CHEM 120: Introduction to
Inorganic Chemistry
Instructor: Upali Siriwardane (Ph.D., Ohio State
University)
CTH 311, Tele: 257-4941, e-mail:
upali@chem.latech.edu
Office hours: 10:00 to 12:00 Tu & Th ; 8:00-9:00
and 11:00-12:00 M,W,& F
Chapters Covered and Test dates
• Tests will be given in regular class periods from 9:30-10:45 a.m. on
the following days:
September 21, 2004 (Test 1): Chapters 1 & 2
• October 6,
2004(Test 2): Chapters 3, & 4
• October 20,
2004 (Test 3): Chapter 5 & 6
• November 3,
2004 (Test 4): Chapter 7 & 8
• November 15, 2004 (Test 5): Chapter 9 & 10
• November 17,
2004 MAKE-UP: Comprehensive test (Covers all
chapters
• Grading:
• [( Test 1 + Test 2 + Test3 + Test4 + Test5)] x.70 + [ Homework + quiz average] x 0.30 = Final Average
•
5
Chapter 2: The composition and
structure of the atom
Matter and structure
• Picture of matter has evolved (and is still
evolving) over the years.
• Democritus (Greek philosopher, 470-380
B.C.) --atomic theory
Composition of the atom
• The smallest unit of an element that retains
the properties of that element is called an
____
• Basic structural unit of an element is an
____
• An _____ is incredibly small.
Makeup of an atom for a chemist
• An atom is composed of
• The _________ (positively charged) and _______
(uncharged) are found in a very small, dense
portion of the atom called the nucleus .
• __________ (negatively charged) surround the
nucleus in a very diffuse region and have a much
smaller mass than the proton and neutron.
• Name
charge
mass(amu)
mass(g)
• Electron (e)
-1
5.4x10-4
9.1095x10-28
• Proton (p)
+1
1.00
1.6725x10-24
0
1.00
1.6750x10-24
• Neutron (n)
Elemental Symbols
• How are elements given symbols?
• Chemical symbols can be one or two letters. The first
letter is always a capital case and the second letter is
always a small case. Some symbols are taken from the
Latin or German names of elements.
• Na = sodium, K = potassium, Fe = iron, Cu = copper, Ag
= silver, Sn = tin, Sb = antimony, W = tungsten,
Au = gold, Hg = mercury, Pb = lead
Symbolic notation for element
• Mass no.
A
c
charge
X
Atomic no.
Z
• X is the chemical symbol for the element
• The atomic no.(Z) is the
• The mass no.(A) is the
• If the mass number (A) = total no. of
protons and neutrons in the nucleus,
• The number is neutrons is?
• Number of neutrons =
• For neutral atoms, the no. of protons in the
nucleus = no. of electrons outside of the
nucleus and the overall charge is zero.
Isotopes (or are all atoms of a
given element the same?)
• Atoms with the same atomic number but
different mass number (therefore diff. nos.
of neutrons) are called isotopes
• Most elements have two or more isotopes.
•
1 H
1
2
1H
3
1H
•
33
S
16
• How many protons?
• How many neutrons?
• How many electrons?
• Is it necessary to include the atomic no.?
• What is the atomic number of bromine?
• How many protons does a Br atom have?
• How many neutrons does a Br atom with
mass number 79 have?
• How many electrons does a (neutral) Br
atom have?
• How many protons, neutrons, and electrons
are in
•
209 Po
84
•
136 Ba
56
Average atomic mass
• The atomic masses of an element are the
weighted averages of all the isotopes of
that element taking into account the relative
abundance each isotope.
• Average at.mass =
%abundance isotope 1/100% x mass isotope 1 +
%abundance isotope 2/100% x mass isotope 2 +
etc…...
Isotopic mass problem
• The atomic masses of the two stable
isotopes of boron, 105B (19.78%) and 115B
(80.22%) are 10.0129amu and 11.0093 amu
respectively. What is the “average” atomic
mass of boron?
Question 2.3
• The element neon has three naturally
occurring isotopes. One of these has a mass
of 19.99 amu and a natural abundance of
90.48%. A second isotope has a mass of
20.99 amu and a natural abundance of
0.27%. A third has a mass of 21.99 amu
and a natural abundance of 9.25%.
Calculate the atomic mass of neon.
Harder isotope problem
• A hypothetical atom has two isotopes only.
• Isotope one has a percent abundance of
60%.
• If the average isotopic mass is 120.0 amu
what are the masses of isotopes one and
two?
Ions
• Ions are charged particles that are a result of
the atom
• _______ have more electrons than protons
and are negatively charged. The original
atom has gained electron(s).
• ________ have more protons than neutrons
and are positively charged. The original
atom has lost electron(s).
• Charge on an ion = no. of protons - no. of
electrons.
gains one electron g
•
37 Cl
17
•
138 Ba
56
loses two electrons g
Like charges repel each other, opposite attract.
• 2.24: Write the symbol for an isotope
• that contains 92 protons and 146 neutrons
• 2.30: Which are true?
• An atom with an atomic number of 7 and a mass
number of 14 is identical to an atom with an
atomic number of 6 and a mass number of 14.
• Neutral atoms have the same number of electrons
as protons.
• The mass of an atom is due to the sum of the no.
of protons, neutrons and electrons.
• 2.28 from end of chapter
Development of the atomic
theory
• Democritus (Greek philosopher) fifth
century B.C. : matter consists of very small,
indivisible particles--atomos (atoms-uncuttable)
• Plato and Aristotle not accept this idea.
