ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
MODULE 1 (Week 1)
ESSENTIAL CONCEPTS AND FUNDAMENTAL MEASUREMENTS IN CHEMISTRY
MODULE 5 (Week 5-6)
Lesson 1: Matter…………………………………………………………………….2 - 4
Lesson 2: Separation Techniques ……………………....…....…4 - 7
Lesson 3: Methods of Heat Transfer………...…….……….… 7
Lesson 4: Conversion of Units ……………………………...……. 8 – 13
FORCES OF ATTRACTION AND CHEMICAL BONDS
39-45
MODULE 6 (Week 7)
NAMING AND CHEMICAL COMPOUNDS
Activities and Assessment …………………………………………..14 -16
46-53
MODULE 2 (Week 2)
SIGNIFICANT FIGURES AND SCIENTIFIC NOTATION
Lesson 1: Significant Figures …………………………………………17 - 18
Lesson 2: Scientific Notation…………….…………….……………...18 - 21
Activities and Assessment ………………………………………….22
MODULE 3 (Week 3)
ATOMS AND ITS RELATED CONCEPTS
Lesson 1: Structure and History of Discovery……... 23 - 24
Lesson 2: Atomic Models …………………………………………………. 24 – 24
Lesson 3: Atomic Mass and Atomic Number ……... 26 – 27
REMINDER:
Activities and Assessment …………………………………………27 – 29
Read and examine carefully the rubrics for your project (final output) on pages
86 - 88 so that you can prepare or do it in advance. Furthermore, kindly review
modules 1 – 7 for your 3rd quarterly examination. Keep posted to get updates
and schedule(s) for the said exam and project submission.
MODULE 4 (Week 4)
ELEMENTS IN THE PERIODIC TABLE
Lesson 1: Electron Configuration …………………………….30 – 31
Lesson 2: Rules in Orbital Diagram …………...……….… 31
Lesson 3: Lewis Dot Structure …………...……………….… 32 – 35
Activities and Assessment …………………………………...…36 – 38
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ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
THIRD QUARTER
MODULE
Essential Concepts and Fundamental
Measurements in Chemistry
At the end of this module, you are expected to:
INCLUSIVE DATE:
•
•
•
•
Differentiate the different types of solution.
Determine the various separation techniques of a mixture,
Explain the processes that states of matter undergo, and
Perform measurement conversion (larger to smaller or vice-versa,
and English to Metric or vice-versa).
Overview
According to the American Chemical Society (ACS), chemistry is the
study of matter, defined as anything that has mass and takes up space, and
the changes that matter can undergo when it is subject to different
environments and conditions (Lim, 2020). Matter does not include photons
(light), heat, sound, thought, and microwave radiation.
KEY QUESTIONS:
• How does energy become responsible for the changing of one
matter state into another?
• Why is unit conversion important?
COLLOIDS
SUSPENSION
Fig 2. Matter Flowchart
Antoine-Laurent de Lavoisier is the Father of Modern Chemistry.
Lavoisier wrote the book Elements of Chemistry (1787).
• He compiles the first complete (at that time) list of elements,
• Discovered and named oxygen and hydrogen
• Helped develop the metric system
• Helped revise and standardize chemical nomenclature, and
• Discovered that matter retains its mass even when it changes forms.
Fig 1. Common Laboratory Apparatus and Equipment
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ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Table 1: Main Branches of Chemistry
BRANCHES
ANALYTICAL
CHEMISTRY
BIOCHEMISTRY
INORGANIC
CHEMISTRY
DESCRIPTION
It involves the analysis of chemicals, and
includes qualitative methods like looking at
color changes, as well as quantitative methods
like examining the exact wavelength(s) of light
that a chemical absorbed to result in that color
change. These methods enable scientists to
characterize many different properties of
chemicals, and can benefit society in a number
of ways.
• For example, analytical chemistry
helps food companies make tastier
frozen dinners by detecting how
chemicals in food change when they
are frozen over time.
• Analytical chemistry is also used to
monitor the health of the environment
by measuring chemicals in water or
soil.
Biochemistry is the branch of science that
explores the chemical processes within and
related to living organisms.
• An example of biochemistry is the
study of cell metabolism.
