SPSTEM05 REVIEWER - STEM11E
Colegio de San Juan de Letrán
151 Muralla St. Intramuros, Manila
Senior High School Department
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SPSTEM05 - GENERAL CHEMISTRY
REVIEWER FOR SUMMATIVE AND
PERIODICAL EXAM
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1ST SEMESTER
Published by: Meagan Elise T. Perdito
STEM12E
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I.
LESSON 1 : Laboratory
Apparatus
1. Beaker
- used for holding
various chemicals.
- Not for measuring
precisely.
2. Graduated Cylinder
- Used to precisely
measure the volume
of liquids or run
experiments.
- Read from the
meniscus at eye level.
(lower)
3. Erlenmeyer Flask
- Used to
approximately
measure the volume
of various liquids.
- Useful for mixing by
swirling
4. Florence Flask
- Used to boil liquids.
- Also used to collect
gases, if applicable.
5. Volumetric Flask
- Used to prepare
standard solutions.
- Used for dilution
- They are only good
for 1 specific volume.
6. Reagent Bottle
- Used to store,
transport, or view
reagents such as acids
or bases.
7. Rubber Stoppers
-
Used to close flasks
and test tubes.
- The holes allow the
insertion of glass
tubing, probes, or
thermometers as
needed by the
experiment.
8. Test Tubes and Rack
- Used to hold
chemicals/tubes while
experimenting.
- Not for measuring
precisely.
- Aim away from faces.
9. Buret and Buret Clamp
- Used for precisely
measuring dispensed
liquids
- Holds buret to ring
stand.
10. Ring Stand and Ring
Clamps
- Base/Pole of set-up
for experimentation.
- Holds glassware in
place for heating or
evaporating.
11. Test Tube Brushes
- Cleaning.
- You must clean the
tubes before and after
you use them.
12. Test Tube Holder
- Used for carrying or
holding hot test tubes.
13. Thermometer
- Measuring
temperature.
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14. Hot Plate
- Used to heat
substances.
15. Bunsen Burner
- Used to heat
substances quickly or
if > 400oC is needed.
16. Rubber Tubing
- Used for a variety of
things, such as
connecting a Bunsen
burner to a gas valve
stem or Connecting
glass tubing together.
17. Wire Mesh or Gauze
- Used to absorb and
spread the heat of
flame.
- Keeps glassware from
cracking and
breaking.
- Part of ring stand
set-up.
18. Clay Triangle
- Used to hold a
crucible in place on a
ring stand.
- Also helps absorb and
spread heat of flame.
- Part of ring stand
set-up.
19. Crucible and Cover
- Used for heating
substances.
- Can withstand high
direct heat.
20. Crucible Tongs
- Used to carry the
crucible.
21. Beaker Tongs
- Used to carry beakers.
22. Mortar and Pestle
- Used to grind
substances into
powder or slurry.
23. Scoopula/spatula
- Used to scoop
chemical powders.
- Not a measuring
instrument.
24. Stirring Rods
- Used to stir
substances.
25. Watch Glass
- Used to evaporate
liquids and cover
beakers during sample
preparation.
- It is also used to hold
solids during
weighing.
26. Evaporating Dish
- Used to evaporate
excess liquids.
27. Centrifuge
- Used to separate
suspensions (solids
from liquids).
28. Funnel
- Used to safely
transfer substances
from one container to
another.
29. Dropper and Bottle
- Used to measure out
small amounts of
liquids for
experiments.
30. Wash Bottle
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-
Usually contains
deionized water.
- Handy for rinsing
glassware and for
dispensing small
amounts of dH2O for
chemical reactions.
31. Digital Balance
- Used to accurately
measure mass.
- Only up to 200g in
our labs.
32. Triple Beam Balance
- Measures the mass of
an object.
- Make certain the
balance is calibrated
correctly before use.
33. Double Pan Balance
- Used to compare the
masses of two
substances.
34. Centigram Balance
- Used to precisely
measure the mass of a
substance.
- More precise than the
triple beam balance.
