UPCAT REVIEWER 2017: Chemistry
Chemistry
- study of composition of matter and the changes that it undergoes. It is the study of composition,
behavior, structure and properties of matter, as well as the changes it undergoes during chemical
reactions.

1)
2)
3)
4)
5)
Scientific Method
Identify the Problem
Formulate hypothesis
Test hypothesis
Hypothesis is rejected/supported
Generalizations
Matter- anything that occupies space and matter
Law of conservation of matter states that matter is always conserved. This statement means that the
total amount of matter in the universe remains constant. Matter is neither created nor destroyed. It is only
changed in form.
Law of conservation of energy states that energy is always conserved. This statement means that the
total amount of energy in the universe remains the same. Energy is neither created nor destroyed. It is
only changed in form.

Classifications of Matter
Mixtures- contain more than one kind of material

Homogeneous mixtures- the composition of the mixture is the same all throughout (e.g. solutions)

Heterogeneous mixtures- the composition of the mixture is not uniform (e.g. sand grains and iron
fillings)
Pure substances- has a definite (constant) composition and distinct properties

Compound- a substance composed of two atoms or more elements chemically united in fixed
proportions

Elements- a substance that cannot be separated into simpler substances
 Physical and Chemical Properties of Matter
 Physical Properties- depends on the substance itself
Two groups:
a. extensive properties- depend upon the amount of matter present
(e.g. mass, length, and volume)
b. intensive properties- do not depend upon the amount of matter present
(e.g. density, malleability, ductility, and conductivity, color, crystalline shape, melting point,
boiling point and refractive index- ability to bend light)
Physical Changes- changes on the physical appearance of the matter; does not alter the chemical
character of a matter
(e.g. Distillation is a change-of-state operation. It is used to separate substances with different
boiling points.)
Changes of state- changes in the physical state of a matter (e.g. liquid changes to a gas, or vice versa)
Specific solubility- amount of solute that dissolves
Precipitate- An insoluble substance which forms from a solution
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 Chemical Properties- properties of matter that depend upon the action of substances in the presence
of other substances
Chemical Change- changes that occur in substances whenever one or more new substances with different
properties are formed (e.g. burning, digesting, and fermenting)
The separation of compounds into their component elements always requires a chemical change. Mixtures
can always be separated by physical means. However, it is sometimes more convenient to separate
mixtures by a chemical change.
 Elements:
Chemical symbols- may be used in place of the name of the element; may consist of either one or two
letters, usually two (e.g. S for sulfur and Se for selenium)
The first letter of a symbol is always capitalized.
 Compounds:
Chemical formula- represents compounds; indicate the relative number of atoms of each element present
in a compound (e.g. The compound ammonia has a chemical formula of NH3. The elements forming
ammonia are nitrogen and hydrogen. Notice that one atom of nitrogen will react to three atoms of
hydrogen to form one molecule of ammonia.)
Organic compounds- compounds that contain carbon (e.g. acetic acid: CH3COOH)
Oxidation number- represents apparent charge on an atom
The algebraic sum of the oxidation numbers of the elements of a compound is zero.
Ion- charged atom or charged group of atoms
Polyatomic ion- an ion made up of more than one atom
Diatomic molecule- consist of two atoms per molecule (e.g. chlorine gas, Cl2)
 Naming compounds:
Binary compounds- compounds that contain only two elements
To name a binary compound, first write the name of the element having a positive charge. Then
add the name of the negative element. The name of the negative element must also be modified to end in
–ide. (e.g. The compound aluminum (Al3+) and nitrogen (N3-), with the formula AlN, is name aluminum
nitride.)
Elements that have more than one possible charge:
1. Add a prefix to the name of each element. The prefix will indicate the number of atoms of that element
in a molecule of the substance being named. (e.g.mono- for one atom, di- for two atoms, ect.)
2. Write the oxidation number of the element having positive charge after the name of that element.
Roman numerals in parentheses are used. (e.g. For N2O, its name using Roman numerals is nitrogen (I)
oxide while dinitrogen monoxide when prefixes is used.)
A few negative ions with names ending in –ide do not form inary compounds. (e.g. OH- (hydroxide), NH2(amide), N2H3- (hyrdazide), and CN- (cyanide).)
Empirical formula- indicates the simplest whole- number ratio of atoms in a formula unit
Molecular formula- describes a molecule; shows the actual number of each kind of atom in one molecule
of a compound; always a whole- number multiple of the empirical formula
Formula mass- for ionic compound, the sum of the atomic masses of the atoms in the formula
Empirical Formula and Molecular Formula from Mass Percent Composition
Steps:
1. Determine the mass of each element in a 100g sample.
2. Convert each of the masses to mole.
3. Write a tentative formula based on the number of moles.
4. Divide each of the subscripts of the tentative formula by the smallest number of moles.
5. Round off any subscripts from Step 4 that differs slightly from whole numbers. This gives the empirical
formula.
*in case of values not close to whole numbers, convert the value to whole numbers by multiplying
the value to a certain factor. Important: multiply all subscripts by the SAME factor
6. Calculate empirical formula mass.
7. Divide the given formula mass by the calculated formula mass.
8. To get molecular formula, multiply subscripts in empirical formula by the number you get in Step 7.
 Reactions:
Chemical reaction- process by which one or more substances are changed into one or more new
substances
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UPCAT REVIEWER 2017: Chemistry
There are two factors which seem to drive a reaction, causing it to occur.
 All systems strive to release energy; in other words, all chemical processes would prefer to undergo an
exothermic process.
 All systems strive to increase entropy; in other words, all chemical processes would prefer to undergo a
change that leads to a state of greater disorder.
 Chemical equations- used to represent chemical changes
Consists:
a. Reactants- starting substances
b. Products- resulting substances
Balancing equation
Steps:
1. Determine exactly what the reactants and the products are.
2. Assemble the parts of the chemical equation. Reactants are written on the left side of chemical
