Transcript Chapter2_20130808125951x

LOOK AROUND YOU
 The great variety of colors in a garden scene
 The texture of the fabric in your clothes
 The solubility of sugar in a cup of coffee
 The transparency of a window
HOW DO WE EXPLAIN?
 What makes diamonds transparent and hard,
while table salt is brittle and dissolves in water?
 Why does paper burn, and why does water
quench fires?
 Where does the beautiful colors of flowers come
from?
 The structure and behavior of atoms are key to
understanding the properties of matter.
VARIETY OF ELEMENTS
 The diverse properties results from only about
100 different elements
 How do atoms combine with one another?
 What rules govern the ways in which atoms can
combine?
 How do the properties of a substance relate to
the kinds of atoms it contains?
 What is an atom like, and what make their
difference?
2.1 THE ATOMIC THEORY OF MATTER
 HISTORY
 Democritus and Greek philosophers (BC 400)
• The material world must be made up of tiny indivisible
particles
• atomos: indivisible or uncuttable
 Plato and Aristotle
• There can be no ultimately indivisible particles
• The “atomic” view of matter faded for many centuries
 Newton (1642-1727)
• Air is composed of something invisible and in constant
motion
• Still very different from thinking of atoms as the
fundamental building blocks
2.1 THE ATOMIC THEORY OF MATTER
 DALTON’S ATOMIC THEORY
 Chemists learned to measure the amounts of elements
that reacts with one another to form new substances
 Dalton’s atomic theory
• Introduced during the period from 1803 to 1807
• The theory was based on the four postulates given in
the figure in the next page
2.1 THE ATOMIC THEORY OF MATTER
 DALTON’S POSTULATES
 Dalton’s theory explains several simple laws of chemical
combination.
 The law of constant composition
• In a given compound, the relative numbers and kinds of
atoms are constant
 The law of conservation of mass (matter)
• The total mass of materials present after a chemical
reaction is the same as the total mass present before
the reaction
 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
(H2O and H2O2)
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
 ATOMIC IMAGES
 Dalton had no direct evidence for the existence of atoms
 Scientists have developed methods for more detailed
probing of the nature of matter
 Today, we can measure the properties of individual
atoms and even provide images of them
FIGURE 2.2 An image of the
surface of silicon obtained by
scanning tunneling microscopy
(STM)
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
 CATHODE RAYS AND ELECTRONS
 Thomson found that cathode rays are streams of
negatively charged particles
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
 CATHODE RAYS AND ELECTRONS
 The charge/mass ratio: 1.76 X 108 C/g
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
 MILLIKAN’S OIL DROP EXPERIMENT
 Robert Millikan (University of Chicago) determined the
charge on the electron in 1909.
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
 RADIOACTIVITY
 Radioactivity is the spontaneous emission of radiation by
an atom.
 It was first observed by Henri Becquerel.
 Marie and Pierre Curie also studied it.
 Three types of radiation were
discovered by Ernest Rutherford:
• a particles
• b particles
• g rays
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
 THE NUCLEAR MODEL
 The prevailing theory
was that of the “plum
pudding” model, put
forward by Thomson.
 It featured a positive
sphere of matter with
negative electrons
imbedded in it.
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
 THE NUCLEAR MODEL
 Ernest Rutherford
shot a particles at a
thin sheet of gold foil
and observed the
pattern of scatter of
the particles.
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
 THE NUCLEAR MODEL
 Since some particles were deflected
at large angles, Thompson’s model
could not be correct.
 Rutherford postulated a very small,
dense nucleus with the electrons
around the outside of the atom.
 Most of the volume of the atom is
empty space.
 Protons were discovered by
Rutherford in 1919.
 Neutrons were discovered by
James Chadwick in 1932.
2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
 SUBATOMIC PARTICLES
 Protons and electrons are the only particles that have a
charge.
 Protons and neutrons have the same mass.
 The mass of an electron is so small we ignore it.
FIGURE 2.11
The structure of an atom.
2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
 SUBATOMIC PARTICLES
 Every atom has an equal number of electrons and
protons, so atoms have no net charge.
