Chapter 1: Atoms and Elements

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Chapter 1 Atoms and Elements
KTT 111/3 Inorganic Chemistry I
Dr. Farook Adam
August 2005
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Chapter 1 Atoms and Elements

• Chemistry is a science that studies the
  composition and properties of matter
• Matter is anything that takes up space and has
  mass
• Mass is a measure of the amount of matter in a
  sample
• Chemistry holds a unique place among the sciences
  because all things are composed of chemicals
• A knowledge of chemistry will be valuable
  whatever branch of science you study

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• Chemistry is constantly changing as new
  discoveries are made by researchers
• Researchers use a commonsense approach to the
  study of natural phenomena called the scientific
  method
• A scientific study normally
• Begins with a question about nature
• Involves a search of the work of others
• Requires observing the results of experiments
• Often results in a conclusion, or a statement
  based on what is thought about a series of
  observations

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• Experiments provide empirical facts
• Facts are called data
• A broad generalization based on the results of
  many experiments is called a (scientific) law
• Laws are often expressed as mathematical
  equations
• Laws summarize the results of experiments

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• Theoretical models attempt to explain why
  substances behave as they do
• A hypothesis is a tentative explanation
• A theory is an experimentally tested explanation
  of the behavior of nature

The scientific method is dynamic observations
lead to laws, which suggest new experiments,
which may lead to or change a hypothesis, which
may produce a theory.
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• Chemical substances are comprised of atoms
• Atoms combine to form molecules which can be
  represented in a number of ways, including

• Using chemical symbols and lines for
  connections
• A 3-D ball-and-stick model
• A 3-D space-filling model

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• Characteristics or properties of materials
  distinguish one type of substance from another
• Properties can be classified as physical or
  chemical
• Physical properties can be observed without
  changing the chemical makeup of the substance
• Chemical properties involve a chemical change and
  result in different substances
• Chemical changes are described by chemical
  reactions

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• Properties can also be described as intensive or
  extensive
• Intensive properties are independent of sample
  size
• Examples sample color and melting point
• Extensive properties depend on sample size
• Examples sample volume and mass
• In general, intensive properties are more useful
  in identifying a substance
• Matter is often classified by properties

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• The three common physical states of matter have
  different properties
• Solids have a fixed shape and volume
• Particles are close together and have restricted
  motion
• Liquids have indefinite shape but fixed volume
• Particles are close together but are able to flow
• Gases have indefinite shape and volume
• Particles are separated by lots of empty space

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• Elements are substances that cannot be decomposed
  by chemical means into simpler substances
• Each element is assigned a unique chemical symbol
• Most are one or two letters
• First letter is always capitalized
• All remaining letters are lowercase
• Names and chemical symbols of the elements are
  listed on the inside front cover of the book

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• Compounds are substances formed from two or more
  different elements combined in a fixed proportion
  by mass
• The physical and chemical properties of a
  compound are, in general, different than the
  physical and chemical properties of the elements
  of which it is comprised
• Elements and compounds are examples of pure
  substances whose composition is the same,
  regardless of source

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• A mixture consists of varying amounts of two or
  more elements or compounds
• Homogeneous mixtures or solutions have the same
  properties throughout the sample
• Heterogeneous mixtures consist of two or more
  phases
• Matter can be classified

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Tutorial

• Anda boleh daftar di bilik 015A Pudat Pengajian
  Sains Kimia mulai hari ini!!!

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• We take for granted the existence of atoms and
  molecules
• The concept of the atom had limited scientific
  usefulness until the discovery of two important
  laws the Law of conservation of mass and the Law
  of Definite Proportions
• These laws summarized the results of the
  experimental observations of many scientists

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• Law of Conservation of Mass
• No detectable gain or loss of mass occurs in
  chemical reactions. Mass is conserved.
• Law of Definite Proportions
• In a given chemical compound, the elements are
  always combined in the same proportions by mass.
• In the sciences mass is measured in units of
  grams (symbol, g)
• One pound equals 453.6 g

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• The laws of conservation of mass and definite
  proportions provided the experimental foundation
  for the atomic theory
• Daltons Atomic Theory (1803 AD)
• Matter consists of tiny particles called atoms.
• Atoms are indestructible. In chemical reactions,
  the atoms rearrange but they do not themselves
  break apart.
• In any sample of a pure element, all the atoms
  are identical in mass and other properties.
• The atoms of different elements differ in mass
  and other properties.
• In a given compound the constituent atoms are
  always present in the same fixed numerical ratio.