equations; products are indicated on the right. The symbols and formulas must be correct.
3. Balance the equation. Balancing means showing an equal number of atoms for each element on both
sides of the equation. In balancing an equation, change only the coefficients. Never change the subscripts.
E.g. C3H8 + 5O2 3CO2 + 4H2O
Reactants: propane- C3H8; oxygen- O2
Products: carbon dioxide- CO2; water- H2O
Number of C on the left hand side= Number of C on the right hand side= 3
Number of H on the left hand side= Number of H on the right hand side= 8
Number of O on the left hand side= Number of O on the left hand side= 10
Classifying chemical change:
1. Single displacement- one element displaces another in a compound
(e.g. 3Li + CmF3 3LiF + Cm –lithium displaces curium from curium (III) fluoride)
General form: element + compound  element + compound
2. Double displacement- positive and negative portions of two compounds are interchanged
(e.g. PbCl2 + Li2SO4 2LiCl + PbSO4)
General form: compound + compound compound + compound
3. Decomposition- substances will break up into simpler substances when energy is supplied; energy may
be supplied in the form of heat, light, mechanical shock, or electricity
(e.g. CdCO3CdO + CO2)
General form: compound  two ore more substances
4. Synthesis- two or more substances combine to form one new substance
(e.g. NH3 + HCl NH4Cl)
General form: element or compound + element or compound compound
 Energy and Chemical change
a. Endothermic- heat energy must be supplied in the reaction
b. Exothermic- heat energy is given off in the reaction
Activation energy- minimum amount of energy required to start a chemical reaction
 Atoms
Atomic Structure
Democritus proposed the earliest recorded atomic theory. Modern atomic theory dates from
John Dalton’s hypothesis. His idea made use of the law of conservation of mass and the law of definite
proportions.
Law of Definite Proportion states that specific substances always contain elements in the same ration by
mass.
Law of Multiple Proportions states that the ratio of masses of one element which combined with a
constant mass of the other element can be expressed in small whole numbers.
Avogadro stated that equal volumes of gases under the same conditions contain the same number
of molecules.
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UPCAT REVIEWER 2017: Chemistry
Atoms-the basic unit of an element that can enter into chemical combination
Electron- a subatomic particle representing the unit of negative charge
Proton- a positive particle found in the nuclei and having a mass of approximately one atomic mass unit
Neutron-a neutral particle found in the nucleus of an atom and having a mass of approximately one
atomic mass unit
Nucleon- a particle found in the nucleus of an atom. A proton or neutron.
Positron- a subatomic particle identical to an electron except possessing a positive charge. The antiparticle
of the electron.
Atomic number (Z)- the number of protons in the nucleus
Mass number- sum of the nucleons
Atomic mass- defined to be 1/12 the mass of the carbon- 12 nuclide
Number of neutrons= A- Z
All atoms of an element contain the same number of protons in their nuclei. Atoms containing the same
number of protons but different numbers of neutrons are isotopes of the same element.
Isotopes of Hydrogen
Name
Protium
Deuterium
Tritium
Protons
1
1
1
Neutrons
0
1
2
Mass Number
1
2
3
Average atomic mass can be determined from relative amounts of each isotope.
The Rutherford- Bohr Atom
Rutherford and Bohr pictured the atom as consisting of a central nucleus surrounded by electrons
in orbits.
Substances excited by an energy source emit light in definite wavelengths called a spectrum.
Planck’s Hypothesis
Planck stated that energy is radiated in discrete units called quanta. A photon is a quantum of
light energy.
Electrons and Clouds
Waves and Particles
De Broglie’s Hypothesis
o
He suggested that particles have characteristics of waves.
Heisenberg’s Uncertainty Principle
o
The principle states that the exact location and momentum of an object cannot be determined at the same
time.
Four Wave Characteristics:
1. Wavelength
2. Frequency
3. Velocity
4. Amplitude
Amplitude of a wave is its maximum displacement from a base line.
Schrodinger’s Work
o
Schrodinger considered the electron as a wave.
o
The four quantum numbers in Schrodinger’s equation are used in describing electron
o
behavior
Quantum numbers represent the different electron energy.
Four Quantum Numbers
1. n- number of the energy level of an atom (1, 2, 3, 4)
2. l- the number of sublevel of n (s, p, d, f)
3. ml- the number of an orbital in a sublevel
Each orbital may contain a pair of electrons.
4 ms- describes the spin of an electron
+1/2 for clockwise spin; -1/2 for counter clockwise spin
Summary of Allowed Combinations of Quantum Numbers
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UPCAT REVIEWER 2017: Chemistry
n
l
M
Subshell
Notation
Number Number of
of
Electrons
Total
Orbitals Needed to Number of
in the
Fill
Electrons in
Subshell Subshell
Subshell
1
0
0
1s
1
2
2
0
0
2s
1
2
2
1
1,0,-1
2p
3
6
3
0
0
3s
1
2
3
1
1,0,-1
3p
3
6
3
2
2,1,0,1,-2
3d
5
10
4
0
0
4s
1
2
4
1
1,0,-1
4p
3
6
4
2
2,1,0,1,-2
4d
5
10
4
3
3,2,1,0,1,-2,-3
4f
7
14
2
8
18
32
In an atom, the number of protons in the nucleus equals the number of electrons in the charge cloud. The
energy level closest to the nucleus will fill first.
Pauli’s Exclusion Principle
o
The principle states that no two electrons in an atom can have the same set of four quantum numbers.
The Diagonal Rule
This rule is used to approximate electron configuration. The atomic number (Z) is equal to the sum of the
superscripts in the electron configuration.
 The Periodic Table
Surveying the Table
The periodic can be used to determine electron configuration of an element.
Period- horizontal rows of elements
Group- vertical columns of elements
o
Atoms in the same period have the same principal quantum number (n).
o
For groups IIIA- VIIIA the endings are p1- p6 and the coefficient equals the period number.
o
For groups IIIB- IIB the endings are d1- d10 and the coefficient is one less than the period number.
o
When endings are f1-f14, the coefficient is two less than the period number.
Octet Rule-Eight electrons in the outer level of an atom represent a stable arrangement.
Metals and Non-metals
Metals- have fewer electrons in the outer level
Metalloids- elements which have both metallic and non-metallic properties
Nonmetals- have more electrons in the outer level
Periodic Properties
Many properties of atoms are due to the average distance of the outer electrons from the nucleus
and to the effective nuclear charge experienced by these electrons.
The effective nuclear charge experienced by the outer electrons is determined primarily by the difference
between the charge on the nucleus and the charge of the core electrons.
Zeff = Z – S
Z – protons in the nucleus
S – average number of electrons