 Atomic mass unit, 1 amu = 1.66054 X 10-24 g
 Atom’s size: 1-5 Å
2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
Sample Exercise 2.1 Atomic Size
The diameter of a US dime is 19 mm, and the diameter of a silver atom is 2.88 Å.
How many silver atoms could be arranged side by side across the diameter of a
dime?
Practice Exercise
The diameter of a carbon atom is 1.54 Å. (a) Express this diameter in picometers. (b)
How many carbon atoms could be aligned side by side across the width of a pencil
line that is 0.20 mm wide?
2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
 THE DIAMETERS OF ATOMIC NUCLEI
 About 10-4 Å
 Density of nucleus: 1013 ~ 1014 g/cm3
 A match box full of nuclei would weigh over 2.5 billion tons!
2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
 ATOMIC NUMBERS, MASS NUMBERS, & ISOTOPES
 What makes the difference between carbon and oxygen?
• The atoms of each element have a characteristic number
of protons
 Atomic number: the number of protons in the nucleus
 Isotopes: atoms with identical atomic numbers but different
mass numbers.
2.5 THE PERIODIC TABLE
 The most significant tool that chemists use for organizing
and remembering chemical facts
2.4 ATOMIC WEIGHTS
 THE ATOMIC MASS SCALE
 The atomic mass unit (amu)
• Defined by assigning a mass of exactly 12 amu to an
atom of 12C

1H:
1.0078 amu, 16O: 15.9949 amu
 AVERAGE ATOMIC MASSES
 Most elements occur in nature as mixtures of isotopes.

12C:
98.93%, 13C : 1.07%
(0.9893)(12 amu) + (0.0107)(13.00335 amu) = 12.01 amu
atomic weight
▲FIGURE 2.12 A mass spectrometer.
▲FIGURE 2.12 Mass spectrum of atomic
chlorine.
2.5 THE PERIODIC TABLE
 PERIODICITY
 Many elements show very strong similarities to one another
 When one looks at the chemical properties of elements,
one notices a repeating pattern of reactivities.
2.5 THE PERIODIC TABLE
 The most significant tool that chemists use for organizing
and remembering chemical facts
2.5 THE PERIODIC TABLE
 It is a systematic catalog of the elements.
 Elements are arranged in order of atomic number.
2.5 THE PERIODIC TABLE
 The rows on the
periodic chart are
periods.
 Columns are
groups.
 Elements in the
same group have
similar chemical
properties.
2.5 THE PERIODIC TABLE
 GROUPS
 Many groups are known by their names.
 “Coinage metals”: Group 11
2.5 THE PERIODIC TABLE
 METALS, NONMETALS, AND METALLOIDS
 Nonmetals generally
differ from the metals in
appearance and in other
physical properties.
 A metalloid is a
chemical element with
properties that are inbetween or a mixture of
those of metals and
nonmetals
2.5 THE PERIODIC TABLE
 Isolated Pu
 Identified the elements having
atomic numbers 95 through 102
◄ FIGURE 2.17
 Identified element
number 106
 ACS proposed
that element
number 106 be
named
seaborgium
2.6 MOLECULES AND MOLECULAR COMPOUNDS
 MOLECULES
 Only the noble-gas elements are normally
found in nature as isolated atoms.
 A molecule is an assembly of two or more
atoms tightly bound together.
 Molecules behave in many ways as a single,
distinct object
a molecule
atoms
2.6 MOLECULES AND MOLECULAR COMPOUNDS
 MOLECULES AND CHEMICAL FORMULAS
 Many elements are found in nature in molecular
form
 Two different molecular forms of oxygen
• O2: a diatomic molecule, essential for life, odorless
• O3: a triatomic molecule, toxic, pungent smell
 Diatomic molecules
2.6 MOLECULES AND MOLECULAR COMPOUNDS
 MOLECULES AND CHEMICAL FORMULAS
 Molecular compounds are composed of two
or more different atoms
 Molecules vs Compounds
 Most molecular substances that we will
encounter contain only nonmetals.