Dalton’s Atomic Theory (1808)
• Marked beginning of modern era of
chemistry
Dalton’s hypotheses
• Elements are composed of extremely small particles, called atoms.
• All atoms of an element are identical (same size, mass, chem. prop).
• The atoms of one element are different from the atoms of all other
elements.
• An atom cannot be created, divided, destroyed or converted into any
other type of atom.
• Compounds are composed of atoms of more than one element in
simple whole-number ratios.
• A chemical reaction involves the separation, combination, or
rearrangement of atoms, not their creation or destruction.
X2Y
Law of definite proportions
Which of Dalton’s postulates are
considered true today?
• Elements are composed of extremely small particles, called atoms.
• All atoms of an element are identical (same size, mass, chem. prop).
• The atoms of one element are different from the atoms of all other
elements.
• An atom cannot be created, divided, destroyed or converted into any
other type of atom.
• Compounds are composed of atoms of more than one element in
simple whole-number ratios.
• A chemical reaction involves the separation, combination, or
rearrangement of atoms, not their creation or destruction.
Subatomic particles
• Is the atom really indestructible, or does it
consist of even smaller particles?
Electrons
• 1890’s: cathode ray tube experiments (tube
sealed with metal electrodes in it and
evacuated of air)
• Apply high voltage source, invisible ray
produced (see effect by fluorescence when
ray strikes coated surface).
Electrons (fig 2.3)
A. Magnetic
Field
C. Electric
field
B. No field or
balanced fields.
Thomson determined
charge/mass ratio of -1.76x108C/g)
• Find rays have same properties regardless of metal
used in constructing the cathode.
• Experiments show that cathode rays are made of
charged particles that interact with electric and
magnetic field when moving.
• Particles are negatively charged (repelled by the
negative plate, attracted toward the positive plate).
• These negative particles are fundamental particles
of matter. Called electrons. (1897 Thomson)
Are there other particles that
make up the atom?
• Atoms are neutral and contain electrons
which are negatively charged.
• Therefore there must be something positive
present also.
• Protons, which are positively charged, were
discovered by Goldstein
• Protons have mass of 1.67262 x 10-24 g
(1840 times greater than electron mass) and
charge equal but opposite in sign from e-.
Atomic strucure model:
or where might the positive charge be?
Thomson’s plum pudding
model is a uniform positive
sphere where negative
electrons embedded in it.
Natural radioactivity
• Spontaneous emission of particles and/or radiation.
a rays = positively charged helium nuclei
b rays = electrons (negatively charged particles)
g rays = high-energy radiation (photons), with no
charge
Scattering of a particles by gold foil
Experiment of Geiger
Majority of a particles undeflected or slightly deflected but some
displayed large deflections.
Rutherford picture of atom from
scattering expt (1910)
• Atom is mostly empty space with positive charge
(protons) concentrated in a dense central core
called the nucleus. (Neutrons are also in the
nucleus--Chadwick experimental evidence in 1932
for neutron)
• Positive core of atom deflects a particles strongly.
• ‘It was almost as incredible as if you fired a 15inch shell at a piece of tissue and it came back and
hit you.”
2.4: Relationship between light
and atomic structure
Light and atomic structure
• Rutherford atom doesn’t say much about
electron location except that they’re in a
region outside of the nucleus that is mostly
empty space.
• Let’s look at electromagnetic spectrum of
light. (energy and wavelength)
roygbiv
High energy
Short wavelength
Electromagnetic spectrum
Low energy
Long wavelength
7.1
• Pass ordinary light through a prism get
continuous spectrum of all wavelengths
• But if look at emitted light from a tube
containing hydrogen or another gas get an
emission spectrum like
Where did all the colors go?
Bohr atom
• Electrons orbit nucleus like a planet around
the sun in circular orbits (held
electrostatically).
• Hydrogen atom consists of 1 electron
orbiting 1 proton
• Electron can only be located in certain
stable orbits. Bohr assumed that the energy
of the e-’s orbit and its radius are ________.
Not all energies or radii are allowed.
• Moving from one orbit (quantum level) to
another causes atom to absorb or emit a
photon (particle of light) of energy
Color and energy
of line in spectrum
correspond to
difference in
energy btn 2 levels
Summary of results of Bohr’s
theory (p 47)
• Light energy can be absorbed and emitted by
promotion and relaxation of electrons from one
energy level to another--see a line in spectrum
corresponding to energy difference (photon
emission) btn levels.
• Don’t see all colors, just those that correspond to
energy difference btn levels
• Electrons can be found only in certain allowed
energy levels (orbits). Not all energies or radii of
orbits are allowed--quantized.
• As orbits get further from nucleus, energy of orbit
increases.
• According to Bohr one can know the location and
energy of an electron in an atom with certainty.
• Read the summary of Bohr theory on p 46
Modern atomic theory
• Bohr model only good for one electron
atoms and quantization assumed.
• Later developments:
• deBroglie noted that electrons had both
wave and particle properties: wave-particle
duality of matter. Need both concepts to
describe electrons.
• Heisenberg Uncertainty Principle: It is
impossible to know simultaneously how fast
an electron is moving and its position with
certainty.
• This leads to:
• Electrons do not move around the nucleus
in well-defined orbits but are located in
orbitals which are regions in space where
there is a large probability of finding an
electron (electron cloud).
• These electron clouds are denser in some
regions than others. The electron density is
proportional to the probability of finding the
electron at any point in time.