35. Pipet, Pump, and Bulb
- Used to precisely
measure the volume
of liquids in small
amounts.
II.
LESSON 2 - MATTER AND ITS
PROPERTIES
★ Matter
- defined as anything
that occupies space
and has mass.
A. States of Matter
1. Solid
- Definite shape and
volume
- Highest density
- Cannot flow and
cannot be compressed
- Particles are tightly
packed with
maximum force of
attraction between
them
- Particles cannot
move, but vibrate
only at their fixed
position
- Minimum kinetic
energy
2. Liquid
- Indefinite shape;
adjusts to the shape of
its container
- Definite volume
- Lower density than
solid
- Can flow but cannot
be compressed
- Less force of
attraction than solids
- Particles are loosely
packed and can slide
over each other
- More kinetic energy
than solids
3. Gas
-
Indefinite shape and
volume
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-
-
Lowest density
Can flow and can be
compressed
Negligible force of
attraction
Particles are loosely
packed and can move
freely
Maximum kinetic
energy
4. Plasma
- Formed by heating
and ionizing a gas
- A superheated matter
comprising 99% of
the visible universe.
5. Bose-Einstein Condensate
- Forms only when
materials are cooled
to temperatures very
close to absolute zero
- Particles moving
together in the same
direction and at the
same speed
B. Classification of Matter
Matter may be classified as a pure
substance or a mixture.
1. Pure substances
- Have definite
composition and
distinct properties.
- Elements and
Compounds fall under
pure substances
Elements - considered to be the simplest
form of matter. There are 118 known
elements and 92 of them occur naturally on
earth.
Compounds - consist of 2 or more elements
chemically combined. Household cleaning
materials (bleaches, soaps, and detergents)
and Personal care products (talc, glycerin,
and artificial dyes and lead) fall under
compounds.
a. Glycerin - Provides hydration and
gives the smooth feel of the skin.
b. Artificial Dyes and Lead Dangerous and hazardous substances
c. Talc - Generally recognized as safe.
It prevents caking of the product and
ensures smooth distribution of
makeup.
2. Mixtures
- Combinations of two or more
substances in which the
substances retain their
distinct identities.
- Components of a mixture can
be separated from one
another by physical means as
they are not chemically
combined.
- Mixtures can be classified as:
- Homogenous - uniform
composition of the mixture
throughout. It is difficult to
distinguish the components
of the mixture due to the even
distribution.
- Heterogenous - lacks uniform
composition. Components are
easily identifiable and it has
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at least 2 phases remaining
separate from each other.
C. Physical and Chemical Properties
1. Physical Properties
- Can be observed or
measured without
changing the
composition of the
substance.
- These properties
include: density,
hardness, malleability,
viscosity, boiling
point.
2. Chemical Properties
- Can only be observed
when the substance
undergoes a chemical
change.
- These properties
include: chemical
reactivity,
flammability, ability
to oxidize.
D.
Intensive and Extensive Properties
1. Intensive Properties
- Do not depend on the
amount of matter in a
sample.
- Temperature, boiling
point, concentration,
luster
2. Extensive Properties
- Depends on how
much matter a sample
contains
- Weight, length,
volume, entropy
E. Physical and Chemical Changes
1. Physical Change
- change that does not
lead to the formation
of new substances.
- involves only a
change in the physical
properties and not the
composition.
- Melting ice, boiling
water, breaking glass
2. Chemical Change
- occur when a
substance reacts and
produces one or more
new substances. (Iron
rusting, fireworks,
burning, cooking an
egg)
III.
LESSON 3 - Measurements
A. System of Measurement
- A measured quantity is
usually written as a number
with an appropriate unit.
There are 7 fundamental
quantities that we can
measure as shown in the table
below:
Base
Quantity
Unit
Unit
Symbol
Length
Meter
m
Mass
Kilogram
kg
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Time
Seconds
s
Electrical
Current
Ampede
A
Temperature
Kelvin
K
Amount of
Substance
Mole
mol
Luminous
Intensity
Candela
cd
International System of Units or SI is the
revised metric system. This system of
measurement uses prefixes together with the
base quantity.
