Download Chapter 5: Atomic Structure

Chapter 2: Atoms,
Molecules and Ions
•
Early
Models
of
Atoms
Democritus (460-400B.C.) first suggested the existence of
these particles, which he called “atoms” for the Greek word for
“uncuttable”. They lacked experimental support due to the lack
of scientific testing at the time.
• Plato and Aristotle formulated the notion that there can be no
ultimately indivisible particles, so the “atomic” view faded for a
number of years.
• John Dalton (1766-1844) performed experiments to study the
ratios in which elements combine in chemical reactions. He
formulated hypotheses and theories to explain his observations,
which became Dalton’s Atomic Theory.
– All elements are composed of tiny indivisible particles called atoms.
– Atoms of the same element are identical. The atoms of any one
element are different from those of any other element.
– Chemical reactions occur when atoms are separated, joined or
rearranged. Atoms of one element, however, are never changed into
atoms of another element as a result of a chemical reaction.
– Atoms of different elements can physically mix together or combine
in simple, whole number ratios to form compounds.
Dalton’s Atomic Theory
• According to Dalton’s atomic theory atoms are
the smallest particles of an element that retain
the chemical identity of the element. His theory
explains several simpler laws of chemical
combination from his time.
– Law of constant composition: In a given compound, the
relative numbers and kinds of atoms are constant. (4)
– Law of conservation of mass: The total mass of
materials present after a chemical reaction tis the
same as the total mass present before the reaction (3)
– 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.
• Example: CO2 and CO
H2O2 and H2O
Discovery of Atomic Structure
• As scientists began to develop methods for more detailed
probing of the nature of matter, we discovered more. Atoms are
now known to be divisible as they can be broken down to even
smaller particles by atom smashers.
• J.J. Thomson (1856-1940) discovered electrons using cathode
ray tubes. Another CRT
• Robert Millikan (1868-1953) carried out experiments to
determine the charge of an electron (-). He also determined the
ratio of the charge to the mass of an electron.
• In 1886, E. Goldstein observed a cathode ray tube and found
rays traveling in the opposite direction to that of the cathode
rays. He called these rays canal rays and concluded that they
must be positive particles, which are now called protons.
• In 1932, James Chadwick confirmed the existence of yet
another subatomic particle: the neutron. Neutrons are
subatomic particles with no charge but with a mass nearly
equal to that of a proton. See simulation
• After discovering these subatomic particles,
scientists wondered how they were put together.
• JJ Thompson thought since the electrons
contributed such a small fraction of the atoms mass,
they were probably an equal fraction of it size so it
was like “Plum Pudding”.
• In 1911, Ernest Rutherford and his coworkers
performed the Gold Foil Experiment to further study
the phenomenon.
• Concluded that most of the mass of each atom and
all of its positive charge reside in a very small,
extremely dense region which is called the nucleus.
The rest of the atom is mostly empty space.
Modern View of Atomic Structure
• Since the time of Rutherford, physicists have learned much
about the nucleus. Although many other parts have been
discovered, chemists tend to only work with three main
particles since they determine chemical behavior: Electron,
Neutron and Proton
• Electron has a charge of -1.602 X 10-19 C and a proton has a
charge of 1.602 X 10-19 C so this quantity of Coulombs is
known as one electronic charge and atomic and subatomic
particles usually have a charge that is multiples of this.
Neutrons have no charge and are electrically neutral.
• Atoms have extremely small masses so instead of using
the real numbers, atomic mass units (amus) are used.
Protons and neutrons are very similar in mass but it would
take 1836 electrons to equal 1 proton so most of an atoms
mass is in the nucleus.
• Atoms are also extremely small with diameters between 1
X 10-10 and 5 X 10-10 so they are usually expressed with
angstroms, which is 10-10.
Illustrating the Size of an Atom
• The diameter of a US penny is 19 mm.
