STEM 11 - SENIOR HIGH SCHOOL Lesson 1: Matter and its Properties Lesson 2: Measurements Lesson 3: Atoms, Molecules, and Ions Lesson 4: Mole Concept

1. GENERAL
CHEMISTRY
1
SENIOR
HIGH
SCHOOL

2. Jens Martensson 2
General Chemistry 1
Subject Description
Composition, structure, and properties of
matter; quantitative principles, kinetics, and
energetics of transformations of matter; and
fundamental concepts of organic chemistry
SCIENCE
TECHNOLOGY
ENGINEERING &
MATHEMATICS
SPECIALIZED SUBJECT
Grade Level: 11
Semester: 1st/2nd
Hours/Semester: 80 Hrs.

3. Jens Martensson 3
Contents
UNIT 1
INTRODUCTIONTO CHEMISTRY
UNIT 2
CHEMICALCALCULATION & REACTION
UNIT 3
THE GASEOUS STATE OF MATTER
UNIT 4
ATOMICAND MOLECULARSTRUCTURES
SCIENCE
TECHNOLOGY
ENGINEERING &
MATHEMATICS
SPECIALIZED SUBJECT
STEM-ACADEMIC
Grade Level: 11
Semester: 1st/2nd
Hours/Semester: 80 Hrs.

4. Jens Martensson
UNIT 1
INTRODUCTION TO
CHEMISTRY
Lesson 1: Matter and Its Properties
Lesson 2: Measurements
Lesson 3:Atoms, Molecules, and Ions
Lesson 4: Mole Concept
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LESSON 1: MATTER AND ITS
PROPERTIESOBJECTIVES OF THE DAY
• I will be able to describe the particulate
nature of the different forms of matter;
• I will be able to classify the properties of
matter;
• I will be able to differentiate pure
substance and mixtures; elements and
compounds; homogeneous and
heterogeneous mixtures;
1 2
3 4
5 6
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LESSON 1: MATTER AND ITS
PROPERTIESOBJECTIVES OF THE DAY
• I will be able to recognize the formulas of
some common substances;
• I will be able to discuss methods to
separate the components of a mixtures;
and
• I will be able to recognize chemical
substances present in some consumer
products
1 2
3 4
5 6
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MATTER
Activity 1: What is Matter?

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Matter is anything that has mass and occupies space. Everything on earth has mass and takes
up space.

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PARTICLES COMPOSING MATTER
These are the smallest unit
of matter that can’t be
broken down chemically.
These are groups of two or
more atoms that are
chemically bonded.
These are particles that
have gained or lost one or
more of their valence
electrons.
ATOMS MOLECULES IONS

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STATES OF MATTER
SOLID LIQUID GAS
Activity 2: Table Completion

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PLASMA
THE 4TH STATE OF MATTER
• It is a hot ionized gas consisting of
approximately equal numbers of
positively charged ions and negatively
charged electrons.
• The characteristics of plasmas are
significantly different from those of
ordinary neutral gases so
that plasmas are considered a distinct
"fourth state of matter."
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BOSE-EISTEIN CONDENSATE
THE 5TH STATE OF MATTER
• It is a state of matter in which separate
atoms or subatomic particles, cooled to
near absolute zero.
• When they reach that temperature the
atoms are hardly moving relative to each
other; they have almost no free energy to
do so. At that point, the atoms begin to
clump together, and enter the same
energy states.
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PROPERTIES OF MATTER
PHYSICAL
PROPERTIES
These can be measured
and observed without
changing the composition
of the substance.
CHEMICAL
PROPERTIES
These are the ability of a
substance to react with
other substances such as
air, water, and base.
INTENSIVE
PROPERTIES
It does not depend on
the size or amount of
the sample.
EXTENSIVE
PROPERTIES
These can be
affected by the size
and amount of
samples.
According to changed involved
during measurements of the
property.
According to dependence on
amount of matter

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PHYSICAL PROPERTIES
INTENSIVE PHYSICAL PROPERTIES EXTENSIVE PHYSICAL PROPERTIES
Color Melting Point Density
Solubility Conductivity Malleability
Luster Viscosity Boiling Point
Temperature Odor
Mass
Volume
Length
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CHEMICAL PROPERTIES
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CHEMICAL PROPERTIES DESCRIPTION
1. Combustibility Whether the substance undergoes combustion or not
2. Stability Whether the substance can be easily decomposed or not
3. Reactivity Whether it reacts with acids, bases, and oxygen, gas or not
4. Relative Activity Whether the material is more active or less active than other members
of its chemical family
5. Ionization Whether it will break into charged particles when in solution with water
or not.
6. Toxicity Whether substance can damage an organism or not.