Prefixes used with SI units.
-
(Multiply 1378 to 10^-3) =
1378 x 10^-3 = 1.378 km
3. 4936 mg = ____ kg ?
- (Subtract exponents of each
unit) = (-3) - (3) = -6
- (Multiply 4936 to 10^-6) =
0.004936 kg
4. 25 mL = ____ nL ?
- (Subtract exponents of each
unit) = (-3) - (-9) = 6
- (Multiply 25 to 10^6) =
25,000,000 nL
B. Density
- Intensive property
- Increases when temperature
decreases
- SI Unit: kg/m^3
- Density = Mass/Volume
Sample Density Problems:
Sample Conversion Problems with Steps:
1. 3.05 cm = ____ mm ?
- (Subtract exponents of each
unit) = (-2) - (-3) = 1
- (Multiply 3.05 to 10^1) =
3.05 x 10^1 = 30.5 mm
2. 1378 m = ____ km ?
- (Subtract exponents of each
unit) = (0) - (3) = -3
1. The density of mercury, the only
metal that is a liquid at room
temperature, is 13.6 g/mL. Calculate
the mass of 5.50 mL of the liquid.
- (Derive the formula for
density) d = m/v → m = d ⋅ v
- (Solve for the mass) m = 13.6
g/mL ⋅ 5.50 mL
- 74.8 g
2. A piece of metal ore weighs 9.00 g.
When a student places it into a
graduated cylinder containing water,
the liquid level rises from 20.25 mL
to 25.47 mL. What is the density of
the ore?
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-
(Solve for the volume) 25.47
mL - 20.25 mL = 5.22 mL
(Solve for the density) 9
g/5.22 mL
1.724 g/mL
b. the Kelvin scale
- K = °C + 273
- K = °90 +273
- 363°K
2.
3. A gold ingot weighs 4.50 lbs. If the
density of gold is 19.31 g/cm3, and
the length and width of the ingot are
11.0 cm and 2.00 cm respectively,
what is the height of the ingot? (1lbs
= 453.59 g)
- (Convert lbs to g) 4.50 x
453.39 = 2040.255
- v=l⋅w⋅h
- (Derive the formula for
density to find the height) d =
m/l⋅w⋅h → h = m/d⋅l⋅w
- (Solve for the height) h =
2041.16 /19.31 ⋅ 11 ⋅ 2
- 4.80 cm
Convert 67°F and 82°F to Celsius
a. 67°F
- °C = (°F - 32) ⋅ 5/9
- °C = (°67 - 32 ) ⋅ 5/9
- 19.44°C
b. 82°F
- °C = (°F - 32) ⋅ 5/9
- °C = (°82 - 32 ) ⋅ 5/9
- 27.78°C
3. What temperature is Celsius and
Fahrenheit equal?
- -40°C / -40°F
D. Scientific Notation
C. Temperature
- Average kinetic energy of particles in
an object/system.
Conversion Factors
Sample Temperature Problems:
1. The temperature difference between
the heated solution and the beaker is
90 C°. Express this temperature
difference on
a. the Fahrenheit scale
- °F = (9/5) °C+32
- °F = (9/5) °90+32
- 194°F
Sample Scientific Notation Problems:
1. 29,372,800,000,000,000
- 2.9372800000000000
- 2.93728 ⋅ 10^16
2. 7,832,940,000
- 7.832940000
- 7.83294 ⋅ 10^9
3. 144,570,000,000
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1.44570000000
1.4457 ⋅ 10^11
a. Obtained by using a
measuring tool (ie.
volume of water,
mark length)
b. Follows the rules of
significant figures
c. In measured numbers,
one of the significant
figures is known with
certainty. The last
significant figure is
only the best possible
estimate.
