Chemistry 112
Overview of
Chapters 1-4
Chapter 1 Highlights
 Chemistry is the study of matter, the
physical substance of all materials.
 The building blocks of matter are atoms,
which combine to form compounds.
 The different types of atoms are called
elements, which are arranged
systematically in the periodic table.
Chapter 1 Highlights (cont)



Atoms are composed of protons, neutrons,
and electrons.
All atoms of the same element contain the
same number of protons (and electrons)
but may vary in the number of neutrons.
The protons and neutrons are found inside
the tiny but dense nucleus, whereas the
electrons are found in orbitals outside the
nucleus.
Chapter 1 Highlights (cont)

The arrangement of electrons in the orbitals
is called the electronic configuration and
determines the chemistry of an atom.
Types of Matter
States of Matter
 Chemistry and Matter
 Physical Changes versus Chemical
Changes
 Physical changes involve changes in
appearance (i.e., changes in state such as
melting).
 Chemical changes result in new substances.
 The Building Blocks of Matter
 Atoms
 Smallest representative units of the elements.
 Compounds
 Different atoms linked together; e.g., H2O.
 The Building Blocks of Matter (cont)
 Dalton’s Atomic Theory
 All matter is composed of indivisible atoms.
 All atoms of one element are identical to each
other but different than the atoms of other
elements.
 Compounds are formed when atoms of different
elements combine in whole number ratios.
 Atoms are rearranged during chemical reactions
but atoms cannot be created or destroyed.
 The Periodic Table
 Used to organize the elements by recurring
chemical properties.
 Elements in the same vertical column of the
periodic table have similar chemical properties
and are said to be in the same group or
family.
The Periodic Table
 The Atom
 Components
 Positive protons, negative electrons, and
neutral neutrons
 Atomic Number
 The number of protons in an atom, which
determines what element it is
 Mass Number
 Number of protons + the number of
neutrons
 The Atom (cont)
 Isotopes
 Isotopes of the same element have the
same number of protons but differ in the
number of neutrons.
 Atomic Mass
 The atomic mass for each element on
the periodic table reflects the relative
abundance of each isotope in nature.
Isotopes
 Models of the Atom
 The Plum Pudding Model
 Electrons are embedded in a sphere of positive
charge.
 The Nuclear Model
 All of the positive charge is in a tiny central
nucleus with electrons outside the nucleus.
 This model was developed by Rutherford after his
landmark experiment.
The Rutherford Experiment
 Models of the Atom (continued)
 Bohr’s Solar System Model
 Electrons circle the nucleus in orbits, which are
also called energy levels.
 An electron can “jump” from a lower energy
level to a higher one upon absorbing energy,
creating an excited state.
 The concept of energy levels accounts for the
emission of distinct wavelengths of
electromagnetic radiation during flame tests.
The Solar System Model
Electromagnetic Radiation
 Models of the Atom (continued)
 The Modern Model
 Orbits are replaced with orbitals, volumes of
space where the electrons can be found.
 The arrangement of electrons in the orbitals is
the electronic configuration of an atom, which
determines the chemistry of an atom.
The Orbital Model:
Electronic Configurations
Chapter 2 Highlights
 Having eight valence electrons is
particularly desirable (“the octet rule”).
 Atoms form bonds with other atoms to
satisfy the octet rule.
 The two major types of chemical bonds are
ionic and covalent.
Chapter 2 Highlights (cont)
 Electronegativity is the ability to attract
shared electrons.
 The type of bond formed between two
atoms depends on their difference in
electronegativity.
 Ionic bonds form between atoms with a
large difference in electronegativity
(generally a metal and a nonmetal).
Chapter 2 Highlights (cont)
 Nonpolar covalent bonds form between
atoms with little difference in
electronegativity (generally two
nonmetals).
 Polar covalent bonds form between
atoms with intermediate difference in
electronegativity.
 There are many ways to depict
molecules.
 The Octet Rule
 Atoms with eight valence electrons are
particularly stable, an observation called
the octet rule.
 Atoms form bonds with other atoms to
achieve a valence octet.
Electronic
Configuration of Noble Gases
Types of Compounds
Lewis Dot Structures
 Ionic Bonds
 Ionic compounds result from the loss of
electrons by one atom (usually a metal)
and the gain of electrons by another
atom (usually a nonmetal).
 Ionic bonds arise from the attraction
between particles with opposite charges
(electrostatic forces); e.g., Na+ Cl-.
Ionic Compounds
 Covalent Bonds
 Covalent bonds are formed when two
atoms share one or more electron pairs.
 When two atoms share one pair of
electrons, the result is a single bond.
 Two shared pairs of electrons is a double
bond; three is a triple bond.
 Equal Sharing versus Unequal Sharing
 When two different kinds of atoms are
bonded, the electrons are usually shared
unequally.
 When a bond exists between two identical
kinds of atoms, the electrons are shared
equally.
 An atom with greater electronegativity has a
greater ability to attract shared electrons.
Electronegativity
Polar vs. Nonpolar Bonds
 Representing Structures
 In a structural formula, atoms are represented by
chemical symbols, and bonds are represented by
lines.
 In a line drawing, any point where lines connect
or terminate is understood to be a carbon atom
with sufficient bonded hydrogen atoms to achieve
the four bonds necessary for carbon.
Drawing Molecules
Chapter 3 Highlights
 Reaction equations have with the initial
materials (reactants) on the left, followed
by a reaction arrow pointing from left to
right, and the final materials (products) on
the right.
 A balanced equation has the same number
and kinds of atoms on both sides of the
equation.
Chapter 3 Highlights