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Activity 3
Group the characteristics of the give substance according to their physical (extensive or
intensive) or chemical properties.
CHARACTERISTICS OF SOME
SUBSTANCES
PHYSICAL PROPERTIES CHEMICAL
PROPERTIESINTENSIVE EXTENSIVE
1. The water in the container has a
volume of 100 mL and a mass of
99.8 g. It is colorless, and
tasteless. It has a density of
0.998g/mL, boils at 100 degrees
Celsius, and freezes at 0 degree
Celsius. It does not burn. It causes
Iron to rust.
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Activity 3
Group the characteristics of the give substance according to their physical (extensive or
intensive) or chemical properties.
CHARACTERISTICS OF SOME
SUBSTANCES
PHYSICAL PROPERTIES CHEMICAL
PROPERTIESINTENSIVE EXTENSIVE
2. NaCl with a mass of 37.9 g is
colorless, odorless, and salty solid
crystals. It has melting point of 801
degree Celsius. When dissolved in 100
mL water, it conducts electricity. It reacts
with silver nitrate to form a white
precipitates. It also react with water to
form chlorine gas, hydrogen gas, and
sodium hydroxide.
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MATTER
It is a matter that has a definite
composition and distinct properties
These are composed of two or more
substances combined physically in various
composition
It is the simplest
form of matter since
it composed of only
one kind of atom.
It contains two or
more kinds of atom
chemically combined
in definite proportion
by mass
It is a solid, liquid, or
gaseous mixture that has
the same proportions of its
components throughout any
given sample.
It is a mixture whose
composition varies from
one position to another
within the sample.
PURE
SUBSTANCE
MIXTURES
ELEMEN
T
COMPOU
ND
HOMOGENEOUS
MIXTURE
HETEROGENOU
S MIXTURE

