This document provides a summary of key concepts from 11 chapters of a chemistry textbook. It outlines fundamental chemistry topics like matter, elements, compounds, chemical bonds and reactions. Specific concepts covered include the nature of chemistry, atomic theory, the periodic table, stoichiometry, gases and more. Formulas, equations and models are explained that are important for understanding the structure and behavior of matter at the molecular level.

1. Breaking Through
Chemistry
By: John Mark B. Sebolino

2. Chapter 1. The Nature of Chemistry
Key Concepts
 Chemistry - is the study of matter and the
changes it undergoes. It is sometimes called the
central science because it overlaps with many
other sciences.
 Technology - is the application of science. It has
improved the quality of human life.
 SI units - are used to express physical quantities
in all sciences. Metric prefixes are used to make
units smaller or larger.

3.  Precision - is how close several measurements
are to the same value.
 Accuracy - tells how close a measurement is to
the true or accepted digit.
 Significant - figures include both the certain
digits and the estimated digit.
 Scientific notation - is used to write very small
or very large numbers.

4.  Dimensional analysis - is the technique that
uses conversion factors. The guide to ensure that
conversion factor are properly formulated is the
cancellation of units.

5. Chapter 2. Matter: Its Composition and
Organization. Key Concepts
 Matter - is anything that has mass and volume.
Properties of matter differ for solids, liquids, and
gases.
 A pure substance is either an element or a
compound. An element is a substance that
cannot be broken down to simpler substances. A
compound is formed when two or more elements
combine in a chemical change.

6.  A change in the properties of substance without a
change in composition is a physical change. If
there is a change in the composition of a
substance, a chemical change has occurred.
Chemical changes produce matter with new
properties.
 The physical combination of two or more
substances is a mixture. A mixture has a variable
composition. It may be heterogeneous or
homogeneous. Heterogeneous mixtures (coarse
mixtures, suspensions, and colloids) do not have
uniform properties throughout, while
homogeneous mixtures (solutions) have uniform

7.  Solutions may be gases, liquids, or solids.
 The components of a mixture can be separate by
physical methods.
 Colloids - are mixtures of two or more
solids, liquids, or gases whose particles are
bigger than the particles of a solution but smaller
than those of a suspension.
 Tyndall effect, Brownian
movement, adsorption, and electrical charge are
the properties of colloids.

8.  Colloids are prepared and purified by
condensation and dispersion methods.
 Condensation - is the process of combining
molecules to form colloidal particles.
 Dispersion - is the process of breaking down
large particles to colloidal size.
 Energy - is the capacity to do work or to transfer
heat. It is involved whenever matter undergoes a
change.

9. Chapter 3. Atomic Theory.
Key Concepts
 Over 2400 years ago, the concept of the atom
was proposed by Greek philosophers.
 In the early 19th century, Daltons proposed the
atomic theory. This theory is related to the three
fundamental laws of matter.
 (1) The total mass of the reactants and products
are constant during a chemical reaction (law of
conservation of mass).

10.  (2) Any sample of compound, has elements in
the same proportion (law of definite
composition).
 (3) In different compounds of the same
elements, the mass of one element that combines
with a fixed mass of the other can be expressed
as a ratio of small whole numbers (law of
multiple proportions).

11.  Thomson’s experiment on the behavior of
cathode rays in magnetic and electric field led to
the discovery of the electron and the
measurement of its charge to mass ratio.
 Millikan’s oil drop experiment measured the
charge of the electron.
 Becquerel and the Curies discovered
radioactivity.
 Rutherford’s studies on alpha rays led to the
discovery of nucleus.

12.  Atoms have a nucleus that contains protons and
neutrons. Electrons move in the space around the
nucleus.
 Elements can be classified by atomic number or
the number of protons in the nucleus of an atom.
 All atoms of a given element have the same
atomic number. The mass number of an atom is
the number of protons and neutrons.
 All atoms of the same element that differ in mass
number are known as isotopes.

13. Chapter 4. Electronic Configuration. Key
Concepts
 The properties of visible light and other forms of
electromagnetic radiation led to the electronic
structure of atoms.
 Max Planck proposed that energy is absorbed
and emitted in discrete amounts or individual
packets called quanta (plural for quantum).
 Albert Einstein used Planck’s theory to explain
the photoelectric effect. He proposed that light
consists of quanta of energy which behave like
tiny particles of light. He called these energy
quanta photons.