The relationship between the amounts of
reactants and products is the stoichiometry,
which comes from a balanced reaction
equation.
The SI unit for measuring atoms and
molecules is the mole.
In an oxidation-reduction reaction,
electrons are transferred from one material
(the substance that is oxidized) to another
material (the substance that is reduced).
 Balanced Reaction Equations
 Writing a Chemical Reaction
 The starting materials, the reactants, are
written on the left.
 The materials that are produced, the products,
are written on the right.
 Reactants are separated from products by a
horizontal arrow pointing from left to right.
Na + Cl
Reactants
NaCl
Product
 Balanced Reaction Equations (cont)
 Balancing the Equation
 The law of conservation of matter states that
matter can neither be created nor destroyed in
a chemical reaction.
 The number and kind of atoms on the lefthand side of an equation must be equal to the
number and kind of atoms on the right.
H2 + O2
H2O
Incorrect
2 H2 + O2
2 H2O Correct
 Balanced Reaction Equations (cont)
 Stoichiometry
 The stoichiometry of a chemical reaction is
the relationship between the number of
molecules of the reactants and products in
the balanced reaction equation.
 A reactant present in insufficient amounts is
the limiting reagent.
The Mole





The mole is the SI unit of measure to describe
the amount of matter that is present.
One mole is equal to 6.02 x 1023 particles
(Avogadro’s number).
One mole of an element has a mass that is
equal to the atomic mass of that element in
grams.
One mole of a compound has a mass that is
equal to the molecular/formula mass of that
compound in grams.
The Mole
 Stoichiometry Calculations
 The units of molar mass are grams/mole.
 Moles x molar mass = mass.
 Example: 2.0 mol CO2 x 44 g/mol = 88 g CO2
 Mass/molar mass= moles.
 Example: 132 g CO2 / 44 g/mol = 3.0 mol CO2
 Stoichiometry Calculations
 The expected mass of a product or
reactant can be calculated for any reaction
by using the balanced equation and the
molar mass.
 Oxidation-Reduction Reactions
 Defined
 Oxidation-reduction (“redox”) reactions
involve the transfer of electrons from one
substance to another.
 Oxidized substances lose electrons and
reduced substances gain electrons.
Oxidation-Reduction
 Oxidation-Reduction Reactions (cont)
 The Chemistry of Batteries
 Combining a readily oxidized substance with an
easily reduced substance can create a battery.
 The oxidized material is the anode and the
reduced material is the cathode of the battery.
Batteries
Chapter 4 Highlights
 Intermolecular forces hold the molecules of a
material together.
 Stronger intermolecular forces lead to higher
melting and boiling temperatures.
 The relative strengths of intermolecular forces
generally follow the trend:
hydrogen bonds > dipole-dipole interactions > London
forces
Chapter 4 Highlights (cont)
 Like dissolves like. That is, polar solutes
dissolve in polar solvents.
 Acids are proton (H+) donors; bases are proton
acceptors that produce OH- in solution.
 The pH measures the acidity of a solution: pH
< 7.0 is acidic; pH > 7.0 is basic; pH = 7.0 is
neutral.
 Acids react with bases in neutralization
reactions.
 States of Matter
 Review of Types of Bonds
1.Chemical bonds (intramolecular forces) hold
atoms together.
2.The three types of chemical bonds are ionic,
polar covalent, and nonpolar covalent.
3.Intermolecular forces hold molecules together.
Review of Types of Bonds
Chapter Outline
States of Matter (cont)


