Transcript Slide 1

Two experimentally determined facts have achieved the status of
scientific laws:
•The Law of Conservation of Mass: Mass is neither created or
destroyed in a chemical reaction. The total mass of all of the reacting
substances is the same as the total mass of all of the products.
•The Law of Constant Composition: The elements present in a
compound are present in fixed and exact proportion by mass,
regardless of the compound’s source or method of preparation.
•The percent composition of a compound represents the mass percent
that each element present contributes to the total mass.
•The percent composition of a compound is the same regardless of the
size of the sample analyzed.
•For a compound containing only carbon (C), hydrogen (H), and oxygen
(O), the percent composition would be calculated:
•Mass % C = (mass of carbon in sample) / mass of sample * 100%
•Mass % H = (mass of hydrogen in sample) / mass of sample * 100%
•Mass % O = (mass of oxygen in sample) / mass of sample * 100%
It should be obvious that:
•Mass C + Mass H + Mass O = mass sample
•Mass % C + Mass % H + Mass % O = 100%
•John Dalton was an English scientist who lived from 1766–1844. Dalton
developed the atomic theory to explain the chemical behavior of
•His theory contained five major ideas:
•All matter is composed of infinitesimally small particles called atoms.
•The atoms of any one element are identical.
•Atoms of different elements have different masses.
•Compounds are combinations of atoms of different elements and possess
properties different from their component elements.
•In chemical reactions, atoms are exchanged between starting compounds to
form new compounds. Atoms are neither created nor destroyed.
Dalton’s ideas about chemical reactions correspond to the following
The shapes (atoms) have simply been rearranged.
Because atoms are so small their masses cannot be directly determined
by ordinary means. Dalton recognized this and proposed a relative mass
•The mass of a hydrogen atom was assigned a relative mass of 1.0.
•He assumed that binary chemical compounds consisted of a 1:1
ratio of the two atoms involved.
•By decomposing a compound into its elements and weighing them
he could then assign a relative mass to other elements.
•Hydrogen chloride can be determined to be 97.24% Cl and 2.76% H.
•In 100 grams of HCl there are 97.24 g of Cl and 2.76 g of H.
•We don’t know how many H and Cl atoms there are, but we assume
that there are equal numbers of each, N.
(Mass of N H-atoms) / (Mass of N Cl-atoms) = 2.76 / 97.24
(Mass of 1 H-atom) / (Mass of 1 Cl-atom) = 2.76 / 97.24
(Mass of 1 H-atom) / (Mass of 1 Cl-atom) = 1 / 35.23
Therefore, if the mass of 1 H atom is 1 mass unit, the mass of 1 Cl
atom must be 35.23 mass units.
•This process can be repeated with compounds containing an element of
unknown atomic mass and an already known element, and eventually
the relative masses of all of the elements can be determined.
•The obvious problem for Dalton was that most chemical compounds do
not contain equal numbers of atoms of their constituent elements (i.e.
•Dalton’s original table of relative weights gave a value of about 8 for
oxygen rather than the correct value of 16.
•The relative masses of atoms are called the atomic masses of the
elements, and are often referred to as atomic weights.
•The current periodic table defines the relative atomic masses to that of
a form of carbon atoms whose mass is defined to be exactly 12 atomic
mass units, or amu (the SI symbol is u).
Atoms consist of three types of particles: protons, neutrons, and
Protons and neutrons make up most of the mass of the atom and are
located in a small volume at the center of the atom called the nucleus.
•Each proton in the nucleus is positively charged.
•The negatively charged electrons are spread out around the nucleus and
take up most of the volume of the atom.
•In a neutral atom there are equal numbers of protons and electrons.
The number of protons in the nucleus of an atom is called its atomic
•An atom of an element may gain or lose electrons to become electrically
charged, but never gains or loses protons in the nucleus.
•In determining the mass of an atom, the masses of the electrons
around the nucleus are usually ignored because they are so small
compared to the masses of the protons and neutrons.