14.  The concept of quantized electrons grew from the
study of line spectra of atoms. A line spectrum
consists of quanta of energy which can be used
like fingerprints to identify the element.
 Niels Bohr used the line spectra to explain
specific energy levels within the atom. He
proposed the planetary model of the atom.
 Louis de Broglie discovered the wave nature of
matter which initiated the development of a new
mathematical description of electron
configuration.

15.  Heisenberg’s uncertainty principle explained the
impossibility of simultaneously measuring the
momentum and location of an electron.
 Erwin Schrodinger devised the quantum
mechanical model of the atom which described
electrons as waves that exist in quantized energy
levels.
 The regions in space around the nucleus where
electrons are most likely to be found are called
orbitals. These orbitals have various shapes and
are labeled s, p, d, and f. Each principal energy
level or shell consists of these orbitals.

16.  The manner in which electrons are arranged
around the nucleus of an atom is called electron
configuration.
 The Aufbau principle, the Puali exclusion
principle, and the Hund’s rule are applied in
writing electron configurations.
 The Aufbau principle tells the sequence in which
orbitals are filled.
 The Pauli exclusion principle states that a
maximum if only two electrons can occupy an
orbital.
 Hund’s rule explains that electrons pair up only

17. Chapter 5. The Periodic Table.
Key Concepts
 Different periodic table were developed by
Dobereiner, Newlands, Mendeleev, and Meyer.
The periodic table was based on similarities in
properties and reactivities of elements in the
increasing order of their atomic mass.
 Discrepancies in these periodic tables were
resolved when Moseley established that each
element has a unique atomic number and
showed that elements should be arranged
according to their increasing atomic number.

18.  The periodic table is organized into 18 groups or
families and 7 periods or rows. The groups are
organized further into s, p, d, and f blocks based
on how valence electrons fill each sublevel.
Elements in a group have similar properties
because they have the same valence electrons.
 Atomic radius decreases from left to right across
a period because the positive charge of the
atoms increases, which attracts electrons more
strongly.

19.  Atomic radius increases down a group because
the electrons of the atoms fill more energy levels.
 Ionization energy - is the energy absorbed to
remove an electron to form a positive ion.
 Electron affinity - is the energy when an atom
gains an electron forming a negative ion.
 Electronegativity is the attraction of an atom for
electrons in a chemical bond.

20.  The trends for ionization energy, electron affinity,
and electronegativity ate the same. They increase
from left to right of the periodic table and
decrease down a period.
 Metals are found on the left side of the periodic
table. Nonmetals are found on the upper right
side of the periodic table. Metalloids have some
properties of metals and nonmetals.

21. Chapter 6. Chemical Bonds.
Key Concepts
 Chemical bonds are classified into three groups: ions
of opposite charges; covalent bonds, which result
from the sharing of electrons by two atoms; and
metallic bond, which are the attractions among
positively charged ions for delocalized electrons.
 These bonds involve the valence electrons with the
tendency of atoms follow the octet rule. This can be
represented by electron – dot symbols or Lewis
symbols.
 Resonance structures are used when a simple Lewis
structure is not adequate to represent a particular
molecule or ion (specie). Some covalent molecules
formed from atoms of the representation groups
1, 2, and 3 lack octet configurations while atoms from
5, 6, and 7 form expanded octet configurations.

22.  A polar covalent bond is formed when electrons
are not shared equally between two atoms.
 Electronegativity difference of bonded atoms
determines the kind of bond formed between the
atoms.
 The sharing of one pair of electrons produces a
single bond, the sharing of two pairs, a double
bond, and three pairs, a triple bond. Double and
triple bonds are also called multiple bonds.

23. Chapter 7. Molecular Geometry.
Key Concepts
 The shapes of small molecules can de explained
in terms of the VSEPR model which states that
electron pairs arrange themselves as far apart as
possible to minimize electrostatic repulsion.
 The geometry of molecules is determined by the
arrangement of bonding pairs and lone pairs.
 The five common shapes of small molecules are
linear, trigonal planar, tetrahedral, trigonal
bipyramid, and ictahedral.