The Zeff increases as we move left to right across a period.
Atomic Size

Atomic size increases as we go down a column in the periodic table and decrease as we proceed left to
right across a row.
Ionization Energy- minimum energy needed to remove an electron from the atom in the gas phase,
forming a cation

Ionization energy decreases as we go down a column and increase as we proceed left to right across a
row.
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UPCAT REVIEWER 2017: Chemistry
Electron Affinity- energy change upon adding an electron to an atom in the gas phase, forming an anion

Electron affinities become more negative as we proceed from left to right across the periodic table, and
more negative as we go up a column.
Electronegativity- tendency of an atom to attract shared electrons

Electronegativity decreases as we go down a column and increase as we proceed left to right across a row.
Remember: The greater the difference in electronegativity between two atoms, the more polar their bond.
Metallic Character

Metallic character increases as we proceed down a column and decreases as we proceed from left to right.
The Process of Bonding
Bond character between atoms depends upon the electronegativity differences.
Chemical Bond Summary
Bond Type
Generally Formed
Between
Properties Associated
with Bond Type
Bond Formed By
Covalent
Atoms of non-metallic
elements
of
similar
electronegativity
Sharing
pairs
Ionic
Atoms of metallic and
non-metallic elements of
widely
different
electronegativities
Electrostatic attraction
between ions resulting
from
transfer
of
electrons
Metallic
Atoms
of
elements
Common exchange of
outer
electrons
between
atoms
of
lower electronegativity

metallic
of
electron
Stable
nonionizing
moleculesnot
conductors of electricity
in any phase
Charged ions in gas,
liquid, and solid. Solid
electrically nonconducting.
Gas
and
liquid
are
conductors. High melting
points.
Electrical conductors in all
phases- lustrous- very
high melting points
Examples of
Substances
Utilizing Bond
Type
OF2, C2H6, AsCl3,
SiC, GeCl4
BaS, NaCl, CdF2,
Ca3N2, BaBr2
Cu, Zn, Au, Na,
Fe, Gd, Dy, Be
The van der Waals forces are attractions between the molecules of nonpolar covalent substances. These
forces are generally believed to be caused by a temporary dipole, or unequal charge distribution, as
electrons constantly move about in an atom, ion, or molecule.
Results of Bonding
Shared Pairs- outer electron pairs attracted by two nuclei
Unshared Pairs- outer electron pairs attracted to one nucleus
VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY
Electron pairs spread as far apart as possible to minimize repulsive forces.
The shape of a molecule containing three or more atoms is determined by the number of shared and
unshared pairs.
 If the central atom has two shared pairs and no unshared pairs, the molecule is linear.
 If there are three shared pairs and no unshared pairs, the molecule is trigonal planar.
 If there are four shared pairs, the molecule is tetrahedral. If there are three shared pairs and one
unshared pair, the molecule is pyramidal.