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Activity 4: Pure Substance or Mixture?
1. TABLE SUGAR 2. TABLE SALT
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PURE SUBSTANCE OR MIXTURE?
3. IODIZED SALT 4. DISTILLED WATER
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PURE SUBSTANCE OR MIXTURE?
4. SOFTDRINKS 5. OXYGEN GAS (TANK)
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PURE SUBSTANCE OR MIXTURE?
6. BROWN SUGAR 7. HUMAN BREATH
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Activity 5: HOMOGENEOUS OR HETEROGENEOUS?
1. RUBBING ALCOHOL 2. WATER &OIL
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HOMOGENEOUS OR HETEROGENEOUS?
3. SALT & PEPPER 4. CARBONATED SOFTDRINKS
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HOMOGENEOUS OR HETEROGENEOUS?
5. HUMAN BREATH
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SEPERATING MIXTURES
Chemist separate mixtures by using different methods.
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SEPERATING MIXTURES
Chemist separate mixtures by using different methods.
1.Filtration is a process of separating the components of a suspension
2.In Decantation the solid particles are allowed to settled first at the
bottom and later, the liquid which is called supernatant is poured into
another container leaving behind solid particle.
3.Evaporation is the process of converting liquid to gas, is useful in
sorting mixtures such as salt solution.
4.Distillation is a process of separating a homogeneous mixture
composed of two substances with different boiling points.
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SEPERATING MIXTURES
Chemist separate mixtures by using different methods.
5. Magnetic Separation is the process of separating elemental metals
from other particles in a mixture.
6. Melting is a process that can be used in extricating mixture that
contain two substances with different melting points.
7. Sublimation is a process of changing solid to gas without passing
through the liquid state.
8. In Centrifugation, the mixture is poured into a special tube in the
centrifuge apparatus, and is allowed to spin using centrifugal force. The
spinning motion forces the sediments to settle at the bottom. The liquid
can be poured off from the solid particles.
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9. Chromatography is another method of separating complex mixtures. It has
various methods that can be used in separating mixture such as paper
chromatography, which makes used of an adsorbent (filter paper or chromatogram
paper), then separation depends upon the solubility of each component in the
solvent.
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PURE SUBSTANCES & MIXTURES IN CONSUMER
PRODUCTS
CONSUMER PRODUCT
• It is any item often bought for consumption.
Convenience Product – those that appeal to a large segment
of the market or those that are routinely bought.
1. Household Cleaning
2. Personal Care Product
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PURE SUBSTANCES & MIXTURES IN CONSUMER
PRODUCTS
HOUSEHOLD CLEANING MATERIAL
• The most commonly used cleaning products are bleach,
soaps, and detergents. These products have different
compositions, specific uses, precautions for use, and costs.
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PURE SUBSTANCES & MIXTURES IN CONSUMER
PRODUCTS
HOUSEHOLD CLEANING MATERIAL
• Bleach helps clean and whiten surfaces by generally lowering
the stability of the chemical bonds in stain molecules.
• It can convert dirt into particles that can be easily washed
away in conjunction with use of detergents.
• NaOCl (Sodium Hypochlorite) and H2O2 (Hydrogen
Peroxide) are most common bleaching agents that are strong
oxidizers; they can burn then skin and eyes especially if used
in concentrated form.
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PURE SUBSTANCES & MIXTURES IN CONSUMER
PRODUCTS
HOUSEHOLD CLEANING MATERIAL
• Soap and Detergent are mixture of surfactants, water
softeners, stain removers, enzymes and perfumes, among
others.
• Surfactants render soaps and detergents capable of lowering
the surface tension of water, which allows them to wet the
surface to be cleaned. They also loosen and disperse water-
insoluble solids making them washable with water.
• Soap and Detergents are generally not toxic and severely
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PURE SUBSTANCES & MIXTURES IN CONSUMER
PRODUCTS
PERSONAL CARE PRODUCTS
• Personal Care product constitute a diverse group of materials
that improve the overall appearance of a person. These
products are used to generally cleanse and beautify.
Examples of highly demanded personal care products are
makeup, lotions, and toothpastes.
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LESSON 2: MEASUREMENTS
OBJECTIVES OF THE DAY
• I will be able to describe the need
for measurement;
• I will be able to carry out simple
measurements of length, volume,
and mass; and
• I will be able to differentiate the
accuracy and the precision of a
measurement
1 2
3 4
5 6
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Keywords for the concepts to be learned;
a.Measurements
b.Unit of Measurements
c.Accuracy
d.Precision
e.Significant figures
f. Errors
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Measurements
• The study of matter requires a certain
degree of measurements, a process of
determining the extent of the dimensions,
quantity, or extent of something.
• Questions such as “How much….?” ,“How
long…?” and “How many…?” simple
cannot be answered without resorting to
measurement.
Q1. Can you cite some situations in daily life
where a measurement is important?
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Units of Measurements
• The most convenient system
of units is the International
System of Units (SI).
• This system is the modern
versions of metric system.
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Units of Measurements
The name of the fractional parts and the multiples of the base units are
constructed by adding prefixes. These prefixes, shown in table, indicate
the size of the unit relative to the base unit.
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Uncertainty in Measurements
• A measured quantity contains some digits that are exactly
known and one digit that is estimated. The estimated digit
produces uncertainty in measurements.
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Random Error and Systematic Error
• RANDOM ERROR (indeterminate error) is the uncertainty
that arises from a scale reading which results from the
uncontrolled variables in the measurement.
• It causes one measurement to differ slightly from the next. It
comes from unpredictable changes during an experiment.
Examples
a. When weighing yourself on a scale, you position yourself
slightly different each time.
b. Measuring your height is affected by minor posture changes.
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Random Error and Systematic Error
• SYSTEMATIC ERROR (determinate error) is the uncertainty
that may come from a flaw in the equipment used or design of
an experiment. These error are usually caused by measuring
instruments that are incorrect calibrated or are used
incorrect.
• Examples
a. A worn out instrument
b. An incorrectly calibrated or tared instrument
c. A person consistently take an incorrect measurements
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Precision and Accuracy
• Precision is the consistency of a result. If you measure a
quantity several times and the values agrees closely with one
another, then your measurement is precise.; however, if the
values varied widely, then it is imprecise.
• Accuracy is determined when a certain quantitative value is
relatively close to the “true value”
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Activity 6
• Each dot is the result of a measurement whose value is indicated in
the horizontal (or x-) axis. The plot presents the results of six
measurements of the weight of a pebble whose true weight is 8.0 g.
• Determine whether each measurement is accurate or inaccurate, and
precise or imprecise.
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Activity 7
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Scientific Notation
• It is a simple way to write or keep track of very large or very
small numbers without having to deal with a lot of zeros.
• It provides a convenient way of recording results and doing
calculations.
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Activity 8
1. 0.012345698632
2. 1 230 945
3. 87 576 788 432 234 543
4. 0.O6OO789653
5. 11 987
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Significant Figures
• Significant figures are the digits in any measurement that
are known certainty with an additional one digit which is
uncertain.
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RULES MEASURED NUMBERS NUMBER OF SIGNIFICANT
1. All nonzero digits are significant.
247 3
2. Zeroes between nonzero digits
are significant. 20303 5
3. Zeroes to the left of the first
nonzero digits are NOT significant 0.0200 3

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Significant Figures
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RULES MEASURED NUMBERS NUMBER OF SIGNIFICANT
4. If the number is less than 1, then
only the zeros at the end of the
number and the zero between
nonzero digits are significant.
0.003560 4
5. If the number is greater than 1,
then all the zeros written to the right
of the decimal point are significant. 35600.00 7