•All atoms of a given element have the same number of protons in their
nuclei, but may have different numbers of neutrons.
•Atoms of an element containing different numbers of neutrons are
called isotopes of each other.
•Many of these isotopes are radioactive and are used in medicine and in
scientific research.
•The symbol for an isotope is usually written as the symbol for the
element with its atomic number (number of protons) as a subscript at
the lower left and the sum of the numbers of its protons and neutrons as
a superscript at the upper left.
•An alternate notation is to follow the name of the element with the
atomic mass number of the isotope:
Carbon-12, carbon-13, etc.
•Two scientists, Dmitri Mendeleev in Russia and Lothar Meyer in
Germany, independently studied the properties of the elements as a
function of increasing mass.
•They determined that the elemental properties did not change smoothly
and continuously with atomic mass, but instead repeated periodically.
This variation is now known as the Periodic Law.
•The Periodic Law is embodied in the periodic table.
The elements are arranged in the periodic table so that those with similar
chemical properties are arranged in the same vertical column, called a
group or a family.
Each horizontal row is called a period.
Elements in group VIII are called the noble gases. They are very
unreactive and do not form chemical compounds as most of the other
elements do.
The elements in group I are called the alkali metals.
Metals are shiny, malleable, can be melted and cast into shapes, and are
excellent conductors of heat and electricity.
Metals tend to lose electrons to form cations when the react chemically.
Alkali metals react violently with water to form basic solutions.
The elements in group II are called the alkaline earth metals.
•The elements in group VII are called the halogens and are nonmetals.
•Nonmetals cannot be cast into shapes and do not conduct electricity.
•Nonmetals tend to gain electrons to form anions when they chemically
The halogens, chlorine, bromine, and iodine:
•The elements between groups II and III are called the transition
•Many of these elements form cations of more than a single electrical
charge: Fe+2 and Fe+3, Cu+1 and Cu+2, etc.
•Many heavier elements are not shown in the abbreviated periodic table
•Among these are the inner transition elements: 57-70 are the
lanthanides and 89-102 are the actinides.
•Most elements are metals and are located to the left in the periodic table.
•Nonmetals are located to the upper right in the periodic table.
•Metalloids or semimetals are found between the metals and nonmetals in the
periodic table and have properties between the two.
White light, as from the sun or an incandescent bulb, can be separated
into its component colors, or visible spectrum, by means of a prism.
Atomic Spectra
When light is created from excited, or energized, atoms in a flame, a
series of sharply defined, separated lines of different colors appears in
the visible spectrum, rather than a continuous rainbow of colors.
When passed through an instrument called a spectrometer, the lines of
color can be recorded on a photographic plate as an atomic emission
A second kind of spectrum can be recorded and measured by passing
white light through the atoms of a specific element.
What would have been brightly colored lines on a black background
become the corresponding missing colors, or dark lines, on a brightly
colored rainbow.
This type of spectrum is called an absorption spectrum and can be
measured using an instrument called an atomic absorption spectrometer.
Electromagnetic Radiation and Energy
Radiation describes the transfer of energy from one point in space to
Heat and visible light are only small parts of the full spectrum of
electromagnetic radiation, which has its origin in the oscillation, or
vibration, of charged particles.
Electromagnetic radiation is usually described by its frequency, or by
how many times per second a vibration is completed.
Atomic Energy States
The basis of the line structure of atomic emission spectra was discovered
by Niels Bohr, a Danish scientist.
•Albert Einstein had discovered that a beam of light consisted of many small
packages of energy called photons.
•Bohr surmised that the electrons in atoms could exist in certain discrete
energy states.
•Moving between these energy states would only be possible by the
absorption or emission of a photon of light containing the same energy as the
difference between the two states.
•The lowest energy state in an atom is called its ground state. An atom in
its ground state cannot emit a photon of light but can absorb a photon of light
of the correct energy to reach one of many different excited states.