4. 345,289,400,000,000
- 3.45289400000000
- 3.452894 ⋅ 10^14
E. Significant Figures
Non-Significant
- Zeros to the left of the first non-zero
digit
- For no decimal point numbers, zeros
after the last non-zero digit
Significant
- Any digit that is not zero
- Zeros between non-zeros
- If a number is greater than 1, all
zeros written to the right of the
decimal point
○ Exact and Measured Numbers
- Exact numbers
a. Obtained by counting
and definition (ie. 3
girls, 16 ounces in a
pound, money)
b. Are considered to
have an unlimited
number of significant
figures and are never
used as a limited
factor in determining
the number of
significant figures in
the result of an
operation.
-
F. Accuracy and Precision
● Accuracy is how close a
measurement is to the true value of
the quantity that was measured.
● Precision refers to how closely two
or more measurements of the same
quantity agree with one another.
IV.
LESSON 4 - Atom and Atomic
Theory
Origin of The Word “Atom”
★ Democritus proposed that all matter
is composed of very small particles
called “atomos”, which means
indivisible or uncuttable.
★ The idea of “atomism” by
Democritus was not accepted by
many Greek philosophers including
Plato and Aristotle.
Measured numbers
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Key Dates in the Fundamental Laws of
Matter
The Sceptical Chymist (1661) - Robert
Boyle
- “Element” as the simplest
composition of matter and its atoms
can be combined to form different
“compounds”
Dephlogisticated Air (1774) - Joseph
Priestley
- Isolated oxygen gas, which Priestley
called “dephlogisticated air” by
heating mercury oxide (HgO)
Law of Conservation of Mass (1774) Antoine-Laurent Lavoisier
- States that in a chemical reaction, the
mass of the substances is equal to the
mass of the substances reacted
- Matter can be neither created nor
destroyed. In chemical reactions,
atoms are only rearranged, thus it
follows that the mass must be
conserved in a chemical reaction.
“Oxygen is involved in combustion and
respiration” (1778) - Antoine-Laurent
Lavoisier
Law of Definite Proportions (1788) Joseph-Louis Proust
-
-
Any sample of a given compound
will always be composed of the same
elements in the same proportion by
mass
Different samples of the same
compound always contain its
constituent elements in the same
proportion by mass.
Law of Multiple Proportions (1808) - John
Dalton
- For elements that can form different
compounds, the masses of the second
element that can combine with a
fixed mass of the first element are in
a ratio of small whole numbers
- If two elements combine to form
more than one compound, the masses
of one element that combine with a
fixed mass of the other element are
in ratios of small whole numbers.
Dalton’s Atomic Theory (1808) - John
Dalton
- the first atomic theory which was
based on experimental evidence from
previous scientific investigations.
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-
1. Elements are composed of
extremely small particles
called atoms.
2. All atoms of a given element
are identical, having the same
size, mass and chemical
properties. The atoms of one
element are different from the
atoms of all other elements.
3. Compounds are composed of
atoms of more than one
element. In any compound,
the ratio of the numbers of
atoms of any two of the
elements present is either an
integer or a simple fraction.
4. A chemical reaction involves
only the separation,
combination, or
rearrangement of atoms; it
does not result in their
creation or destruction.
This theory became one of the
foundations of modern chemistry.
on the Cathode Ray
Experiment.
-
qe/me = -1.76 x 10^11 C/kg
-
Plum-Pudding Model Atoms were made from a
positively charged substance
with negatively charged
electrons scattered about, like
raisins in a pudding.
-
Robert Millikan and Harvey
Fletcher (1909) conducted
the oil drop experiment to
determine the charge of an
electron.
-
1e = -1.602 × 10^−19 C
Atomic Structure
- Atoms actually are not indivisible;
that is, they are made up of even
smaller particles. These are called
the subatomic particles.
★ Electron Atomic Structure
- Joseph John Thomson
(1856-1940) started working
Radioactivity
★ Wilhelm Röntgen
- He noticed that cathode rays
caused glass and metals to
emit highly energetic
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radiation. He called it an
X-ray.
★ Antoine Becquerel
- He found that exposing
thickly wrapped photographic
plates to a certain uranium
compound caused them to
darken, even without the
stimulation of cathode rays.