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Activity 9
Give the number of significant figures for each of the following measurements.
1. 2 365 mm
2. 309 cm
3. 5.030 g/mL
4. 0.0670 g
5. 3.60 x 10-4
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Activity 10
Give the number of significant figures for each of the following measurements.
1. 0.476 kg
2. 89.7808 ft
3. 0.430 mg
4. 60.0 min
5. 1 x 107
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Rules for Significant Figures in Fundamental Operations
• In addition and subtraction, the answer must have the same
number of decimal places as the measured number with the
least number of decimal places.
• In multiplication and division, the answer must have the
same number of significant figures as the measured number
with the lowest number of significant figures.
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Activity 11
Perform the following operations and write the answers in the proper
number of significant figures.
1. 4.87 m + 36.578 m + 4.34 m
2. 8.9 mL ÷ 45 mL
3. 68.980 cm – 67.16 cm
4. 45.00 ft. x 3.00 ft.
5.14.4 g + 6.0 g
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Rules in Rounding Off
• Oftentimes, the answers to
computations contain too many
insignificant digits. Hence it
becomes necessary to round off
numbers to attain the insignificant
figures. Rounding off, therefore, is
the process of removing,
insignificant digits from calculated
number.
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57. Jens Martensson
Rules in Rounding Off
• The following rules should be applied to round off values to
the correct number of digits.
1. For a series of calculations, carry extra digits through to the
final result, then round off.
2. If the first digit to be deleted is….
a. 5 or greater, the last retained figure is increased by one
b. 4 or less, the last retained figure is retained.
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Conversion of Units (Dimensional Analysis)
• Dimensional Analysis is a process in which a conversion
factor written in a form of ratio is used to change units given in
the data to the units desired.
• The following are steps to be followed in doing dimensional
analysis.
a. Write the unknown quantity that is sought, including the
units.
b. Write all known conversion factors needed
c. Begin with what is known and then multiply it by the
identified conversion factor, cancelling similar units to get the
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METRIC AND ENGLISH CONVERSIONS
QUANTITY METRIC ENGLISH CONVERSION
MASS g, Kg lb, oz 1lb = 454 g
1kg = 2.2 lb
1 oz = 28.35g
LENGTH cm, m, km in, ft, mi, Å 1 in = 2.54cm
1 m = 39.37 in
1 ft = 12 in
1 mi = 1.609 km
1 km = 0.62137 mi
1 Å = 10-10 m
VOLUM
E
mL, L qt, pints, cups,
tsp, tbsp, fl oz,
gal
1 qt =946 mL
1 L = 1.057 qt
1 L = 2.12 pints
1 L = 4.23 cups
1 tsp = 4.93 mL
1tbsp = 14.79 mL
1 fl oz = 29.06 mL
1 gal = 3.79 L 59

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Activity 12
Sample Problems!
1.The lemon juice drink contains 500.0 mg of vitamin C.
Express the vitamin C content in grams.
2. A fitness drink measures 0.300 L. Express the volume
in L
3. Calculate the number of centimeters in 53.5 inches.
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Density Measurement
• Density measurement is one of the common measurements
done in the laboratory. It involves getting Mass, Volume and
Temperature of an Object.
A. Mass is the quantity of matter in the object. It is determined
by weighing the object, using balance. The SI basic unit of
mass is the Kilogram, but the gram is more convenient to
use.
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Density Measurement
B. Volume is the amount of space occupied by a substance. In
liquids, the volume can be determined using a graduated
cylinder, while solids, the volume can be determined by two
methods.
1. For regularly shaped-solids, the volume formula for the
particular shaped is used.
Some formula that may be used are the following:
Rectangular solid = L x W x H Cylindrical Solid = πr2h
Cubic solid = S x S x S Spherical solid = 4/3 πr3
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Density Measurement
2. For irregularly-shaped solids, the water displacement method
is used.
C. Temperature tells how hot or cold an object is. It is
commonly marked either by oC (Celsius) or o F (Fahrenheit);
although the SI basic unit for temperature is the K (Kelvin).
To convert oC or oF to K, the following are used.
• K = oC + 273.15
• K = (oF + 459.67) x 5/9
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Density Measurement
DENSITY is the ratio of the mass of an object to the volume it
occupies.
DENSITY =
𝑴𝑨𝑺𝑺
𝑽𝑶𝑳𝑼𝑴𝑬
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Activity 13
Sample Problems
1.A sample amount of sugar has a mass of 250.0 g and a
volume of 157.3 cm3. What is its density in grams per cubic
centimeter?
2. Gold metal has a density of 19.3 g/cm3. What is the volume
in cubic centimeter of a 500.0 g bar of gold metal?
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Activity 14
Sample Problem
1. The volume and the mass of two objects (A & B) have been
obtained in order to determine their densities, respectively.
Identify which object is denser.
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OBJECT METHOD USED FOR DETERMINING THE VOLUME MASS
A By measuring its dimension
L = 2.0 cm, W = 2.5 cm H = 15 cm
90.0 g
B By water displacement method:
Final Volume (Water + object) = 100 mL
Initial Volume (Water) = 80.0 mL
65.0 g

