
are called chemical bonds. Only, electrons
participate in bonding.
There are two types of chemical bonding:
1.
Electron sharing: this lead to covalent bond formation, and the
resulted compound is called a molecule. There are two ways to
present the molecules:
(a) Chemical Formula: (Simplest way)
e.g. H2O, CO2, NH3, C2H6O, etc
(b) Structural Formula: (more information)
It may or may not indicate the actual shape of molecule.
H2
H2O
NH3
CH4
16
2.6 Molecules and Ions
1.
2.
Electron sharing:
Attraction Among Ions: this lead to ionic bonding and the formed
compound is called ionic composure ionic solid, or salt.
Ions are atoms or group of atoms that have a net (+ve) or (ve) charge.
e.g.
cation – ion with a positive charge: If a neutral atom loses
one or more electrons it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge: If a neutral atom gains one
or more electrons it becomes an anion.
Cl
17 protons
17 electrons
Cl-
17 protons
18 electrons
17
2.5
2.6 Molecules and Ions
• Using Molecular Modeling Computer software one can build
many molecules.
18
2.5
2.7 An Introduction to the Periodic Table
The simple periodic table (figure 2.21) contains symbols that
represent elements.
• Element symbols are abbreviations based on the element name or
the original name e.g.
Hydrogen (name)
Sodium = [original name is Natriam]
19
2.7 An Introduction to the Periodic
Table
• Most of the elements are metals. They have characteristic
physical properties e.g.
– High heat and electric conduction.
– Malleability (hammered to sheet)
– Ductility (pulled into wires)
• Chemically metals tend to lose electrons to form +ve-ions. Fe2+ ,
Fe3+ , Na+ ,K+ , Ca2+
• Upper right – hand corner contains nonmetals. They lack the
physical properties of metals.
• Nonmetals tend to gain electrons to for –ve-ions. Cl-, F-, O2-, S2-…
• Nonmetals tend to bond with each other by forming covalent
bonds.
20
Periodic Table arranged in:
• vertical columns called groups or families
e.g. alkali metals, Noble gases, halogen …etc.
• Horizontal rows called periods
e.g. first period, second period …etc.
• Noble gases group exist normally as monoatomic (Single
atoms) because they have little chemical reactivity.
• all elements per group have similar chemical properties e.g.
gp I: alkali metals only form +ve ions with one +ve charge.
gp II: Alkali earth metals form +ve ions with two +ve charge.
21
2.4
22
Noble Gas
Halogen
Group
Alkali Metal
Alkali Earth Metal
Period
2.8 Naming Simple Compounds:
 The following systematic naming will be for both ionic and
covalent compounds:
1. Binary ionic compounds (type I):
Contain
positive ion
Cation
Negative ion
anion
Rule:
Cation named first as name of element.
Anion: named after cation, root + ide.
e.g. Na ClSodium Chloride.
23
1. Binary ionic compounds (type I):
2. Binary Ionic Compounds (Type II):
–
Many metals (transition metals) form more than one
type of ionic compound
e.g. Fe Cl2 and Fe Cl3.
To name these compounds we have:
Rule:
•
•
•
Determine the charge on the metal.
Use Roman numeral to indicate the charge of the cation (I, II,
III, IV,…).
Use same naming as type I.
e.g. Fe Cl2
Fe2+
Iron (II) Chloride
Fe Cl3
Fe3+
Iron (III) Chloride
24
Figure 2.22: The common cations and anions
25
Example 2.4
Give the systematic name of the following
compounds.
a. CuCl b. HgO
c. Fe2O3 d. MnO2.
Note:
for elements that will not give more than one ionic
species DON'T use Roman Numeral.
e.g.
Ag
Ag+ only
Al
Al3+ onl
Zn
Zn2+ only
gp I elements
gp II elements
26
Note:
For polyatomic ion formulas Special names are given and must
be memorized.
27


Binary ionic compounds (type I):
Binary Ionic Compounds (Type II):
 Binary compounds
(Type III):
Compounds formed from
Nonmetals and give Covalent
molecules.
Naming is similar to type I for
RHS-element and LHS-element
+ prefixes to denote the number
of atoms present. (mono, di, tri,
tetra, penta, heca, etc..)
e.g.
N2O
Dinitrogen monoxide
NO
nitrogen monoxide
28
2.8 Naming Simple
Compounds:
Figure 2.23: A flowchart for naming binary compounds
Figure 2.24: Overall strategy for naming chemical compounds
29
Example 2.8
Give the systematic name for each of
the following :
Tetra phosphorus decaoxide.
a. P4O10
b. Nb2O5 Niobium (V) oxide.
Lithium peroxide (O22- peroxide).
c. Li2O2
d. Ti(NO3)4 Titanium (lV) nitrate.
30
2.8 Naming Simple Compounds
Acids:
Acids are the molecules that produce H+ when they dissolved
in water (chapter 4).
There are two types of acids:
(a) H+ + anion does not contain oxygen. Names given by
adding prefix (hydro-) and suffix (-ic).
e.g. HCl
Hydrochloric acid
H2S
Hydrosulfuric acid
(b) H+ + anion contain oxygen. Naming by using
root + -ic …, or, root + -ous … or, Depending on
the name of the anion.
e.g.
anion ends is –ate: acid -ic
SO42- Sulfate Sulfuric acid
anion ends is – ite acid -ous
SO32-  Sulfite  Sulfurous acid
31
32
Compounds
Figure 2.25: A flowchart for naming acids. An acid is best
considered as one or more H+ ions attached to an anion.
33