Inorganic chemistry deals with synthesis and
behavior of inorganic and organometallic
compounds. It refers to materials not
containing carbon-hydrogen bonds, including
metals, salts, and minerals.
• Inorganic chemistry is used to create a
variety of products, including paints,
fertilizers and sunscreens.
ORGANIC CHEMISTRY
PHYSICAL CHEMISTRY
Organic chemistry deals with chemical
compounds that contain carbon, an element
considered essential to life. Organic chemists
study the composition, structure, properties
and reactions of such compounds, which along
with carbon, contain other non-carbon
elements such as hydrogen, sulfur and silicon.
• Most products you use involve organic
chemistry. Your computer, furniture,
home, vehicle, food, and body contain
organic compounds. Every living thing
you encounter is organic. Inorganic
items, such as rocks, air, metals, and
water, often contain organic matter,
too.
Physical chemistry uses concepts from physics
to understand how chemistry works.
• For example, figuring out how atoms
move and interact with each other, or
why some liquids, including water, turn
into vapor at high temperatures.
•
Unlocking of Terms
• PURE SUBSTANCES are further broken down into elements and compounds.
• MIXTURES are physically combined structures that can be separated into their
original components. A chemical substance is composed of one type of atom or
molecule. A mixture is composed of different types of atoms or molecules that are
not chemically bonded.
o A heterogeneous mixture is a mixture of two or more chemical substances where
the various components can be visually distinguished. For instance, mixtures of
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sand and water, mixtures of sand and iron filings, a conglomerate rock, water and oil, a
salad, trail mix, mixtures of gold powder and silver powder.
Factors Affecting Solubility
1. Temperature
• Basically, solubility increases with temperature. It is the case for most of
the solvents. The situation is though different for gases. With increase of
the temperature they became less soluble in each other and in water, but
more soluble in organic solvents.
o A homogeneous mixture is a type of mixture in which the composition is uniform
and every part of the solution has the same properties.
• SOLUBILITY - is a property referring to the ability for a given substance, the solute,
to dissolve in a solvent. It is measured in terms of the maximum amount of solute
dissolved in a solvent at equilibrium. The resulting solution is called a saturated
solution. The solute is the substance that is being dissolved, while the solvent is
the dissolving medium. Solutions can be formed with many different types and
forms of solutes and solvents. There are three (3) conditions of a solution, namely:
1. A saturated solution contains the maximum amount of solute that will
dissolve at that temperature. Any further addition of solute will result in
undissolved solid on the bottom of the container.
2. An unsaturated solution contains less than the maximum amount of solute
that can be dissolved at that temperature.
3. A supersaturated solution contains more than the maximum amount of
solute that can be dissolved at that temperature. It is unstable and the
solute will usually begin to crystallize, especially if disturbed.
2. Polarity (Nature of Solvent and Solute)
• In most cases solutes dissolve in solvents that have a similar polarity.
Chemists use a popular aphorism to describe this feature of solutes and
solvents: "Like dissolves like". Non-polar solutes do not dissolve in polar
solvents and the other way round.
3. Pressure
• Solid and liquid solutes: For majority of solid and liquid solutes, pressure
does not affect solubility.
• Gas solutes: As for gasses the Henry's law states that solubility of gas is
directly proportional to the pressure of this gas. This is mathematically
presented as: p = kc, where k is a temperature dependent constant for a
gas. A good proof of Henry's law can be observed when opening a bottle
of carbonated drink. When we decrease the pressure in a bottle, the gas
that was dissolved in the drink bubbles out of it.
Table 2: Examples of Solutions
SOLUTIONS
Gas dissolved in gas: dry air
Gas dissolved in liquid:
carbonated water
Liquid dissolved in gas:
moist air
Liquid dissolved in liquid:
vinegar
Solid dissolved in liquid:
sweet tea
SOLUTE
oxygen
SOLVENT
nitrogen
carbon dioxide
water
water
air
acetic acid
water
sugar
tea
4. Molecular or Particle Size
• The larger the molecules of the solute are, the larger is their molecular
weight and their size. It is more difficult it is for solvent molecules to
surround bigger molecules. If all of the abovementioned factors are
excluded, a general rule can be found that larger particles are generally
less soluble. If the pressure, and temperature are the same than out of
two solutes of the same polarity, the one with smaller particles is usually
more soluble.