2.6 MOLECULES AND MOLECULAR COMPOUNDS
 MOLECULAR AND EMPIRICAL FORMULAS
 Molecular formulas
• H2O
• H2O2
• C2H4
• C6H12O6
 Empirical formulas
• H2O
• HO
• CH2
• CH2O
 Why do we need empirical formulas?
• Certain common methods of analyzing
substances lead to the empirical formula only
2.6 MOLECULES AND MOLECULAR COMPOUNDS
 PICTURING MOLECULES
 A structural formula shows which
atoms are attached to which within the
molecule.
 A perspective drawing gives some
sense of three dimensional shape
 Ball-and-stick models show the
accurate angles between bonds
 A space-filling model shows the
relative sizes of the atoms.
2.7 IONS AND IONIC COMPOUNDS
 CATIONS AND ANIONS
 When atoms lose or gain electrons, they become
ions.
• Cations are positive and are formed by elements on the
left side of the periodic chart (metal atoms).
• Anions are negative and are formed by elements on the
right side of the periodic chart (nonmetal atoms).
2.7 IONS AND IONIC COMPOUNDS
 PREDICTING IONIC CHARGES
 Many atoms gain or lose e- to make the same
number of e- as the noble gas.
Figure 2.20. Predictable charges of some common ions
2.7 IONS AND IONIC COMPOUNDS
 IONIC COMPOUNDS
 A compound that contains both positively and
negatively charged ions.
 Generally combinations of metals and nonmetals
2.7 IONS AND IONIC COMPOUNDS
 WRITING EMPIRICAL FORMULAS
FOR IONIC
COMPOUNDS
 The ionic compound formed from Mg2+ and N3-
 97% of the mass of most organisms: O, C, H, N, P, and S
 70% of the mass of most cells: H2O
 C is the most prevalent element in the solid components of
cells
2.8 NAMING INORGANIC COMPOUNDS
 NAMES AND FORMULAS OF IONIC COMPOUNDS
 Positive ions (Cations)
• Cations formed from metal atoms
• Ions of the same element that have different charges
exhibit different properties
• Metals that form only one cation
- group 1A/2A, Al3+, Ag+, Zn2+
• Cations from nonmetals
2.8 NAMING INORGANIC COMPOUNDS
Fe(III)
Fe(II)
2.8 NAMING INORGANIC COMPOUNDS
 NAMES AND FORMULAS OF IONIC COMPOUNDS
 Negative ions (Anions)
• Monatomic and simple polyatomic anions
• Oxyanions
2.8 NAMING INORGANIC COMPOUNDS
 NAMES AND FORMULAS OF IONIC COMPOUNDS
 Negative ions (Anions)
• Anions containing H+
• Older method:
HCO3- bicarbonate ion; HSO4- bisulfate ion
(a) SeO42- (b) SeO32-
2.8 NAMING INORGANIC COMPOUNDS
2.8 NAMING INORGANIC COMPOUNDS
 NAMES AND FORMULAS OF IONIC COMPOUNDS
 Ionic compounds
• Cation name followed by anion name
2.8 NAMING INORGANIC COMPOUNDS
 NAMES AND FORMULAS OF ACIDS
2.8 NAMING INORGANIC COMPOUNDS
 NAMES AND FORMULAS OF ACIDS
(a) Hydrocyanic acid or hydrogen
cyanide.
(b) Nitric acid
(c) Sulfuric acid.
(d) Sulfurous acid
(a) HBr, (b) H2CO3
2.8 NAMING INORGANIC COMPOUNDS
 NAMES AND FORMULAS OF
BINARY MOLECULAR COMPOUNDS
(a) sulfur dioxide, (b) phosphorus pentachloride,
and (c) dichlorine trioxide.
(a) SiBr4, (b) S2Cl2
2.9 SOME SIMPLE ORGANIC COMPOUNDS
 ALKANES
 Organic compounds
• Compounds that contain carbon
 Alkanes
• Compounds that contain only carbon and hydrogen
2.9 SOME SIMPLE ORGANIC COMPOUNDS
 SOME DERIVATIVES OF ALKANES
 An alcohol is obtained by replacing an H atom with a –OH
group