unshared- unshared
repulsion
>
unshared- shared
repulsion
>
shared- shared
repulsion
Outer
Atoms
2
3
Lone
Pairs
0
0
Charge
Clouds
2
3
Shape
Linear
Trigonal Planar
2
4
3
1
0
1
3
4
4
Bent
Tetrahedral
Trigonal Pyramidal
2
5
2
0
4
5
Bent
TrigonalBipyramidal
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4
3
2
2
5
5
Seesaw
T-Shaped
2
6
3
0
5
6
Linear
Octahedral
5
4
1
2
6
6
Square Pyramidal
Square Planar
Combining Atomic Orbitals (Hybrid Orbitals)
Sigma and Pi Bonds
A covalent band is formed when an orbital of one atom overlaps an orbital of another atom.
Orbitals may over lap when electrons are shared.
sigma bond (σ)- overlap of two s orbitals
overlap of an sand a p orbital
end to end overlap of two p orbitals
pi bond (π)- sideways overlap with parallel axes of two p orbitals
Resonance- equivalent alternative structures for a molecule or polyatomic ion which lead to “average” bond
lengths
STRUCTURE AND PROPERTIES OF MOLECULES
Isomers- have the same molecular formula but different structures. Isomers are not resonance forms.
Isomers may be: geometric, structural, positional, functional, or optical.
Polarity- A polar covalent bond occurs when a shared pair of electrons is attracted more strongly to one of
the atoms. Polar bonds, unless symmetrically arranged, produce polar molecules. (e.g. CCl4 - four C- Cl
bonds are polar, but their symmetrical arrangement, tetrahedral, produces a nonpolar molecule)
Dipole- a polar molecule

States of Matter
SOLIDS
Types of Crystalline Solids
Type of Solid
Form
of
Particle
Covalent
Atoms
or
molecular
molecules
Covalent
network
Ionic
Metallic
Atoms
connected
in a network
of covalent
bonds
Positive and
negative
ions
atom
s
Forces between
Particles
London
dispersion,
dipoledipole
forces,
hydrogen
bonds
Covalent bonds
Properties
Examples
Fairly soft, low to
moderately
high
melting point, poor
thermal
and
electrical conduction
Ar, CH4, sucrose
(C12H22O11), CO2
Very hard, very high
melting point, often
poor thermal and
electrical conduction
Diamond,
quartz, SiO2
Electrostatic
attractions
Hard and brittle,
high melting point,
poor thermal and
electrical conduction
Soft to very hard,
low to very high
melting
point,
excellent
thermal
and
electrical
conductivity,
malleable
and
ductile
Typical
saltsNaCl, Ca(NO3)2
Metallic
bonds
C;
All
metallic
elements- Cu, Fe,
Al., Pt
Amorphous Materials
Amorphous or are substances which appear to be solids but are not crystalline. Their structure reveals a
disordered arrangement of particles. Amorphous substances are called supercooled liquids. Butter is a
good example.
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Viscosity- resistance of a liquid to flow
The viscosity of a substance usually decreases as the temperature increases.
LIQUIDS
For a substance, the freezing point of the liquid state is equal to the melting point of the solid state.
Vapor Equilibrium
Vapor refers to the gaseous state of substances which are liquids or solids at room temperature.
In a closed container, the vapor phase is in equilibrium with its solid or liquid phase. In dynamic
equilibrium, forward and reverse reactions occur at the same rate. When a substance is in equilibrium with
its vapor, the gaseous phase of the system is said to be saturated with the vapor of that substance.
Vapor Pressure
Vapor pressure is directly dependent on temperature.

Le Chatelier’s Principle states that, if stress is applied to a system at equilibrium, the system readjusts so
that the stress is reduced. The stress may be a change in temperature, pressure, concentration, or other
external forces.
Melting point- the temperature at which the vapor pressure of the solid and the vapor pressure of the
liquid are equal
Sublimation- phase change directly from solid to gas
Evaporation- the process of a molecule leaving the surface of the liquid or solid and entering the gaseous
state
Liquefaction- changing a gas to a liquid

Condensation Point- also called the boiling point of a liquid
Boiling Point
Boiling occurs when the vapor pressure of a liquid is equal to the atmospheric pressure over the liquid. The
boiling point of a substance decreases as the atmospheric pressure decreases.
Volatile- liquid which boils at a low temperature and evaporates at room rapidly temperature
Nonvolatile- liquids which boil at high temperature and evaporate slowly at room temperature
Heat of fusion- the heat required to melt one gram of a substance at its melting point
Heat vaporization- the heat required to vaporize one gram of a substance at its boiling point
Properties of Liquids
Hydrogen Bonding
Hydrogen bonded to a strongly electronegative element causes some substances to differ from predicted
behavior. A hydrogen atom sandwiched between two electronegative atoms makes up a hydrogen bond.
Nitrogen, oxygen, and fluorine have sufficient electronegativity to cause hydrogen bonding. A hydrogen
bond is a dipole attraction.
Surface Tension
Surface tension of liquids is due to the unbalanced forces on surface particles. Surface tension causes drops
of liquid to be spherical. Capillary rise of liquids in small tubes is due to surface tension.
Other properties:
 phase change
 specific heat capacity
 viscosity
 density
 heat of vaporization
 conductivity
 capillary action
 cohesion
GASES
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Under the same temperature change, gases change volume more than solids or liquids. Gas particles move
about at random. An ideal gas is composed of point masses with no mutaual attraction.
STP= 1.01325 bar and 0°C
The Gas Laws
Boyle’s Law states:
If the temperature of a gas remains constant, the pressure exerted by the gas varies inversely as the
volume.
Pressure exerted by a gas depends on:
1. number of particles/ unit volume
2. average kinetic energy of particles
Dalton’s Law of Partial Pressure states: The total pressure in a container is the sum of the partial
pressures of the gases in the container. When a gas is one of a mixture, the pressure it exerts is called its
partial pressure.
Charles’ Law states: The volume of a quantity of gas, held at a fixed pressure, varies directly with the
Kelvin temperature.
Combined Gas Law states: A change in volume resulting from a change in both temperature and pressure
can be found by combining the ratios.
Properties of Gases
Diffusion- the random scattering of gas particles
Graham’s Law states: The relative rates at which two gases under identical conditions of
temperature and pressure will pass through a small hole varies inversely as the square roots of the
molecular masses of the gases.
Gas Density
If the number of particles remains the same:
1. density increases as pressure increases
2. density decreases as temperature increases
Energy and Disorder
Natural processes:
High energy  low energy
Order  disorder
Isothermal: constant temperature
Isobaric: constant pressure