24.  The electron pair cloud repulsion model suggests that
the denser the electron clouds, the greater the
repulsive force. The order from greatest to least
repulsive force is that triple bond > double bond >
lone pair > single bond (≡>═>1.p.>─).
 Molecules that contain polar bonds (bond dipoles)
may be polar or nonpolar molecules, depending on
the shape of the molecules. The properties of polar
molecules (dipole) are different from those of
nonpolar molecules.
 Valence bond theory - is an extension of the Lewis
covalent bond. In this theory, bonds are formed when
neighboring atoms overlap and the potential energy of
the system decreases. The greater the overlap, the
stronger the bond formed.

25.  Shapes of molecules are also described in terms
of hybrid orbitals. The process of hybridization
involves the promotion of electron to empty
orbital(s) and mixing of the orbitals to form
equivalent numbers of hybrid orbitals. Hybrid
orbitals can overlap with orbitals of other atoms to
make bonds. Or they can accommodate lone
pairs.
 Covalent bonds that overlap end to end along the
line connecting the atoms are called sigma (σ)
bonds. When p orbitals overlap on a side to side
orientation perpendicular to the line connecting
the atoms, these are called pi (π) bonds.

26. Chapter 8. Chemical Names and
Formulas.
Key Concepts
 The charges or oxidation numbers of the ions of
representative elements are determined by their
position in the periodic table.
 Most transition metals have more that one common
ionic or oxidation numbers. A polyatomic ion is a
group of atoms that behaves as an ion – ide. If
cations have more than one ionic charge, a Roman
numeral is used in the name.
 Ternary ionic compounds contain at least one
polyatomic ion. The names of these compounds end
in – ite or – ate.
 Binary molecular compounds are composed of two
nonmetallic elements. Prefixes are used to indicate
the number of atoms each element that are present in

27.  Binary acid are compounds that contain hydrogen
and nonmetal ions. They are named by using the
prefix hydro followed by the name of the anion
ending in – ic acid.
 Ternary acid contain hydrogen and polyatomic
ions. They are named by using the name of the
polyatomic ion ending in – ic or - ous acid
 Based are compounds containing a metal ion and
hydroxide ion(OH‾). Bases are named by writing
the name of the cation followed by hydroxide.
 Salts are named by using the name of the cation
followed by the name of the anion.

28. Chapter 9. Chemical Reactions.
Key Concept
 Chemical reactions are represented by chemical
equations.
 The substances that undergo chemical changes
are the reactants and the substances formed are
the products.
 Chemical equations must be balanced to be
consistent with the law of conservation of mass.
In balancing an equation, appropriate coefficients
are placed before the formulas of the reactants
and products so that the same number of atoms
of each element appears on each side of the
equation.

29.  The state of a substance in an equation is
detonated by (s), (1), and (g) for solid, liquid, and
gas, respectively. A substance dissolved in water
is denoted by (aq) for aqueous. If heat, light, or
electricity is used to initiate the reaction, its
process or symbol is written above the arrow. If a
catalyst is used to increase the speed of
reaction, its formula or symbol is also written
above the arrow.
 In a combination reaction, two or more
elements or compounds combine to produce a
single product.

30.  In a decomposition reaction, a single
compound is broken into two or more simpler
substances.
 In a single replacement reaction, a more
chemically active element displaces a substance
below it in the activity series.
 A double replacement reaction involves the
exchange of cations and anions between two
compounds. Replacement reactions can be
written as net ionic equations.
 In a combustion reaction, oxygen is always one
of the reactants.

31. Chapter 10. Stoichiometry.
Key Concept
 A mole is the amount of substance that contains
6.02 ×1023 particles or species.
 The representative particles of elements are the
atoms.
 Molecules are representative particles of
molecular compounds and diatomic elements.
 The representative particles for ionic compounds
are formula units. The mass of a mole of
atoms, molecules, or ions is its formula weight
expressed in grams called molar mass.
 A mole is defined in terms of the number of
particles in a substance or the mass in grams of
the substance. The mole can be used in
converting among different units.

32.  Percent composition of a compound is the
percent by mass of each element in a compound.
 Empirical formula is the simplest whole-number
ratio of atoms of elements in a compound. This
can be calculated from the percent composition of
a compound.
 Molecular formula shows the actual number of
atoms of each element in a compound. It may be
the same as or a multiple of an empirical formula.
 Stoichiometry is the study of the quantitative
relationship of individual compounds in chemical
reactions.