•Once in an excited state, an atom can emit a photon of light to return to the
ground state.
•The first model of the atom resembled the solar system with the
positive nucleus attracting an orbiting electron.
•This model, the Bohr model, although it accounted for the spectrum of
hydrogen, was proved to be incorrect.
•A new theory called quantum mechanics is now believed to correctly
account for the electronic structure of the atom.
•Quantum mechanics is the modern theory that describes the properties
of atoms, molecules, and subatomic particles.
•Quantum mechanics provides a set of rules that describe the way in
which electrons are organized within an atom.
•We will examine the periodicity of the chemical properties of the
elements by describing the electronic arrangement of the hydrogen atom
and then by moving one atom at a time through the periodic table.
•Electrons are organized within an atom into shells, subshells, and
•The relationship between these levels of organization depends upon the
energy state of the atom, which in turn depends upon a quantity called
the principle quantum number.
•Shell: Identified by a principal quantum number (1, 2, 3 . . . n) that
specifies the energy level, or energy state, of the shell. The higher
the quantum number, the greater the energy and the farther the
shell electrons are from the nucleus.
•Subshells: Locations within a shell, identified by the lowercase
letters s, p, d, and f.
•Orbital: The region of space within a subshell that has the highest
probability of containing an electron.
The region of space in which an electron is most likely to be found is
called an atomic orbital, which is best viewed as a cloud surrounding
the nucleus.
The size and shape of an orbital depend upon the atom’s energy state:
the greater the energy, the larger and more complex the orbitals.
Each type of orbital, s, p, d, f, retains its
shape as the principle quantum number
increases, but increases in volume.
The sizes and shapes of the atomic orbitals are responsible for the
properties of molecules constructed from the atoms (bond length, bond
angles, etc.)
Combining these two tables, the orbitals
for the first 3 principle quantum
numbers are:
2s 2p1 2p2 2p3
3s 3p1 3p2 3p3 3d1 3d2 3d3 3d4 3d5
•An orbital is a “state” or location in which electrons can exist in an
•Only two electrons can occupy any one orbital on the same atom. In
order for this to happen the two electrons must have opposite “spins”.
•When two electrons have opposite spins, they are said to be spinpaired.
•The concept of electron spin results from the fact that they behave
like a spinning top which can rotate either in a clockwise or
counterclockwise direction.
•Since each orbital can hold only two electrons, it is now possible to
calculate how many electrons can be held by each subshell:
•The rules for constructing a picture of the arrangement of the electrons
about an atom are called the Aufbau rules.
•Using these rules, we describe the atomic structure of hydrogen and
then systematically describe the atomic structure of each successive
atom by adding one proton and one electron to the picture of the
previous atom.
Once a particular orbital is filled, the next electron must be placed in an
empty orbital of next highest energy. No intermediate energy levels are
This property means that the energies of the electrons in an atom are
Fig. 2.13
The electronic configurations of the first four elements illustrate the
Aufbau process:
This configurations could also be written:
As p orbitals are filled, Hund’s rule must be applied, as illustrated for
carbon and oxygen:
These configurations can be written:
Note that the three p orbitals may be written in this notation using a
single letter p with a superscript from 1 to 6.
•The electronic configuration of the larger elements can become very unwieldy to
•We can think of the electronic configuration of an element as that of the
preceding inert gas plus any additional electrons left over.
•Li: 1s22s1 becomes [He]2s1; Mg: 1s22s22p63s2 becomes [Ne]3s2.
Notice that the outer electronic configuration (in blue) is the same for
each column of elements, except for the principle quantum number.
This is the reason that the elements in each of the columns shown have
similar chemical properties.
The electronic configuration of a carbon atom that contains 6 electrons
can be written:
1s2 2s2 2px1 2py1
According to Hund’s rule, each of the three p orbitals, px, py, and pz must
contain one electron (each with the same spin) before any of the p
orbitals can be filled.