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Support for Daltons Atomic Theory The Law of
Multiple Proportions Whenever two elements form
more than one compound, the different masses of
one element that combine with the same mass of
the other element are in the ratio of small whole
numbers.
Each molecule has one sulfur atom, and therefore
the same mass of sulfur. The oxygen ratio is 3 to
2 by both mass and atoms Sample experimental
data Mass
Mass Compound Size S
O Sulfur dioxide 2.00 g 1.00 g 1.00 g Sulfur
trioxide 2.50 g 1.00 g 1.50 g
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• It follows from Daltons Atomic Theory that atoms
  of an element have a constant, characteristic
  atomic mass or atomic weight
• For example, for any sample of hydrogen fluoride
• F-to-H atom ratio 1 to 1
• F-to-H mass ratio 19.0 to 1.00
• This is only possible if each fluorine atom is
  19.0 times heavier than each hydrogen atom

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• It turns out that most elements in nature are
  uniform mixtures of two or more kinds of atoms
  with slightly different masses
• Atoms of the same element with different masses
  are called isotopes
• For example there are 3 isotopes of hydrogen and
  4 isotopes of iron
• Chemically, isotopes have virtually identical
  chemical properties
• The relative proportions of the different
  isotopes are essentially constant

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• A uniform mass scale for atoms requires a
  standard
• For atomic mass units (amu, given the symbol u)
  the standard is based on carbon
• 1 atom of carbon-12 12 u (exactly)
• 1 u 1/12 the mass of 1 atom of carbon-12
  (exactly)
• This definition results in the assignment of
  approximately 1 u for the mass of hydrogen (the
  lightest atom)

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• Example Naturally occurring chlorine is a
  mixture of two isotopes. In every sample of this
  element, 75.77 of the atoms are chlorine-35 and
  24.23 are chlorine-37. The measured mass of
  chlorine-35 is 34.9689 u and that of chlorine-37
  is 36.9659 u. Calculate the average atomic mass
  of chlorine.

Abundance Mass
Isotope () (u)
Contribution
Chlorine-35 75.77 34.9689
0.7577 34.9689 26.50 u Chlorine-37
24.23 36.9659
0.2423 36.9659 8.957 u
(Rounded) Total
35.46 u
The average mass of 1 atom of chlorine in nature
is 35.46 u.
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• Experiments have been performed that show atoms
  are comprised of subatomic particles
• There are three principal kinds of subatomic
  particles
• Proton carries a positive charge, found in the
  nucleus
• Electron carries a negative charge, found
  outside the nucleus, about 1/1800 the mass of a
  proton
• Neutron carries no charge, found in the
  nucleus, a bit heavier than a proton, about 1800
  times heavier than an electron

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• An element can be defined as a substance whose
  atoms all contain the identical number of
  protons, called the atomic number (Z)
• Isotopes are distinguished by mass number (A)
• Atomic number, Z number of protons
• Mass number, A (number of protons) (number of
  neutrons)
• For charge neutrality, the number of electrons
  and protons must be equal

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• This information can be summarized
• Example For uranium-235
• Number of protons 92 ( number of electrons)
• Number of neutrons 143
• Atomic number (Z) 92
• Mass number (A) 92 143 235
• Chemical symbol U
• Summary for uranium-235

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• The Periodic Table summarizes chemical and
  physical properties of the elements
• The first Periodic Tables were arranged by
  increasing atomic mass

Dmitri Mendeleev (Rusia)
1869
Julius Lothar Meyer (German)
Both these researchers drew out the first
periodic table independently of each other.
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Mendeleev have been honored as the first person
to arrange the elements in the form of a table
because he reported his findings to the Russian
Chemical Society a few months earlier than
Meyer!!!!
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The Modern Periodic table is arranged by
increasing atomic number Elements are arranged
in numbered rows called periods The vertical
columns are called groups or families (group
labels vary)
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• Modern Periodic Table with group labels and
  chemical families identified

Lanthanides
Actinides
Note Placement of elements 58 71 and 90 103
saves space
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• Some important classifications
• A groups representative elements or main group
  elements
• I A alkali metals
• II A alkaline earth metals
• VII A halogens
• VIII noble gases
• B groups transition elements
• Inner transition elements elements 58 71 and
  90 103
• 58 71 lanthanide elements
• 90 103 actinide elements

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• Classification as metals, nonmetals, and
  metalloids

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• Metals
• Tend to shine (have metallic luster)
• Can be hammered or rolled into thin sheets
  (malleable) and can be drawn into wire (ductile)
• Are solids at room temperature and conduct
  electricity
• Nonmetals
• Lack the properties of metals
• React with metals to form (ionic) compounds
• Metalloids
• Have properties between metals and nonmetals

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The End

• Do the exercises at the end of chapter ONE.
• For your own practice only!!
• You do not have to pass up this assignment.

Any questions?