Energy Changes
Heat Measurement
Calorimeter- it is used to measure energy changes in chemical reactions
Specific heat capacity (J/ g.C)-the heat needed to raise the temperature of one gram of the substance
1oC
Molar (or atomic) heat capacity- the quantity of heat needed to raise the temperature of one mole of a
substance 1oC
- it is found by multiplying the specific heat capacity of the substance by its formula (or atomic) mass
Heat of Chemical Reaction
C (c) + O2 (g)  CO2 (g) + heat (393.5 kJ)
One mole carbon reacts with one mole of oxygen to produce on a mole of carbon dioxide and 393.5 kJ of
heat. The heat released (393.5 kJ) is called heat of reaction.
Heat of reaction- energy absorbed or released during a chemical reaction
Heat content or enthalpy (H) – that part of the energy of a substance which is due to the motion of its
particles
Heat content of a free element is defined as zero, at standard atmospheric pressure and 25oC.
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Heat of Formation- change in enthalpy (∆H) when one mole of a compound is produced from the free
elements, ∆Hof. (It is expressed usually in terms of kJ/ mole.)
Enthalpy change, ∆Ho – the difference between the enthalpy of the products and the enthalpy of the
reactants
Knowledge of enthalpy change can help determine whether a reaction will be exothermic or endothermic.
Exothermic reaction- negative ∆Ho
Endothermic reaction- positive ∆Ho
Enthalpy (H)-measure of the change in heat content of the reactants and products
Exothermic: H<0
Endothermic: H>0
Entropy (S)- measure of the change in the randomness or disorder of the reactants and products
Increase in disorder: S>0
Decrease in disorder: S<0
Free Energy (G)- the energy available to do useful work
Spontaneous reaction: G<0
Non- spontaneous reaction: G>0

Free Energy Change (G)- the maximum amount of energy that a reaction can theoretically deliver
Hess’s Law
Hess's Law says that the net change in enthalpy of a reaction will be the sum of the enthalpy changes that
occur along the way.
 Solutions
A solution is a homogenous mixture of two or more substances. The substance present in greater quantity
is usually called the solvent. The other substance in the solution is known as the solute; they are to be
dissolved in the solvent.
For example, when a small amount of sodium chloride (NaCl) is dissolved in a large quantity of water, the
water is the solvent and the sodium chloride is the solute.
Solubility of a solute depends on the temperature.
Saturated solution- a solution in which undissolved substance is in equilibrium with the dissolved
substance
Unsaturated solution- a solution containing less than the saturated amount of solute for that
temperature
Supersaturated- a solution containing more solute than a saturated solution
Solution Proportions
Concentrated solution- a relatively large amount of solute is present per unit volume
Dilute solution- a relatively small amount of the solute is present per unit volume
Solution rate is affected by
1. surface area exposed to solvent
2. kinetic energy of the particles
Colloids
Colloids are mixtures of two phases of matter.
1. dispersed phase
2. continuous phase
Colloidal particles are intermediate in size between solutions and suspensions.
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UPCAT REVIEWER 2017: Chemistry
Properties of solutions, colloids, and suspensions
Solutions
Colloids
Suspensions

Do not settle

Do not settle

Settle out on
out
out
standing

Pass

Pass

Separated
unchanged
unchanged
by ordinary
through
through
filter paper
ordinary
ordinary