Particle Cohesion Determines Physical State


In general, the relative strengths of intermolecular
forces follows the trend:
gases < liquids < solids
Changes of State


Adding energy breaks intermolecular forces and
causes molecules to change their state.
The stronger the intermolecular forces of a
compound, the higher are the melting and boiling
points.
Changes of State
 Types of Intermolecular Forces
within Pure Substances
 London dispersion forces
 A temporary dipole in one molecule can
induce a dipole in a neighboring molecule.
 The negative end of one temporary dipole
can attract the positive end of an induced
dipole; these attractions are called London
dispersion forces.
 London forces tend to be fairly weak.
London Dispersion Forces
 Types of Intermolecular Forces within
Pure Substances (cont)
 Dipole-dipole interactions
 Dipole-dipole interactions exist between
molecules with polar covalent bonds.
 Dipole-dipole interactions are typically stronger
than London dispersion forces.
Dipole-Dipole Interactions
 Types of Intermolecular Forces
within Pure Substances (cont)
 Hydrogen Bonds
 Hydrogen bonds are a special type of
dipole-dipole interaction.
 Hydrogen bonds can occur when H is
bonded to one of the highly
electronegative atoms N, O, or F. An
example is H2O.
 Hydrogen bonds are typically quite strong.
Hydrogen Bonds in Water
Mixtures
 Forming Solutions
 Like dissolves like
 Ionic solutes often dissolve in polar solvents;
e.g., NaCl dissolves in H2O.
 Polar solutes generally dissolve in polar
solvents; e.g., NH3 in H2O.
 Nonpolar solutes generally do not dissolve well
in polar solvents; e.g., oil in H2O.
NaCl Dissolving in H2O
 Emulsions
 Emulsifying agents are molecules that
contain a polar portion and a nonpolar
region.
 Soap is an example of an emulsifying
agent that can form a suspension of a
nonpolar material in a polar solvent (an
“emulsion”).
Emulsification with Soap
Measuring Amounts in Solution


Solubility


The maximum amount of a solute that dissolves
in a solvent
Molarity


The amount of a solute dissolved in a solvent is
its concentration.
Concentration is often measured in moles/liter,
also called molarity (M).
 Acid-Base Chemistry
 Definitions of Acids and Bases
 Acids turn litmus paper red; bases turn litmus
paper blue.
 Acids produce H+ in solution; bases produce
OH- in solution.
 Acids are proton donors; bases are proton
acceptors.
 Acid-Base Chemistry (cont)
 The pH Scale: a measure of acidity
 Acid-Base Chemistry (cont)
 Acid-Base Indicators
 Molecular sensors of H+.
H+
 Acid-Base Chemistry (cont)
 Neutralization Reactions: equal molar
amounts of an acid and a base react to
form a neutral solution.
HCl + NaOH
NaCl + H2O
 Acid-Base Chemistry (cont)
 Buffers: contain a weak acid and its
conjugate base, which react with added H+
or OH- to prevent pH changes.
HA
H+ + A-
Adding acid: H+ reacts with A- to make more HA
Adding base: OH- reacts with HA to make more A- and H2O