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Lesson 3: Atoms, Molecules and Ions
OBJECTIVES OF THE DAY!
• I will be able to describe and discuss the
basic laws of chemical change;
• I will be able to discuss how Dalton’s Atomic
Theory could explain the basic laws of
chemical changes;
• I will be able to give the information
provided by the atomic number and mass
number of an atom and its isotopes
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68. Jens Martensson
Lesson 3: Atoms, Molecules and Ions
OBJECTIVES OF THE DAY!
• I will be able to differentiate atoms,
molecules, and ions;
• I will be able to write the chemical formula of
some molecules;
• I will be able to differentiate a molecular
formula and an empirical formula; and
• I will be able to give the name of a
compound, given its chemical formula.
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Keywords
a. Law of Conservation of Matter h. Law of Definite
Proportion
b. Law of Multiple Proportion i. Dalton’s Atomic
Theory
c. Atomic number j. Mass number
d. Isotope k. Atom
e. Molecule l. Ion
f. Chemical formula m. Molecular formula
g. Empirical formula
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LAWS OF CHEMICAL CHANGE
• These laws were inferred from several
experiments conducted during the 18th century
using a balance for the measurements:
1. Law of Conservation of Mass
2. Law of Definite Proportion
3. Law of Multiple Proportion
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A. Law of Conservation of Mass
• ANTOINE LAVOISIER, a brilliant
French chemist, formulated this law
by describing one of his experiments
involving mercuric oxide.
• He placed a small amount of
mercuric oxide, a red solid, inside a
retort and sealed the vessel tightly.
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A. Law of Conservation of Mass
• He weighed the system, and then subjected it to
high temperature.
• During the heating, the red solid turned into a
silvery liquid. This observation indicated that a
chemical reaction took place.
• After which, the setup was cooled and then
weighed. The weight of the system was found to
be the same as before heating.
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A. Law of Conservation of Mass
• In a chemical reaction, no change in mass takes
place. The total mass of the products is equal to
the total mass of the reactant.
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B. Law of Definite Proportion
• A compound always contains the same constituent elements
in a fixed or definite proportion by mass.
• If water samples coming from different sources are analyzed,
all the samples will contain the same ratio by mass of
hydrogen to oxygen.
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Sample Problems
• A pure sample of Sodium Flouride (NaF) contains 35g of
Sodium. How many grams of Flourine are present in this
sample?
• If there are 42g of H in a sample of pure Methane (CH4), How
many grams of Carbon are present?
• If there are 19g of oxygen in a sample of Aluminum Oxide
(Al2O3), How many grams of Aluminum are present?
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C. Law of Multiple Proportions
• If two elements can combine to form more than one
compound, the masses of one element that will combine with
a fixed mass of the other element are in a ratio of small whole
numbers.
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Dalton’s Atomic Theory
• In 1808, John Dalton published his
book A New System of Chemical
Philosophy, where he proposed an
atomic theory of matter that can
explain chemical observations as
predicted by the three fundamental
laws.
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Dalton’s Atomic Theory
The atomic theory comprised the following postulates:
1. Matter is made up of extremely small
indivisible particles called atoms.
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Dalton’s Atomic Theory
The atomic theory comprised the following postulates:
2. Atoms of the same element are
identical, and are different from those of
other elements.
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Dalton’s Atomic Theory
The atomic theory comprised the following postulates:
3. Compounds are composed of atoms
of more than one element, combined in
definite ratios with whole number
values.
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Dalton’s Atomic Theory
The atomic theory comprised the following postulates:
4. During a chemical reaction, atoms
combine, separate, or rearrange. No
atoms are created and no atoms
disappear.
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• During the time of Dalton, the atom
was believed to be the smallest
particle comprising substances.
However, before the end of the 19th
century, experiments provided proof of
the existence of smaller particles
within the atom.
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Activity 15
• Recall the particles contained in an atom (or the subatomic particles) and differentiate the
particles in terms of location, charge, and relative mass by filling up the following table:
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Activity 15
• Recall the particles contained in an atom (or the subatomic particles) and differentiate the
particles in terms of location, charge, and relative mass by filling up the following table:
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Atomic Number and Mass Number
• An atom of an element may be represented in a certain
configuration that includes its atomic number (Z) and Mass
number (A), written as the left superscript and left subscript,
respectively of the element symbol.
4
2
He
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mass number (A)
atomic number (Z)
Symbol of Element