Separated
filter paper
filter paper
by a

Pass

Separated
membrane
unchanged
by a
through
membrane

Scatter light
membrane

Do not affect

Do not

Scatter light
colligative
scatter light

Do not affect
properties

Affect
colligative
colligative
properties
properties
Properties of Colloids
Tyndall Effect- the scattering of light by colloid particles
Brownian movement- the constant random motion of colloidal particles
Adsorption- occurs when solid or liquid surfaces attract and hold substances
Electrophoresis- migration of charged colloidal particles within an electrical field
Le Chatelier’s Principle
If a system is in equilibrium and a condition is changed, then the equilibrium will shift toward
restoring the original conditions.
 Increasing the concentration of reactant will produce a greater concentration of product.
 Increased pressure on a reaction system with a gas phase has the same effect as increased concentration.
 Optimum conditions are those which produce the highest yield of product.
Concentration of Solutions
Molarity
Molarity (symbol M) expresses the concentration of a solution as the number of moles of solute in a liter of
solution (soln)
Molarity =
moles of solute ___
volume of solution in liters
Molality
Molality (symbol m) is expressed as
Molality =
moles of solute____
kilograms of solvent
Mole Fraction of component
Mole Fraction =
moles of component______
total moles of all components
Mass Percentage
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Mass percent of component = mass of component in solution
total mass of solution
x 100
Solutions of known molarity can be formed either by weighing out the solute and diluting it to a known
volume or by the dilution of a more concentrated solution of known concentration (a stock solution).
Outline to the procedure used to solve stoichiometry problems that involve measured (laboratory) units of
mass, solution concentration (molarity), or volume.
Factors Affecting Solubility
Solute- Solvent Interaction
One factor determining solubility is the natural tendency of substances to mix (the tendency of
systems to move towards disorder). In general, when other factors are comparable, the stronger the
attractions between solute and solvent molecules, the greater the solubility.
Like dissolves like.
Substances with similar intermolecular attractive forces tend to be soluble in one another. Polar
solutes dissolve in polar solvents. Nonpolar solutes dissolve in nonpolar solvents.
Pressure Effects
When pressure is increased, the rate at which gas molecules enter the solution increases. The
concentration of solute molecules at equilibrium increases in proportion to the pressure.
Temperature Effects
The solubility of most solid solutes in water increases as the temperature of the solution increases In
contrast to solid solutes, the solubility of gases in water decreases with increasing temperature.
 Acids, Bases, and Salts
Acids, bases, and salts are electrolytes. Electrolytes conduct a current.
SUMMARY OF ACID- BASE THEORIES
Theory
Arrhenius
Theory
Acid
Definition
Any
substance
which
releases
H+
ion in water
Base
Definition
Any
substance
which
releases OHions in water
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UPCAT REVIEWER 2017: Chemistry
solution
solution.
BronstedLowry Theory
Any
Any
substance
substance
which
which
donates
a
accepts
a
proton.
proton.
Lewis Theory
Any
Any
substance
substance
which
can
which
can
accept
an
donate
an
electron pair.
electron pair.
A substance that is an acid or base under the Arrhenius theory is also an acid or base under the Lewis and
Bronsted- Lowry theories.
Binary Acids- acids containing only two elements
The names of binary acids begin with hydro- and ends with –ic.
Naming:
HCl - hydrochloric acid
HI – hydroiodic acid
HS – hydrosulfuric acid
Ternary Acids- acids which contain three elements
Acid Anhydrides and Basic Anhydrides
Anhydride means without water.
Acid anhydride- any oxygen- containing substance which will produce an acid when dissolved in water
Basic anhydride- any oxygen- containing substance which will produce a base when dissolve in water
Na2O (s)
+ H2O (l)  2NaOH (aq)
Basic Anhydride + water 
base
Nonmetals tend to form acids, and metals will tend to form bases.
An acid- base neutralization reaction produces salt. Water solutions of some salts are not neutral.
Salt- a crystalline substance formed from the combination of the negative ions of the acid and the positive
ions of the base when water is evaporated
Acidic salts and basic salts do not produce neutral solutions.
Naming:
Ternary acids
HMnO4

Permanganic acid
H2SO4

Sulfuric acid
HNO2

Nitrous acid
Salt
MnO4permanganate
SO42sulfate
NO2nitrite
 Electrolytic Properties
A substance (such as NaCl) whose aqueous solutions contain ions is called an electrolyte. A substance
(such as C12H22O11) that does not form ions in solution is called a nonelectrolyte.
The presence of ions causes aqueous solutions to become good conductors. Ions carry electrical charge
from one electrode to another, completing the electrical circuit.
Molecular and Ionic Compounds in Water
When an ionic solid dissolves in water, it dissociates into its component ions. Positive ions
(cations) are attracted by the negative end of H2O (the O atom which is rich in electrons), and the
negative ions (anions) are attracted by the positive end (the H atoms). This prevents anions and cations
from recombining.
When a molecular compound dissolves in water, the solution usually consists of intact molecules
dispersed throughout the solution. Most molecular compounds are nonelectrolytes.
EXAMPLES:
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
Sodium sulfate (Na2SO4) dissociates into sodium ions (NA+) and sulfate ions (SO42-) [ionic]
<you must remember the formulas and charges of common ions>

Methanol, CH3OH,in water dissociates into CH3OH molecules [molecular]
Acid Base Reactions
Those acids and bases that are strong electrolytes are called strong acids and strong bases,
respectively. Those that are weak electrolytes are weak acids and weak bases. When solutions of acids
and bases are mixed, a neutralization reaction occurs.
HCl (aq)
(acid)
+
NaOH (aq)
(base)

H2O (l)
(water)
+
NaCl (aq)
(salt)
In general, a neutralization reaction between an acid and a metal hydroxide produces water and a salt.
Strong electrolytes- solutes that exist in solution completely or nearly completely as ions
Weak electrolytes- solutes that exist mostly in the form of molecules with only a small fraction in the
form of ions
Common Strong Acids and Bases
Strong acids
Hydrochloric, HCl;
Hydrobromic, HBr;
Hydroiodic, HI;
Chloric, HClO3;
Perchloric, HClO4;
Nitric, HNO3; Sulfuric,
H2SO4


