CHEM 121 Overview
Part 1
MATTER CLASSIFICATION SUMMARY
MIXTURES
• Heterogeneous mixture – components
making up the mixture separate out
• Sand in water
• Homogeneous mixtures are also called
solutions. No separation of components can
be detected.
• Sugar in coffee
• Soda drinks
PHYSICAL & CHEMICAL PROPERTIES AND CHANGES
• PHYSICAL PROPERTIES OF MATTER
• Can be observed or measured without attempting to change the
composition of the matter being observed.
• Examples: color, shape and mass
• PHYSICAL CHANGES OF MATTER
• Take place without a change in composition.
• Examples: freezing, melting, or evaporation of a substance (e.g.
water)
• CHEMICAL PROPERTIES OF MATTER
• Can be observed or measured only by attempting to change the
matter into new substances.
• Examples: flammability and the ability to react (e.g. when vinegar
and baking soda are mixed)
• CHEMICAL CHANGES OF MATTER
• Always accompanied by a change in composition.
• Examples: burning of paper and the fizzing of a mixture of vinegar
and baking soda
SUBATOMIC PARTICLES
• Nucleus contains protons are
neutrons.
• Protons carry a +1 electrical
charge and have a mass of 1
atomic mass unit (u).
• Neutrons carry no electrical
charge and have a mass of 1
atomic mass unit (u).
• Electrons are located outside the
nucleus of an atom. They carry a
-1 electrical charge and have a
mass of 1/1836 atomic mass unit
(u). They move rapidly around
the heavy nucleus.
SUBATOMIC PARTICLE CHARACTERISTICS
ATOMIC & MASS NUMBERS
• ATOMIC NUMBER OF AN ATOM (Z)
– The atomic number of an atom is equal to the number of
protons in the nucleus of the atom.
A
Z
E
• MASS NUMBER OF AN ATOM (A)
– The mass number of an atom is equal to the sum of the number
of protons & neutrons in the nucleus of the atom.
• Isotopes contain same number of protons but different
number of neurons
ATOMIC MASS UNIT (u or amu) and
Molecular Weight
• An atomic mass unit is a unit used to express the relative masses of atoms.
One atomic mass unit is equal to 1/12 the mass of a carbon-12 atom.
• A carbon-12 atom has a relative mass of 12 u.
• The atomic weight of an element is the relative mass of an average atom of
the element expressed in atomic mass units.
• According to the periodic table, the atomic weight of nitrogen atoms (N) is
14.0 u, and that of silicon atoms (Si) is 28.1 u.
• The relative mass of a molecule in atomic mass units is called the molecular
weight of the molecule.
• Because molecules are made up of atoms, the molecular weight of a molecule
is obtained by adding together the atomic weights of all the atoms in the
molecule.
MOLECULAR WEIGHT
• The relative mass of a molecule in atomic mass units is called the
molecular weight of the molecule.
• Because molecules are made up of atoms, the molecular weight of
a molecule is obtained by adding together the atomic weights of all
the atoms in the molecule.
• The formula for a molecule of water is
H2O. This means one molecule of water
contains two atoms of hydrogen, H, and
one atom of oxygen, O. The molecular
weight of water is then the sum of two
atomic weights of H and one atomic
weight of O:
• MW = 2(at. wt. H) + 1(at. wt. O)
• MW = 2(1.01 u) + 1(16.00 u) = 18.02 u
THE MOLE CONCEPT (continued)
• THE MOLE AND CHEMICAL CALCULATIONS
– The mole concept can be used to obtain factors that are useful in
chemical calculations involving both elements and compounds.
One mole quantities of six
metals; top row (left to
right): Cu beads (63.5 g),
Al foil (27.0 g), and Pb
shot (207.2 g); bottom
row (left to right): S
powder (32.1 g), Cr
chunks (52.0 g), and Mg
shavings (24.4 g).
One mole quantities of
four compounds: H2O
(18.0 g); small beaker
NaCl (58.4 g); large
beaker aspirin, C9H8O4,
(180.2 g); green (NiCl2 ·
6H2O) (237.7 g).
ATOMIC ORBITALS (continued)
• According to the quantum mechanical model, all types of atomic
orbitals can contain a maximum of two electrons.