5. Stirring
• Stirring does not have an effect on solubility of a substance, but everyone
knows that if he puts sugar in his tea and does not stir, it will not dissolve.
Actually, if we left the tea to stand for a long enough time, the sugar would
dissolve. Stirring only increases the speed of the process - it increases
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ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
move of the solvent what exposes solute to fresh portions of it, thus
enabling solubility. As molecules in liquid substances are in constant move,
the process would take place anyway, but it would take more time.
Separation Techniques
1. Handpicking - This method involves simply picking out all the unwanted
substances by hand and separating them from useful ones. For example – if
you separate black grapes from green ones from a mixture of the two.
2. Threshing - This method is mostly done during the harvesting of crops.
Normally, the stalks of the wheat are dried once it is harvested. The grain is
then separated from the stalks and grounded into the floor by beating the dry
stalks to shake off the dried grains.
7. Separating Funnel - Separating funnel is used mainly to segregate two
immiscible liquids. The mechanism involves taking advantage of the unequal
density of the particles in the mixture. Oil and water can be easily separated
using this technique.
3. Winnowing - When the grains are collected from the process of threshing, it
needs to be cleared out of husks and chaffs before it is turned into flour.
Normally the separation of the mixture is carried out with the help of wind or
blowing air. The husk and chaff are blown away by the strong wind when the
farmers drop the mixture from a certain height to the ground. The heavier
grains are collected at one place.
8. Magnetic SeparationFig- When
one substance
3. Sedimentation
Process in the mixture has some magnetic
properties then this method is quite useful. Strong magnets are commonly
used to separate magnetic elements.
9. Simple Distillation - is a method for separating the solvent from a solution. For
example, water can be separated from salt solution by simple distillation. This
method works because water has a much lower boiling point than salt. When
the solution is heated, the water evaporates. It is then cooled and condensed
into a separate container. The salt does not evaporate and so it stays behind.
4. Sieving or Sifting - It is done to separate mixtures that contain substances
mostly of different sizes. The mixture is passed through the pores of the sieve.
All the smaller substances pass through easily while the bigger components of
the mixture are retained.
10. Fractional Distillation - Separating 1 liquid from a mixture of different liquids
that have different boiling points. It is a method for separating a liquid from a
mixture of two or more liquids. For example, liquid ethanol can be separated
from a mixture of ethanol and water by fractional distillation. This method
works because the liquids in the mixture have different boiling points. When
the mixture is heated, one liquid evaporates before the other. One way to
check the purity of the separated liquids is to measure their boiling points. For
example, pure ethanol boils at 78°C and pure water boils at 100°C.
5. Evaporation or Heating - is a technique that is used in separating a mixture
usually a solution of a solvent and a soluble solid. In this method, the solution
is heated until the organic solvent evaporates where it turns into a gas and
mostly leaves behind the solid residue.
6. Filtration or Sedimentation - The most common method of separating a liquid
from an insoluble solid is the filtration. Take, for example, the mixture of sand
and water. Filtration is used here to remove solid particles from the liquid.
Various filtering agents are normally used like filtering paper or other
materials. The sand or mud that stays behind in the filter paper (it becomes
the residue) the water passes through the filter paper (it becomes the filtrate)
11. Chromatography - Paper chromatography is a method for separating dissolved
substances from one another. It is often used when the dissolved substances
are colored, such as inks, food colorings and plant dyes. It works because some
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ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
of the colored substances dissolve in the solvent used better than others, so
they travel further up the paper.
States of Matter
DIAMOND
JUICE
CLOUDS
NEON
Fig 3. A chromatogram, the results of a chromatography experiment
SUBLIMATION
DEPOSITION
A pure substance will only produce one spot on the chromatogram during
paper chromatography. Two substances will be the same if they produce the same
color of spot, and their spots travel the same distance up the paper. In the example
below, red, blue and yellow are three pure substances. The sample on the left is a
mixture of all three.
Fig 4. An illustration of the four common states of matter and
their respective transition processes.