Strong Bases
Group Ia metal
hydroxides (LiOH,
NaOH, KOH, RbOH,
CsOH)
Heavy group 2A metal
hydroxides (Ca(OH)2,
Sr(OH)2, Ba(OH)2)
If an acid is added to water, the [H3O+] increases and [OH-] decreases.
If a base is added to water, the [H3O+] decreases and [OH-] increases
The pH Scale
The pH scale is a simplified way of stating the concentration of H3O+ ions in solution. Hydrogen ion
concentration is measured by pH.
Acidic solution: pH < 7
Neutral solution: pH = 7
Basic solution:
pH > 7
The pOH scale
pOH scale is a way of stating the concentration of OH- ions in solution.
The sum of pH and pOH is 14.
 Buffers
A buffer system can absorb acids or bases without significant change in pH. Buffer solutions are prepared
by using a weak acid or a weak base and one of its salts.
FOR A WEAK ACID:
HA +OH- H2O + AA- + H3O  HA + H2O
where: HA – weak acid
OH- – added base
A- - negative ion from the salt
FOR A WEAK BASE:
MOH + H3O+ M+ + 2H2O
M+
+ OH- MOH
where: MOH- weak base
H3O+- added acid
M+- positive ion from salt
 Oxidation- Reduction Reactions (Redox)
Oxidation is the loss of electrons by a substance.
LEORA- Loss Electrons Oxidation Reducing Agent
Reduction is the gain of electrons by a substance.
GEROA- Gain Electrons Reduction Oxidizing Agent
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Reducing agent is the substance which gives up electrons. Oxidizing agent is the substance which gains
electrons.
Oxidation number of an atom in a substance is the actual charge of the atom if it is a monoatomic ion;
otherwise, it is the hypothetical charge assigned to the atom using a set of rules.
We use the following rules for assigning the oxidation numbers:
1. For an atom in its elemental form the oxidation number is always zero.
Each H atom in H2 has an oxidation number of 0.
Each P in the P4 molecule has an oxidation number of 0.
2. For any monoatomicion, the oxidation number equals the charge on the ion.
K+ as an oxidation number of +1
S2- has an oxidation number of -2
3.
Nonmetals usually have negative oxidation numbers, although they can sometimes be positive.
4.
The sum of the oxidation numbers of all atoms in a neutral compound is zero. The sum of the oxidation
numbers in a polyatomic ion equals the charge of the ion.
H3O+
H= +1
O= -2
Thus, 3(+1) + (-2) = +1, which equals the net charge of the ion.

Alkali metal ions (group 1A) always have a 1+ charge. Alkaline earth metals (group 2A) are always 2+,
and aluminum (group 3A) is always 3+.
Remember: Whenever one substance is oxidized, some other substance must be reduced.
The reaction of a metal with either an acid or a metal salt conforms to the following general pattern:
A + BX
AX + B
Examples: Zn (s)+ 2HBr (aq)
ZnBr2(aq) + H2(g)
Mn(s) + Pb(NO3)2(aq)Mn(NO3)2(aq) + PbS
These reactions are called displacement reactions because the ion in solution is displaced or replaced
through oxidation of an element. Many metals undergo displacement reactions with acids, producing salts
and hydrogen gas.
Mg(s) + 2HCl (aq)
MgCl2 (aq) + H2 (g),
0
+1 -1
+2 -1
0
The oxidation number of Mg changes from 0 to +2 meaning, the atom lost electrons due to increase in
oxidation number (oxidized). The H+ ion of the acid decreases in oxidation number meaning it gained
electrons (reduced). Oxidation number of Cl- remains the same, -1 (spectator ion). The net ionic equation
is:
Mg (s) + 2H+ (aq)
Mg2+ (aq) + H2 (g)
Remember: Whenever one substance is oxidized, some other substance must be replaced.
 Electrochemistry
Salt bridge- solution containing ions in a U- tube
Electrolytes- substances whose solutions conduct electricity

Anode- positive electrode

Cathode- negative electrode
Negative ions (anions) are attracted to the anode. Positive ions (cations) are attracted to the cathode.
Chemical change at the cathode is reduction. At the anode, negative ions are oxidized by losing electrons.
Chemical change at the anode is oxidized.
Voltaic cell- is a cell which produces an electric current and is composed of two dissimilar metals and an
electrolyte
Reduction potential- measures the relative strength of oxidizing and reducing agents
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In an electrochemical reaction, there is always an oxidation half- reaction and a reduction half- reaction.
Any oxidation- reduction reaction is an equilibrium reaction. Therefore, any change in temperature,
pressure, or concentration will affect the flow of electric current.
 Nuclear Chemistry
Nuclear decay is spontaneous.
Three types of natural radiation:
1. alpha particles (He nuclei)
2. beta particles (electrons)
3. gamma rays (high energy x rays)
Positron- positive electrons (β+ or
0
+1e)
Transmutation- it occurs when an atom with a different atomic number is produced
Binding energy- energy needed to separate the nucleus into individual particles
Binding energy maybe increased by several kinds of nuclear reactions:
1. α- particle emission
2. κ- electron capture
3. β+ emission
4. β- emission
5. neutron emission
Half-life- it is the length of time required for one- half of the atoms of a radioactive sample to decay
Alpha decay
In alpha decay, the nucleus emits an alpha particle; an alpha particle is essentially a helium nucleus, so it's
a group of two protons and two neutrons.
An example of an alpha decay involves uranium-238:
The process of transforming one element to another is known as transmutation.
Beta decay
A beta particle is often an electron, but can also be a positron, a positively-charged particle. If an electron
is involved, the number of neutrons in the nucleus decreases by one and the number of protons increases
by one. An example of such a process is:
Gamma decay
The third class of radioactive decay is gamma decay, in which the nucleus changes from a higher-level
energy state to a lower level. Similar to the energy levels for electrons in the atom, the nucleus has energy
levels.
Fission- the break- up of a heavy nucleus into two approximately equal parts
Fusion- two or more smaller nuclei combine to form a larger nucleus