33.  The coefficients in a balanced equation represent
the relative number of moles of each substance.
Coefficients are used in establishing conversion
factors as mole ratios in solving stoichiometric
problems.
 The conversion factor relates the mole of a given
substance to the moles of the required
substance. Units such as grams and particles are
converted to moles when solving stoichiometric
problems.
 When reactants supplied are not in the exact
amounts required by the balanced equation, that
which is used up is the limiting reagent and that
which remains after the reaction is completed is

34.  The theoretical yield is the amount of product
obtained when all of the limiting reagent is used
up.
 The actual yield is the product formed when the
actual reaction is carried out.
 The percent yield is the ratio of the actual yield
to the theoretical yield expressed in percent.

35. Chapter 11. Gases
Key Concept
 The physical properties of gases are given by four
quantities:
 Pressure P
 Volume V
 Temperature T
 Amount of Gases n
 The behavior of gases can be explained by the kinetic
molecular theory.
 The standard temperature and pressure (STP) is 0°C
and 1 atm.
 Atmospheric pressure is the pressure exerted by the
gases (air) around us which is 1 atm or 760 mm HG.

36.  Boyle’s law states that the pressure and volume of a gas
are inversely proportional to its absolute temperature
(constant n and T).
 Charles law states that the volume of a gas is directly
proportional to its absolute temperature (constant n and P).
 Avogadro's law states that equal volumes of gases contain
the same number of particles (constant T and P).
 Ideal gas equation PV=nRT is a combination of the gas
laws.
 Daltons law states that the pressure of a mixture of gases
is the sum of the partial pressure of the component gases.
 Real gases behave like ideal gases in ordinary conditions
except at high pressure and low temperature.
 Lighter gases diffuse and effuse faster than heavier gases
do.

37. Chapter 12. Liquids and Solids
Key Concept
 At room temperature, substances with weak
intermolecular forces of attraction are gases;
those with moderate intermolecular forces are
liquids; and those with strong intermolecular
forces are solids.
 Intermolecular forces include ion-dipole
forces, dipole-dipole forces, London dispersion
forces and hydrogen bonds.
 Physicals properties of liquids and solids are
explained by the kinetic molecular theory.

38.  Liquids possess properties such as
viscosity, surface tension, capillarity
evaporation, boiling point, and critical
temperature and pressure.
 Heating curve is a plot of temperature versus
heat for phase changes.
 The properties of solids are explained based on
their nature and strength if intermolecular forces
of attraction.
 A phase diagram indicates the states or phases
of a substance under specific temperatures and
pressures.

39. Chapter 13. Solutions
Key Concept
 Solutions are homogeneous mixtures of two or more
substances in a single phase.
 A solutions is made of solute, the substance that
dissolves, and solvent, the substance in which the
solute is dissolved. A substance that dissolves in
another substance is soluble (miscible) and if it does
not, it is insoluble (immiscible).
 Solutions are either gaseous, liquids, or solid
solutions.
 In preparing dilute solutions form concentrated
solutions, the number of moles before dilutions is
equal to the number of moles after dilutions.

40.  Saturated solutions contains the maximum
amount of solute it can dissolve at a given
temperature.
 Unsaturated a solutions that contains less than
the maximum.
 Supersaturated a solution with more than the
maximum.
 Solubility is the extent to which a solute dissolves
in a given solvent.

41. Chapter 14. Chemical Kinetics
Key Concept
 Chemical kinetics is the study of rate and
sequence of steps by which chemical reactions
occur.
 The rate of a reaction is the measure of how
reactants turn into products.
 Collisions theory assumes that particles collide at
the proper orientation and with sufficient energy in
order to react.
 Activation energy is the minimum energy required
for a chemical reaction to occur and make the
reactant form an activated complex or transition
state.

42.  The factors that affect the rate at which a chemical
reaction proceed are nature of the
reactants, concentration of the reactants, temperature
at which reaction occurs.
 A rate law for a reaction describes the relationship
between the concentration of reactants and the
reaction rate.
 Most chemical reactions proceed through a series of
elementary steps. The series of steps called the
reaction mechanism.
 The slow reaction in a reaction mechanism called the
rate-determining step.