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Atomic Number and Mass Number
• The atomic number of an element represents the number of
protons in its nucleus. Because an atom as a whole is
electrically neutral, the atomic number also specifies the
number of electron present.
ATOMIC NUMBER = NUMBER OF PROTONS = NUMBER OF
ELECTRONS
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Atomic Number and Mass Number
• The mass number of an atom is the sum of the number of
protons and neutrons in its nucleus. Thus, the mass number
gives the number of subatomic particles present in the
nucleus.
MASS NUMBER = NUMBER OF PROTONS + NUMBER OF
NEUTRONS
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Activity 16
COMPLETE THE TABLE BELOW
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Isotopes are toms of an element having the same atomic
number but different mass number.
The existence of isotopes was shown by mass spectroscopy
experiments, wherein elements were found to be composed of
several types of atoms, each with different masses.
a. The atomic number identifies an element. The atoms of
isotopes of an element have the same number of protons and
electrons.
b. The atoms of isotopes of an element differ in the number of
neutrons.
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Atoms, Molecules and Ions
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Atoms, Ions and Molecules
• Of all the elements, only six exist as single atoms, namely
Helium, Neon, Argon, Krypton, Xenon and Radon. Most
matters are composed of ions formed from atoms.
• A molecule is a combination of at least two atoms in a definite
proportion, bound together by covalent bonds.
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Ions
• When a neutral atom gain or loses one or more electrons, it
becomes an electrically charged particles called ion.
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Ions
• Metals tend to lose electrons and become positively charged
cations. Nonmetals, on the other hand, gain electrons and
become negatively charged anions. The number of electron
lost or gained is the charged number.
• Ions can be made up of only one atom (monoatomic) or more
than one type of atom (polyatomic).
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Naming Monoatomic Ions
Monoatomic ions are named based on the element.
a. For cations, the name of the element is unchanged.
• If an element can form two ions of different charges, the
name, which is usually derived from its Latin name, is
modified by the suffix –ic for the ion with the higher charge,
and –ous for that with the lower charge.
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Naming Monoatomic Ions
b. The monoatomic anions are named by attaching the suffix –
ide to the first few letters (root) of nonmetal name,
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Activity 17.1
Name the following cations below.
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Activity 17.2
Name the following anions below.
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• Several anions are polyatomic and are named based
on the atomic constituents and the suffix – ide.
• The most common examples are:
a. OH- – hydroxide ion
b. CN- – cyanide ion
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• A number of polyatomic anions containing oxygen
atoms are named based on the root word of the central
(or non-oxygen) atom and the suffix –ate for the one
with more oxygen atoms and –ite for the one with less
oxygen atom.
• Some anions have common names ending with the
suffix –ate.
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Chemical Formula
• The composition of a molecule or an ion can be represented
by a chemical formula. The formula consists of the symbols of
the atoms making up the molecule. If there is more than one
atom present, a numerical subscript is used.
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Chemical Formula
• There are two types of chemical formula - the molecular
formula and empirical formula.
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Chemical Formula
• Molecular Formula indicates the actual number of each
element in a compound.
• Emperical Formula is the simplest chemical formula. It only
shows relative ratio between the number of atoms of the
different elements present in the compound.
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Activity 18
Write the empirical formula of the following molecules.
1. C2H4O2
2. C8H12N4
3. C8H10
4. P4O10
5. PH3
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Naming Compounds
A. IONIC COMPOUNDS (cation and anion)
1. For Binary Compounds, metal cations take their
names from the elements, while the anion take the
first part of the name of element , and add the suffix
–ide end.
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CATION ANION COMPOUND NAME OF
COMPOUND
Na+ O-2 Na2O Sodium oxide
Mg+2 N-3 Mg3N2 Magnesium nitride
Al+3 O-2 Al203 Aluminum Oxide

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Naming Compounds
2. For Ternary Compounds, the cation goes first in its
name before the polyatomic ion which usually ends with
–ite or -ate
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CATION ANION COMPOUND NAME OF
COMPOUND
Na+ NO3
-1 NaNO3 Sodium Nitrate
Na+ NO2
-1 NaNO2 Sodium Nitrite