Intramolecular and Intermolecular Forces
Intramolecular Forces
 Forces of electrostatic attraction within a molecule.
 Occurs between the nuclei of atoms and their electrons making up the molecule
i.e Covalent bond, Ionic Bond, Metallic Bond
Intermolecular Forces
 Forces of attraction and repulsion between molecules that hold molecules, ions, and atoms
together.
i.e Dipole-dipole interactions, Hydrogen Bond, London Dispersion Forces
 Much weaker than the intramolecular forces of attraction
Determine the physical properties of molecules like boiling point, melting point, density, and enthalpies of
fusion and vaporization
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Types of Intermolecular forces that exist between molecules
1. Dipole-Dipole Attractions
For polar molecules, attractive forces called dipole-dipole attractions occur between
the positive end of one molecule and the negative end of another.
Example: For a polar molecule with a dipole such as HCL, the partially positive H atom of one HCl molecule
attracts the partially negative Cl atom in another molecule.

2.
Hydrogen Bond


3.
Occurs between the partially positive hydrogen atom of one molecule and a lone pair of
electrons on a nitrogen, oxygen, or fluorine atom in another molecule
Strongest type of forces between polar molecules
Dispersion Forces
 Very weak attractions that occur between nonpolar molecules
 Physical properties such as melting point are low for compounds with weak attractive
forces (i.e dispersion forces).
Chemical Equilibrium
 Achieved when the rates of the forward and reverse reactions are equal and the concentrations of
the reactants and products remain constant.
Law of Mass action
 Formulated by two Norwegian chemists, Cato Guldberg and Peter Waage, in 1864.
 for a reversible reaction at equilibrium and a constant temperature, a certain ration of reactant
and product concentrations has a constant value, K (the equilibrium constant).
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UPCAT REVIEWER 2017: Chemistry
The Equilibrium Constant
Consider the following reaction at equilibrium:
Where a, b, c and d are the stoichiometric coefficients for the reacting species A, B, C, and D. For the
reaction at a particular temperature
Where K is the equilibrium constant.
Note that although the concentrations change, according to the law of mass action, the value of K remains
constant.
 The magnitude of the equilibrium constant tells us whether an equilibrium reaction favors the
products or reactants
 If K is much greater than 1 (that is K>>>1), the equilibrium will lie to the right and favors the
products.
 If K is much less than 1 (that is K<<<1), the equilibrium will lie to the left and favor the
reactants.

Organic Chemistry
Alkanes
 often described as saturated hydrocarbons
Hydrocarbons because they contain only carbon and hydrogen
Saturated because they have C-C and C-H single bonds and thus contain the maximum possible number
of hydrogen per carbon
 Occasionally called aliphatic compounds, a name derived from the Greek aleiphas, meaning “fat”
 General Formula: CnH2n+2
where n is an integer
Naming of Alkanes
I.
Straight-chain alkanes are named according to the number of carbon atoms they contain, as
show in the table.
II.
With the exception of the first four compounds- methane, ethane, propane, and butane –whose
names have historical roots, the alkanes are based on Greek numbers.
III.
The suffix –ane is added to the end of each name to indicate that the molecule identified is an
alkane.
Alkenes
 sometimes called olefin
 a hydrocarbon that contains a carbon-carbon double bond
 occur abundantly in nature
 ethylene and propylene, the simplest alkenes, are the two most important organic chemicals
produced industrially
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UPCAT REVIEWER 2017: Chemistry
Naming Alkenes
Step 1. Name the parent hydrocarbon
Find the longest carbon chain conatining the double bond, and name the compound accordingly, using the
suffix –ene:
Step 2. Number the carbon atoms in the chain.
Begin at the end nearer the double bond or, if the double bond is equidistant from the two ends, begin at
the end nearer the first branch point. This rule ensures that the double- bond carbons receive the lowest
possible numbers.
Step 3. Write the full name.
Number the substituents according to their positions in the chain, and list them alphabetically. Indicate the
position of the double- bond by giving the number of the first alkene carbon and placing that number
directly before the parent name. If more than one double bond is present, indicate the position of each and
us one of the suffixes –diene, -triene, and so on.
Cycloalkenes are named similarly, but because there is no chain end to begin from, we number the
cycloalkane so that the double bond is between C1 and C2 and the first substituent has as low number as
possible. It’s not necessary to indicate the position of the double bond in the name because it’s always
between C1 and C2.
As with open-chain alkenes, newer but not yet widely accepted naming rules place the locant immediately
before the suffix in a diene.
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