Matter has two properties: physical and chemical properties. Physical
properties can be observed or measured without changing the composition of
matter. Physical properties are used to observe and describe matter. Physical
properties of materials and systems are often described as intensive and extensive
properties. Examples of intensive properties include temperature, refractive
index, density, and hardness of an object. When a diamond is cut, the pieces
maintain their intrinsic hardness (until their size reaches a few atoms thick). In
contrast, an extensive property is additive for independent, non-interacting
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subsystems. The property is proportional to the amount of material in the system.
In other words, intensive properties: A physical property that will be the same
regardless of the amount of matter while extensive properties: A physical property
that will change if the amount of matter changes.
the thermal environment of a building is influenced by heat fluxes through the
ground (conduction), and the building envelope (mostly convection and radiation).
Table 3: Methods of Heat Transfer
Change in which the matter's physical appearance is altered, but
composition remains unchanged. Meanwhile, chemical change results in one or
more substances of entirely different composition from the original substances
like the corrosion of metals. Corrosion is the unwanted oxidation of metals
resulting in metal oxides (2Mg+O2→2MgO(1)).
TYPES/METHODS
CONDUCTION
CONVECTION
RADIATION
DESCRIPTION
is heat flux through solid materials. Heat Flux
Sensors can measure conductive heat flux (see
picture on the left).
• Touching a hot cup of coffee
is heat flux through liquids and gases. Heat Flux
Sensors can measure convective heat flux (see
picture on the left). Examples of convective heat
flux are:
• Feeling much colder when it is windy.
• Feeling much colder in water of 25°C than
in air of 25°C.
Radiation is heat flux through electromagnetic
waves. Heat Flux Sensors can measure radiative
heat flux (see picture on the left). Examples of
radiative heat flux are:
• Feeling hot when standing close to fire.
• Measurement of solar power.
In a nutshell, Heat energy can be transferred from one place to another
by three main processes. In CONVECTION, heat energy is carried by the movement
of particles of matter. In CONDUCTION, heat is transferred by particles vibrating.
In RADIATION, heat is carried directly by electromagnetic waves. Heat flows from
hot to cold objects. When a hot and a cold body are in thermal contact, they
exchange heat energy until they reach thermal equilibrium, with the hot body
cooling down and the cold body warming up.
Fig 5. Organizational breakdown of chemical and physical properties of matter
@https://chem.libretexts.org/
Methods of Heat Transfer
Heat is transferred via solid material (conduction), liquids and gases
(convection), and electromagnetic waves (radiation). Heat is usually transferred in
a combination of these three types and seldomly occurs on its own. For example,
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Measurement and Measurements Conversion
ENGLISH or
Customary
Measurement
In science, a measurement is a collection of quantitative or numerical data that
describes a property of an object or event. A measurement is made by comparing
a quantity with a standard unit. The modern International System of Units (SI)
bases all types of physical measurements on seven base units:
1. the kilogram (kg), for mass
2. the second (s), for time
3. the Kelvin (K), for temperature
4. the ampere (A), for electric current
5. the mole (mol), for the amount of a substance
6. the candela (cd), for luminous intensity
7. the meter (m), for distance
The British or Imperial system of measurements was
common before SI units were adopted internationally.
Although Britain has largely adopted the SI system, the
United States and some Caribbean countries still use
the English system for non-scientific purposes. This
system is based on the foot-pound-second units, for
units of length, mass, and time.
Measurement can be obtained by the following methods: the length of a
piece of string can be measured by comparing the string against a meter stick, the
volume of a drop of water may be measured using a graduated cylinder, the mass
of a sample may be measured using a scale or balance, and the temperature of a
fire may be measured using a thermocouple.
Table 4: Measurement Systems
SYSTEMS
DESCRIPTIONS
SI comes from the French name Système International
SI
d'Unités. It is the most commonly used metric system.
SI is a specific metric system, which is a decimal system
of measurement. Examples of two common forms of
the metric system are the MKS system (meter,
kilogram, second as base units) and CGS system
METRIC
(centimeter, gram, and second as base units). There
are many units in SI and other forms of the metric
system that are built upon combinations of base units.
These are called derived units.