★ Marie Curie
- Suggested the name
radioactivity to describe this
spontaneous emission of
particles and/or radiation
★ Nucleus and Proton Atomic
Structure
- Ernest Rutherford
(1871-1937) performed the
gold foil experiment with the
hypothesis that alpha rays
should pass through the plum
pudding-like structure of the
gold atoms.
- All of an atom’s positively
charged particles were
contained in the nucleus. The
negatively charged particles
were scattered outside the
nucleus
- The positively charged
particles in the nucleus are
called protons.
- mass of proton 1.6723 x
10^-24 g — about 1840 times
the mass of electron.
- mass of a nucleus constitutes
most of the mass of the entire
atom
- Atomic radius = 100 pm,
-
radius of an atomic nucleus =
5 x 10^-3 pm
★ Neutron Atomic Structure
-
Discovered by James
Chadwick, a British
physicist, in 1932.
Atomic Model
★ Bohr Model
- Niels Bohr (1913) proposed
an improvement on the
Atomic model. In his model,
he placed each electron in a
specific energy level.
- Electrons move in definite
orbits around the nucleus,
much like planets circle the
sun. These orbits, or energy
levels, are located at certain
distances from the nucleus.
★ Electron Cloud Model
- A modern theory that
describes the behavior of
electrons in atoms which
focuses on the probability
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-
-
V.
distribution of electrons
around the nucleus.
It is impossible to determine
the exact location of an
electron. The probable
location of an electron is
based on how much energy
the electron has.
Electrons with the lowest
energy are found in the
energy level closest to the
nucleus otherwise farther
from the nucleus.
Isotopes
- Atoms of the same element that have
different masses.
- Atoms that have the same number of
protons but different number of
neutrons
- Most elements have two or more
isotopes; they have the same atomic
number but different mass numbers.
LESSON 5 - Molecules and Ions
Atomic Number and Mass number
- Atomic number (Z) = number of
protons in the nucleus
- Mass number (A) = number of
protons + number of neutrons
- In a neutral atom the number of
protons is the same as the number of
electrons.
- Number of neutrons = (A – Z) or A –
number of protons
Molecule
- an aggregate of two or more atoms in
a definite arrangement held together
by covalent bonds.
Characteristics of Molecules
- A molecule can be a compound or an
element. It may contain atoms of the
same element (O2) or atoms of two
or more elements joined in a fixed
ratio (H2O).
- Molecules are electrically neutral.
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-
Molecules can be diatomic or
polyatomic. A diatomic molecule
contains only two atoms. Molecules
containing more than two atoms are
called polyatomic molecules.
-
An atom or a group of atoms that has
gained net positive or negative
charge.
Atoms become ions when it loses or
gains an electron. When an atom
loses one or more electrons, it gains
a positive charge and becomes a
cation.
A negatively charged ion on the
other hand is called anion. Anions
are formed when a neutral atom
gains one or more electrons.
Can also be classified as
monoatomic or polyatomic.
Monoatomic ions contain only one
atom. Polyatomic on the other hand
are ions containing more than one
atom.
Ion
-
-
-
Chemical Formulas:
- use chemical symbols to show the
composition of molecules and ionic
compounds.
- also shows the ratios in which the
atoms are combined in a compound.
- Can be categorized into Empirical
and Molecular Formula
a. Molecular Formula - shows
the exact number of atoms of
each element in a molecule. It
represents the actual formula
of a molecule.
-
b. Empirical Formula - shows
the simplest whole-number
ratio of the atoms present in a
compound. It is written by
reducing the subscripts in the
molecular formulas to the
smallest possible whole
numbers
The empirical formula can also be
the same as the molecular formula.
Writing Chemical Formulas of Ionic
Compounds
1. Criss-cross Method
Remember:
a. Ionic compounds are
electrically neutral.
b. The formulas of ionic
compounds are their
empirical formulas.