When an alkali metal atom (sodium) reacts with a halogen atom (Cl), an
electron is transferred from the sodium atom to the chlorine atom. The
sodium atom becomes a positively charged cation, and the chlorine atom
a negatively charged anion.
Examining the electronic configurations of Na+ and Cl-, it can be seen
that each is identical to that of an inert gas: Na+ to Ne and Cl- to Ar.
The simple chemistry of the metals and nonmetals is dominated by this idea,
which is called the Octet Rule. Each inert gas contains the electronic
configuration: ns2np6, where n is the principle quantum number.
The chemistry of the elements depends only upon the electrons in the
outermost shell called the valence shell. These electrons are called the
valence electrons.
Individual atoms are often represented by their chemical symbol,
surrounded by dots representing the valence electrons for that atom.
These representations are called Lewis symbols.
With the exception of helium, the number of valence electrons for an element is
the same as its group number.
The reaction between an alkali metal and a halogen can also be written using
Lewis symbols. Note that whenever an electron is lost by one atom, it is always
gained by a second atom.
Chapter 2 Summary
Dalton’s Atomic Theory
• All matter consists of particles called atoms, which are indestructible.
• The atoms of any one element are chemically identical.
• Atoms of different elements are distinguished from one another
because they have different masses.
• Compounds consist of combinations of atoms of different elements. In
chemical reactions, atoms trade partners to form new compounds.
• The theory explained the conservation of mass in chemical reactions
and the constant composition of matter.
Chapter 2 Summary
Atomic Masses
• The relative mass of an element is called its atomic mass.
• Atomic masses are calculated in relation to the mass of the most
common form of the element carbon, exactly 12 atomic mass units
Chapter 2 Summary
The Structure of Atoms
• Atoms contain three kinds of particles: protons, electrons, and
• The protons and neutrons are located together in the nucleus.
• In an electrically neutral atom, the number of protons is balanced by
the same number of electrons.
• An atom bearing a net electrical charge is called an ion. A cation is a
positively charged ion, and an anion is a negatively charged ion.
Chapter 2 Summary
• Naturally occurring elements consist of isotopes.
• Isotopes of a given element have the same atomic number but have
different numbers of neutrons.
• An isotope’s mass number is the sum of the numbers of protons and
neutrons in its nucleus.
Chapter 2 Summary
The Periodic Table
• The properties of the elements repeat periodically if the elements are
arranged in order of increasing atomic number.
• In the periodic table, elements of similar chemical properties are
aligned in vertical columns called groups.
• Horizontal arrangements, called periods, end on the right with a noble
• Each of the second and third periods contains eight elements. But the
fourth period includes a new group of ten transition elements.
• Elements are designated as either main-group or transition elements.
• The metals occupy most of the left-hand side of the periodic table, and
the nonmetals occupy the extreme right-hand side of the table.
• Elements called metalloids, or semimetals, have properties between
those of metals and nonmetals and lie between the two larger classes.
Chapter 2 Summary
Electron Organization Within the Atom
• Electrons in atoms are confined to a series of shells around the
• The farther away the electrons are from the nucleus, the greater their
• Electron energy states are characterized by an integer called the
principal quantum number.
• Electrons are contained in shells that consist of subshells, and
electrons in these subshells reside in orbitals.
• Only two electrons are allowed in any orbital, and, to do that, they
must have opposite spins—they must be spin-paired.
• The atom’s orbitals define the probability of finding an electron in a
given region of space around the nucleus.
• With an increase in the number of shells, the number of subshells and
the complexity of their shapes increase.
Chapter 2 Summary
Atomic Structure, Periodicity, and Chemical Reactivity
• The symbolic notation that indicates the distribution of electrons in an
atom is called its electron configuration.
• An alternative method of representation, Lewis symbols, emphasizes
the number of electrons in an atom’s valence shell.
• The valence electrons dictate the chemical reactivity of each element.
• The octet rule summarizes the tendency of atoms to attain noble-gas
outer-shell electron configurations through chemical reactions.