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Naming Compounds
3. For compounds containing a metallic ion of variable
charge, either the classical method or the stock method
of naming may be used.
• In the classical method, the name of metallic ions ends
in –ous (for lower charge) and –ic (for higher
charge)
• In the stock method, the metal is named first followed
by the value of the charge written in roman numeral
(enclosed in parenthesis)
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Naming Compounds
B. MOLECULAR COMPOUNDS (TWO NONMETALS)
• For one pair of elements that form several different
compounds, Greek prefixes are used to determine the number
of each element in the compound. For the first element, the
prefixes “mono” is omitted.
Examples
CO – carbon monoxide
CO2 – carbon dioxide
N2O4 – dinitrogen tetraoxide
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Naming Compounds
B. MOLECULAR COMPOUNDS (TWO NONMETALS)
• For binary compounds, place the name of the first element;
then follow it with the second element. The second element is
named by adding –ide to the root of the element name.
Examples:
a. HCl – hydrogen chloride
b. HBr – hydrogen bromide
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Naming Compounds
B. MOLECULAR COMPOUNDS (TWO NONMETALS)
• For binary compounds considered as acids, use the prefix
hydro- followed it with second element. The second element
is named by adding –ide to the root of the element name.
• Examples
a. HCl – hydrogenchloric acid
b. HBr – hydrobromic acid
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Naming Compounds
B. MOLECULAR COMPOUNDS (TWO NONMETALS)
• Oxy-acids, those that contain hydrogen, oxygen and another
element, is named in two ways –
a. For anions ending with –ate, change –ate to –ic; then, follow
it with the word acid.
b. For anions ending with –ite, change –ite to –ous; then follow
it with the word acid.
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116. Jens Martensson
Lesson 4: Mole Concept
OOTD: OBJECTIVES OF THE DAY
• At the end of this lesson, the students must be able to
• explain the meaning of average atomic mass
• define a mole;
• illustrate Avogadro’s number with examples;
• determine the molar mass of elements and
compounds;
• calculate the mass of a given number of moles of an
element or compound, or vice versa; and
• calculate the mass of a given number of particles of
an element or compound, or vice versa.
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PERFORMANCE TASK
THIRD QUARTER – GENERAL CHEMISTRY 1
• History of the Atomic Structure
1. J.J Thomson’s Plum Pudding Model
2. John Dalton’s Billiard Ball Model
3. E. Rutherford’s Nuclear Model
4. Niels Bohr’s Planetary Model
5. Schrodinger & Heisenberg Quantum Mechanical Model
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• Ms. Lilia sells shelled peanuts in a store.
But she meets customers asking for 150
peanuts, another for 750 peanuts, and
another for 2,000 peanuts.
• Obviously, it will take Ms. Lilia a very
long time to count the peanuts. What
would be another way to count them?
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Atomic Mass and Atomic Mass Unit
Question: Is it possible to use the same procedure to count atoms. Why
or why not?• Whether it is peanuts or mongo beans or candies or atoms, the
procedure should be the same.
• The problem, however, is atoms are very, very small and it is not
possible to see them and count them individually to get the average
mass.
• We need to look for another way to get the average mass of the atom.
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Atomic Mass and Atomic Mass Unit
• Experiments have shown that atoms have different masses
relative to one another.
• For example, a Mg atom is experimentally reported to be
twice as heavy as a carbon atom; a silicon atom is twice the
mass of a nitrogen atom.
• It is possible to make a relative scale if one atom is chosen
as the reference or standard atom against which the masses
of the other atoms are measured.
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Atomic Mass and Atomic Mass Unit
• By international agreement, the
reference atom chosen is the C-12
isotope which contains six protons
and six neutrons.
• By definition, one atom of C-12 has a
mass of exactly 12 atomic mass units
(amu).
• One amu, therefore, is one-twelfth
(1/12) the mass of a C-12 atom.
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Atomic Mass and Atomic Mass Unit
• Example
• The atomic mass of Cu-63 is 62.93 amu. This means that
relative to C-12, one atom of Cu-63 is 62.93/12 or 5.244 times
the mass of a C-12 atom.
Try This!
• One atom of Se-77 is 6.410 times as heavy as an atom of C-
12. What is the atomic mass of Se-77?
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Average Atomic Mass
• The atomic mass of the atoms of an element is the average atomic
masses of the naturally occurring isotopes of this element.
• The periodic table provides the average atomic mass which takes into
account the different isotopes of an element and their relative
abundances.
NOTE: It is not a simple average that is taken but a weighted average
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For example
• The average atomic mass of Oxygen is 15.999, not 16.00
• The 15.999 is calculated by considering the naturally-occurring isotopes of Oxygen, namely
Oxygen-16, Oxygeny-17 and Oxygen-18

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Average Atomic Mass
• Average atomic masses are obtained by multiplying the mass of an
isotopes by its fractional abundance, as shown as follows.
• Isotopes of elements occur in different abundances. Some are more
abundant than others.
1. Chlorine has two isotopes. The natural abundance of Cl-35 is 75%
while that of Cl-37 is 25%. This means that if you have 100 atoms of
chlorine, 75 of them will be Cl-35 and 25 of them will be Cl-37.
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ELEMENT MASS NUMBER ISOTOPIC MASS % ABUNDANCE AVERAGE
ATOMIC MASS
Oxygen
16 15.9949 u 99.76%
15.999 u17 16.9991 u 0.04%
18 17.9992 u 0.20%