The diameter of a silver atom, by
comparison is only 2.88 A. How many
silver atoms could be arranged side by
side in a straight line across the
diameter of a penny?
19 mm 1X10-3m 1A 1 Ag atom = 6.6 X107 Ag atoms
1mm 1X10-10m 2.88 A
This is over 66 million silver atoms could sit side by
side across a penny!
Atomic Number
• The number of protons in the nucleus
of an atom of that element, which is the
primary difference that distinguishes
each element.
• For an atom with no charge, this is also
the number of electrons since the
positive charge of the protons cancels
the negative charge of the electrons.
Mass Number
• Most of the mass of an atom is found in the
nucleus so the total number of protons and
neutrons equals the mass number.
• If you know the atomic number and mass number
you can determine the composition of that atom.
• The composition can be represented by the
shorthand notation using the element symbol,
atomic number and mass number.
• For gold, Au is the symbol for the element and the
atomic number is subscript and mass number is
superscript on the left side.
197
Au
79
• Do Sample Exercises 2.3 and Practice Exercises
in that box on pg 46
Isotopes
• Atoms that have the same number of
protons but different number of
neutrons.
• Affects the shorthand notation of the
element.
• Do Sample Exercises 2.3 and Practice
Exercises in that box on pg 46
Atomic Mass
• Today we can determine the masses of individual atoms
with a relative high degree of accuracy but since they
are so small atomic mass units are used with hydrogen
being 1 amu.
• The average atomic mass for an element due to the
different isotopes, the mass of those isotopes and the
natural percent abundance. It is also known as atomic
weight.
• Add up the different atomic mass of each atom and then
divide by the number of atoms.
• Or, multiply mass by % and then determine average
mass.
• Sample Exercise 2.4 and Practice Exercise on pg 47
Mass Spectrometer
• The most direct and accurate means for
determining atomic and molecular
weights. See pg 48
Periodic Table
• The arrangement of elements in order of
increasing atomic number, with elements
having similar properties placed in vertical
column.
• Atomic number, symbol, name, atomic
weight are found in each square for each
element. Some tables have additional
information as well. Example
• Can be arranged according to metals, nonmetals and metalloids, solid liquid and
gases, and by family. Example
Molecules and Molecular
Compounds
• Even though the atom is the smallest representative sample of
an element, only the noble gas elements are normally found in
nature as isolated atoms. All others form either molecules or
ions.
• A molecule is an assembly of two or more atoms tightly bound
together by a covalent bond created by two atoms sharing
electrons.
• Diatomic atoms form diatomic molecules (remember 7… start at
7 form a 7 and hydrogen).
• Compounds that are composed of molecules that contain more
than one type of element are molecular compounds.
• Most molecules are composed of nonmetals.
• Chemical formulas that indicate actual number and types of
atoms in a molecules are called molecular formulas. Such as
H2O, C6H12O6, and C2H4.
• Empirical formulas give only the relative number of atoms, they
are basically the reduced formula. Such as H2O, CH2O, and CH2.
• Do Sample Exercise 2.6 and Practice Exercise on pg 53.
Picturing Molecules
• The molecular formula of a substance describes
the composition but doesn’t show how they come
together.
• Structural formula: shows which atoms are
attached to which.
– Atoms are represented by their symbol and the bonds
are represented by lines.
• Perspective Drawing: shows actual geometry to
give some sense of three-dimensional shape.
• Ball-and-stick Models: shows atoms a spheres and
bond as sticks. Accurately represents the angles at
which the atoms are attached to one another within
the molecules.
• Space-filling Model: shows what the molecule
would look like if the atoms were scaled up to size.
Ions
• Some atoms can gain or lose electrons to try and get the
same number of electrons as the nearest noble gas, when
an electron is gained or lost from a neutral atom a
charged particle occurs called an ion.
• An ion with a positive charge (lost an electron) is called a
cation, where as an ion with a negative charge (gained an
electron) is called an anion.