• Thus, a single d orbital can contain a maximum of 2 electrons, and a
d subshell that contains five d orbitals can contain a maximum of
10 electrons.
ATOMIC ORBITALS
• The last descriptor of the location and energy of an electron moving
around a nucleus is the atomic orbital in which the electron is
located.
• Each subshell consists of one or more atomic orbitals, which are
specific volumes of space around the nucleus in which electrons of
the same energy move.
ATOMIC ORBITALS (continued)
•Atomic orbitals are designated by
the same number and letter used
to designate the subshell to which
they belong. Thus, an s orbital
located in a 2s subshell would be
called a 2s orbital.
•All s subshells consist of a single s
orbital.
•All p subshells consist of three p
orbitals.
•All d subshells consist of five d
orbitals.
•All f subshells consist of seven f
orbitals.
FILLING ORDER & PERIODIC TABLE
• Notice the order of subshell filling matches the order of the
subshell blocks on the periodic table, if the fill occurs in the order of
increasing atomic numbers.
ELEMENT CLASSIFICATION
VALENCE ELECTRONS
• Valence electrons are the electrons in the outermost shell.
• The valence electrons determine chemical behavior and bonding
behavior.
• The number of valence electrons is the same as the Roman numeral
group number.
• Examples: Calcium, Ca, is in group IIA. The number of valence
electrons is 2. Phosphorus, P, is in group VA. The number of
valence electrons is 5.
LEWIS STRUCTURES
• A representation of an atom or ion in which the elemental
symbol represents the atomic nucleus and all but the valenceshell electrons. The valence electrons are represented by dots
arranged around the elemental symbol.
NOBLE GAS CONFIGURATIONS
• An electronic configuration that is characterized by two
electrons in the valence shell of helium and eight electrons in
the valence shell of all other group VIIIA noble gases.
THE OCTET RULE
• According to the octet rule, atoms will gain or lose sufficient
electrons to achieve an outer electron arrangement identical to
that of a noble gas. This arrangement usually consists of eight
electrons in the valence shell.
• SIMPLE ION
– A simple ion is an atom that has acquired a net positive or
negative charge by losing or gaining one or more electrons.
BINARY IONIC COMPOUND FORMULAS
• Binary ionic compounds typically form when a metal and a
nonmetal react.
• The metal tends to lose one or more electrons and forms a
positive ion.
• The nonmetal tends to gain one or more electrons and
forms a negative ion.
• The symbol for the metal is given first in the formula.
• NaCl, BeS, BaBr2, Al2O3
COVALENT BONDS
• A covalent bond is a type of bond in which the octet rule is satisfied when atoms
share valence electrons. The shared electrons are counted in the octet of each
atom that shares them as illustrated below for fluorine, F2.
The sharing of electrons takes place when electron-containing orbitals of atoms
overlap. This is shown below for the formation of the H2 molecule.
COVALENT MOLECULE POLARITY
• The shared electrons of covalent bonds are not always shared
equally by the bonded atoms.
• Electrons of a covalent bond are attracted toward atoms of
highest electronegativity.
COVALENT MOLECULE POLARITY (continued)
• When the resulting partial charges are distributed
symmetrically in a molecule, the molecule is nonpolar. When
the partial charges are distributed nonsymmetrically, the
molecule is polar.
VSEPR THEORY
• The shapes of molecules or polyatomic ions can be predicted using a theory
called the valence-shell electron-pair repulsion theory, or VSEPR theory
(sometimes pronounced "vesper" theory).
• According to the VSEPR theory, electron pairs in the valence shell of an
atom will repel each other and get as far away from each other as possible.
• When the VSEPR theory is used, two rules are followed:
– Rule 1: All valence-shell electron pairs around the central atom are
considered to behave the same regardless of whether they are bonding or
nonbonding pairs.
– Rule 2: Double or triple bonds between surrounding atoms and the
central atom are treated like a single pair of electrons when shapes are
predicted.
Molecular Shapes
INTERPARTICLE FORCE SUMMARY
• Ionic and covalent bonds represent two of the forces that occur
between atomic-sized particles and hold the particles together to
form the matter familiar to us.
• Other forces also exist that hold the particles of some types of
matter together. These include:
• metallic bonding,
• dipolar forces,
• hydrogen bonding,
• dispersion forces.