126. Jens Martensson
2. Magnesium, on the other hand, has three isotopes with varying
abundances: Mg-24,Mg-25, and Mg-26, 11.01 have 78.99%, 10.00%,
and 11.01% abundance, respectively.
3. For carbon, the natural abundance of C-12 is 98.90% while that of C-
13 is 1.10%. The atomic mass of C-13 has been determined to be
13.00335 amu while that of C-12 is exactly 12 amu.
• Relative Atomic Mass is the ratio of the average atomic mass of an
atom to one atomic mass unit (amu) Hence, its value is similar with
average atomic mass, except that it has no unit.
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Activity 19
1. Copper has two stable isotopes with the following masses and %
abundances: Cu-63 (62.93 amu, 69.09% abundance) and Cu-65
(64.9278 amu, 30.91% abundance). Calculate the average atomic
mass of copper.
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The Avogadro’s Number
• In the SI system, the mole (mole) is defined as the
amount of substance containing the same number
of particles as there are atoms in exactly 12 g of
carbon-12 isotope.
• One mole of a substance is equivalent to the
Avogdro’s number of particles 6.02 x 1023
• This number is so-named in honor of the Italian
scientist, Amadeo Avogadro
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The Avogadro’s Number
• Thus, based from the definition. It can be said that
a. One mole of an element is numerically equal to its
atomic mass unit.
b. One mole of an element contains 6.02 x 1023
atoms
c. One mole of molecular compound contains 6.02 x
1023 molecules
d. One mole of ionic compound contains 6.02 x 1023
cations and 6.02 x 1023 anions
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Molar Mass
• The molar mass of a compound (molecular or ionic) is the
mass in grams of one mole of a substance. It is numerically
equal to the sum of the masses of the elements (in amu) that
make up the compound.
• The molar mass is obtained by multiplying the number of
atoms by the atomic mass of each element, and getting the
sum. The unit for molar mass is g/mol.
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131. Jens Martensson
Molar Mass
• The molar mass of a compound (molecular or ionic) is the
mass in grams of one mole of a substance. It is numerically
equal to the sum of the masses of the elements (in amu) that
make up the compound.
• The molar mass is obtained by multiplying the number of
atoms by the atomic mass of each element, and getting the
sum. The unit for molar mass is g/mol.
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Activity 20
• Calculate the molar mass of the following compounds
1. C3H5N3O9
2. (NH2)2 CO
3. Hg(OCN)2
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Formula Mass and Molecular Mass
• Formula Mass is used for compounds that exists as ions, such as
NaCl. It is expressed in amu or u, and is numerically equal to the molar
mass expressed in grams per mole of a substance.
• Molecular Mass is used for compounds that exist as molecules, such
as water (H2O) It is numerically equal to the molar mass and has a unit
amu.
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IONIC COMPOUND COMMON NAME MOLAR MASS FORMULA MASS
NaCl Table Salt 58g/mol 58 amu
CaO Quicklime 56g/mole 56 amu
MOLECULE COMMON NAME MOLAR MASS MOLECULAR MASS
CO2 Dry Ice 44g/mol 44 amu
C12 H22 O11 Dextrose 342g/mole 342 amu

134. Jens Martensson
Calculation Involving Formulas
• The Avogadro’s number and molar mass are important to enable
conversions between mass and moles of atoms or molecules, ion and
vice versa. The following are the conversion factors that can be used in
calculations involving formulas.
• Where X represents the symbol of atoms, ions, or the formula of the
compound.
1 𝑚𝑜𝑙 𝑜𝑓 𝑋
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑋
and
1 𝑚𝑜𝑙 𝑜𝑓 𝑋
6.23 𝑥 1023 𝑎𝑡𝑜𝑚𝑠 (𝑜𝑟 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒𝑠 𝑜𝑟 𝑖𝑜𝑛𝑠)
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Sample Problems
A. Conversion between atoms, molecules, or ions and mass
1.Zinc is an essential mineral that is naturally occurring found in
foods and is also available as dietary supplement. How many
atoms are in 16.5 g of Zinc?
2. Ammonia (NH3) is used for fertilizers and many other things.
How many molecules of ammonia are present in 0.334 g of
ammonia?
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Sample Problems
B. Conversion between mass and moles
1. Ammonium Nitrate (NH4NO3) is a main component of
explosive mixtures used in mining, quarrying, and civil
construction. If an explosive contains 345.0 g of ammonium
nitrate, how many mole of ammonium nitrate are present in
the explosive?
2. Copper is used for the absorption and used of iron in the
formation of haemoglobin. How many grams of Cu are
present in 3.87 mol copper?
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