• In general, metals atoms tend to lose electrons to form
cations and nonmetals tend to gain electrons to form
anions.
• In addition to simple single atom ions, there are
polyatomic ions, which consist of atoms joined as a
molecule but they have a net positive or negative charge.
• Ionic charge can be predicted by determining how many
electrons an atom has to lose to become like the nearest
stably arranged noble gas.
• Do Sample Exercises 2.7 and 2.8 and Practice Problems
on pg 55.
Ionic Compounds
• When a positive ion such as Na comes close to a
negative ion such as Cl, their opposite charges are
attracted and form an ionic compound connected by a
ionic bond.
• Generally, they are combinations of metals and
nonmetals such as Na and Cl.
• Ions in ionic compounds are arranged in threedimensional structures.
• The formula for an ionic compound is always an
empirical formula (most reduced form) because there is
no discrete molecule of NaCl.
• Chemical compounds are always electrically neutral, so
the empirical formula shows the ratio of the ions for this
to be true.
• For example, Mg2+ and N3- would have to be Mg3N2.
• Sample Exercise 2.10 and Practice Exercises on pg 58
Naming Inorganic Compounds
• To obtain information about a particular
substance you must know its chemical name and
formula, the system used for this is chemical
nomenclature. Some compounds also have
common names in addition to their chemical
nomenclature such as water.
• The rules for naming a compound is based on
divisions of substances into categories. The
major division is between inorganic and organic.
• Among the inorganic compounds the three basic
divisions are ionic compounds, molecular
compounds and acids.
Naming Positive Ions
(Cations)
• Cations formed from metals atoms have the
same name as the metal found on the periodic
table. These are monatomic ions.
– Mg2+Magnesium ion and K+ Potassium ion
• If a metal can form different cations, the
positive charge is indicated by a roman
numeral in parentheses following the name of
the metal. These are usually transition metals.
– Cu2+ Copper (II) ion and Cu+ Copper (I) ion
• Cations formed from nonmetal atoms have
name that end in –ium. These are polyatomic
ions.
– NH4- Ammonium ion and H3O+  Hydronium ion
Naming Negative Ions
(Anions)
• The names of monatomic anions are formed by
replacing the ending of the name of the element
with –ide.
– O2- oxide ion and N3- nitride ion
• Polyatomic anions containing oxygen have
names ending in -ate or -ite.
– NO3-  nitrate ion and NO2- nitrite ion
• Anions derived by adding H+ to an oxyanion are
name by adding hydrogen or dihydrogen as a
prefix as appropriate.
-
– HCO3 -hydrogen carbonate ion and
H2PO4 dihydrogen phosphate ion
Naming Ionic Compounds
• Names of ionic compounds consist of
the cation name followed by the anion
name.
• Do Sample Exercises 2.12 and 2.13 and
Practice Exercises on pg 63
Naming Acids
• You know a molecule is an acid because its
cation is hydrogen.
• Acids containing anions whose names end in
-ide are named by changing the -ide ending to -ic
adding the prefix hydro- to this anion name and
then following with the word acid.
– HClHydrochloric Acid
• Acids containing anions whose name end in -ate
or -ite are named by changing the -ate to -ic and
-ite to -ous and then adding the word acid.
– HNO3Nitric Acid & HClO2Chlorous Acid
• Do Sample Exercise 2.14 and Practice Exercise
on pg 64-65.
•
Naming Binary Molecular
Compounds
The name of the element farther to the left in the
periodic table is usually written first. Except Oxygen is
written last with all except Flourine.
• If both elements are in the same group in the periodic
table, the one having the higher atomic number is
named first.
• The name of the second element is given and -ide
ending.
• Greek prefixes are used to indicate the number of atoms
of each element. Although mono is never used with the
first element.
– mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca
• N2O5  Dinitrogen pentoxide
• Sample Exercise 2.15 and Practice Exercise on pg 65