Chemical Nomenclature
- Binary Ionic compounds – consists
of two elements, a metal and a
nonmetal
How to name binary ionic compounds:
1. Cation = Metal cations take their
names from the elements
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2. Anion = The anion is named by
taking the first part of the element
name and adding the suffix “-ide.”
- The “-ide” ending is also
used for certain anion groups
containing different elements,
such as hydroxide (OH -) and
cyanide (CN -)
Examples:
NaCl - Sodium chloride
MgBr2 - Magnesium bromide
Al2O3 - aluminum oxide
Chemical Nomenclature: Metals with
variable charges
When naming compounds involving metals
with variable charges, the charge associated
with the metal cation in the compound must
be determined
-
following the name of the
metal. The Roman numeral
represents the charge of the
metal.
Name the nonmetal with the
ending -ide.
Ternary Ionic Compounds
- Ternary compounds - meaning
compounds consisting of three
elements.
- The most common types of ternary
ionic compounds consist of a
metallic cation and a polyatomic
anion
1. Classical System
- Assign the ending “-ous” to
the cation with fewer positive
charges and the ending “-ic”
to the cation with more
positive charges.
- The “-ous” and “-ic”
designations provide names
for only two different
elemental cations
2. Stock System
- Name the metal followed by
a Roman numeral in
parentheses immediately
Naming Ternary Ionic Compounds
1. Name the cation. Metal cations take
their names from the elements. If the
cation is a polyatomic ion, use the
name of ion. e.g. NH4+ (ammonium)
2. If the metal cation has a variable
charge, indicate the charge by using
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a roman numeral enclosed in
parentheses, or use the –ous and –ic
ending.
3. Use the name of the polyatomic
anion. e.g. OH- (hydroxide), CN(cyanide), PO43- (phosphate)
Naming Covalent Compounds
- Covalent compounds are composed
of non-metallic elements.
- Naming binary covalent compounds
is similar to naming binary ionic
compounds. We write the name of
the first element in the formula first,
and the second element is named by
adding -ide to the end of the
element’s name.
- In naming binary molecular
compounds, we use Greek prefixes
to denote the number of atoms of
each element present.
Remember:
1. The prefix “mono-” may be omitted
for the first element.
2. In oxides, the ending “a” in the
prefix may be omitted in some cases.
e.g. pentaoxide can be written as
pentoxide
VI. LESSON 6 - Stoichiometry
Atomic Mass
- (sometimes called atomic weight) is
the mass of the atom in atomic mass
units (amu).
Example:
1. Copper, a metal known since ancient
times, is used in electrical cables and
pennies, among other things. The
atomic masses of its two stable
isotopes, 63Cu (69.09 percent) and
65Cu (30.91 percent), are 62.93 amu
and 64.9278 amu, respectively.
Calculate the average atomic mass of
copper. The relative abundances are
given in parentheses.
Ave. Cu = (0.6909)(62.93 amu) +
(0.3091)(64.9278 amu) = 63.54 = 64 amu
Mole
- SI unit for the amount of substance.
- A mole is the amount of a substance
that contains as many elementary
entities (atoms, molecules, or other
particles)
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SPSTEM05 REVIEWER - STEM11E
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The mole is a convenient way to
express the quantity of substances
containing a very large number of
atoms.
Avogadro’s Number
- Named after the scientist Amedeo
Avogadro.
- The actual number of atoms in
exactly 12g of the carbon-12 isotope.
- Discovered through experiment
- Approximately equal to 6.022 x
10^23 (all elements will have this
number of atoms as long as they
have the same mass.)
Number of moles and mass
-
1 mole of carbon-12 atoms has a
mass of exactly 12 g and contains
6.022 x 10^23 atoms.
Molar mass (𝓜) is defined as the
mass of 1 mole of units (such as
atoms or molecules) of a substance.
Molar mass of an element (g) =
atomic mass (amu)
**NOTE: THIS FILE IS FOR REVIEWING PURPOSES ONLY. ANY ACT OF MALICIOUSLY USING IT AS CHEAT SHEET
DOES NOT CORRELATES TO THE PUBLISHER.**