43. Chapter 15. Thermo chemistry
Key Concept
 Thermodynamics is the study of processes which
involve heat transfer and the performance of work.
 Thermochemistry is the study of this heat exchange
and work on chemical reactions.
 Energy + Energy = constant: law of conversation of
energy.
 3 types of system :
 Open
 Closed
 Isolated
 An open system allows the transfer of both energy
and matter into and out the system through a
boundary or wall.

44.  A closed system is only capable of transferring
energy through boundary.
 An isolated system is not capable of transferring
both energy and matter into and out of the system
through a boundary or wall.
 Heat is a transfer of energy between system and
surrounding due to temperature difference.

45. Chapter 16. Chemical Equilibrium
Key Concept
 Equilibrium is a state at which there is “balance of
forces”.
 3 types of equilibrium:
 Mechanical
 Thermal
 Chemical
 Chemical equilibrium is achieved when the rate of the
forward reaction is equal to the rate of the reverse
reaction and the amount of components remains
unchanged.

46.  Reversible reactions is an incomplete reactions.
The reaction is represented by using a double
headed arrow (═).
 Law of mass reaction states that the
compositions of a reaction mixture can vary
according to the quantities of components that
are present.

47. Chapter 17. Acids and Bases
Key Concept
 The operational definitions of acids and bases are
based on experimental results from the laboratory
which includes color change using dyes.
 Arrhenius acids is a neutral substance that
ionizes when it dissolves in water to give the H+
or hydrogen.
 Arrhenius base is a neutral substance that gives
the OH-, or hydroxide ion when dissolves in
water.

48.  Lewis defines an acids as species that can
accept a pair of electrons while a base is a
species that can donate a pair of electrons.
 The degree of ionization, not the concentration,
classifies an acid or a base as weak or strong.
 Compounds with more than one proton to give
are called polyprotic acids.

49. Chapter 18. Electrochemistry
Key Concept
 Electrochemistry is the branch of chemistry that deals
with electricity and its relation to chemical reactions.
 A chemical reactions were loss of electron(s)
is involved id called oxidation while reaction where
electron(s) is gained is called reduction.
 Redox reaction can be balanced by using the
oxidation number method or the ion electron method.

50.  Electrochemical cell, voltaic cell, or galvanic cell converts
chemical energy from spontaneous reaction to produce
electricity.
 Electrochemical cell is composed of the electrodes and
charge carriers.
 Anode is the electrode where oxidation occurs.
 Cathode is where reduction occurs or where electrons are
accepted.
 There 3 types of electrodes:
 Inert
 Metallic
 Membrane

51. Chapter 19. Nuclear Chemistry
Key Concept
 Many elements have at least one radioactivity isotope or
radioisotope. Elements with atomic numbers 83 or greater
are all radioactivity.
 Radioactivity decay of naturally occurring radioisotope
produces alpha particles, beta particles, and gamma
radiations.
 The half-life of a radioisotope is the time it takes for one-
half of a sample of the isotope decay.
 In artificial radioactivity or artificial transmutation, the
nucleus of an atom is bombarded with a particle or
radiation and changed into different nuclei.

52.  In balancing nuclear equation, the sum of the
mass numbers and atomic numbers of reactants
must be equal to the sum of the mass numbers
and atomic numbers of the product.
 The mass defect in a nucleus is due to the strong
forces of attraction that bind nucleons together.

53. Chapter 20. Organic Chemistry
Key Concept
 Organic compounds are basically made up of carbon
atoms bonded mostly to
hydrogen, oxygen, nitrogen, and sulfur.
 Organic chemistry the study of the carbon-based
compounds.
 Hydrocarbons are made up of carbons and
hydrogens.
 Alkanes also called saturated hydrocarbons, have an
sp3 hybridization, four sigma bonds with no pi bonds
that can be bound to H or C atoms.

54.  Alkenes are hydrocarbon containing a carbon-carbon
double bond.
 Alkynes are hydrocarbons containing a carbon-carbon
triple bond.
 Cycloalkanes are aliphatic cyclic (alicyclic)
compounds which have general ring structure
containing –CH-.
 A molecule can only be aromatic if it has the following
properties:
 (1) the molecule is planar and
 (2) has a monocyclic system of conjugation with a total
of (4n + 2) p electrons where n is an integer.