There are 20 accepted prefixes. A prefix may be used to identify multiples of
the original unit or fractions of the original unit. For example, kilo- denotes a multiple
of a thousand, so there are one thousand meters in a kilometer. Milli- denotes a
thousandth; therefore, there are one thousand millimeters in a meter.
Meanwhile, a conversion factor is a number used to change one set of
units to another, by multiplying or dividing. When a conversion is necessary, the
appropriate conversion factor to an equal value must be used. For example, to
convert inches to feet, the appropriate conversion value is 12 inches equal 1 foot.
Below are the lists of conversion factors commonly used in Chemistry and Physics.
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ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Fig 7.1. Conversion Factors @elearnstation.com
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CONVERSION FACTORS FOR TIME AND TEMPERATURE
▪ 1 decade = 10 years
▪ 1 Score = 20 years
▪ 1 Century = 100 years
▪ 1 Millennium = 1,000 years
▪ 1 Leap Year = 366 days
▪ 1 Year = 365 ¼ days
Fig 7.3. Conversion Factors
Useful Temperature Facts
▪ Celsius and Fahrenheit are the same at -40°.
▪ Water boils at 100°C or 212°F.
▪ Water freezes at 0°C and 32°F.
▪ Absolute zero is 0 K.
Fig 7.2. Conversion Factors
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ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
▪
Table 5: Formulae for Temperature Conversion
SCALE
FORMULA
CELSIUS
▪ This temperature scale was developed by Swedish astronomer Andres Celsius in
▪
Kelvin to Celsius
Kelvin to Fahrenheit
Kelvin to Rankine
Kelvin to Reaumur
1742.
Scientists use the Celsius scale for two main reason: In the Celsius scale the
freezing and boiling points of water are 100 units (or degrees Celsius) apart,
freezing point being 0 degrees Celsius and boiling point being set at 100 degrees
Celsius. ... Hence, the Celsius scale is just easier to use.
Celsius to Kelvin
Celsius to Fahrenheit
Celsius to Rankine
Celsius to Reaumur
° C = K – 273.15
° F = K × 1.8 - 459.67
Ra = K × 1.8
R = (K - 273.15) × 0.8
RANKINE
▪ William John Macquorn Rankine developed it in 1857.
▪ Rankine is commonly used in the aerospace industry in the United States. Rankine
K = ° C + 273.15
° F = 9/5 (° C) + 32 OR
° F = 1.8 (° C) + 32
Ra = °C × 9/5 + 491.67 OR
Ra = °C × 1.8 + 491.67
Re = °C × 0.8
is to Fahrenheit what Kelvin is for Celsius. So when people in the United States
were creating programs and using equations that needed an absolute
temperature, they used Rankine before Celsius became dominate for scientific
calculations.
Rankine to Celsius
Rankine to Fahrenheit
Rankine to Kelvin
Rankine to Reaumur
FAHRENHEIT
▪ It was developed by Daniel Gabriel Fahrenheit in 1724
▪ Fahrenheit is superior for measuring temperature precisely. It's also better
because humans tend to care more about air temperature rather than water
temperature.
Fahrenheit to Celsius
Fahrenheit to Kelvin
Fahrenheit to Rankine
Fahrenheit to Reaumur
Kelvin used this as a basis for an absolute temperature scale. He defined
"absolute" as the temperature at which molecules would stop moving, or "infinite
cold."
°C = (Ra - 32 - 459.67) / 1.8
°F = Ra - 459.67
K = Ra / 1.8
Re = (Ra - 32 - 459.67) / 2.25
REAUMUR
• Réaumur temperature scale, scale established in 1730 by the French naturalist
° C = 5/9 (° F - 32)
K = 5/9 (° F - 32) + 273.15
Ra = F + 459.67
Re = (F - 32) / 2.25
•
•
KELVIN
▪ It was developed by William Thomson aka Lord Kelvin in 1848.
▪ The Celsius and Fahrenheit scales were both built around water, either the
René-Antoine Ferchault de Réaumur.
The Rankine scale is used in engineering systems where heat computations are
done using degrees Fahrenheit.
In the Netherlands, the Reaumur thermometer is used when cooking.
Reaumur to Celsius
Reaumur to Fahrenheit
Reaumur to Kelvin
Reaumur to Rankine
freezing point, the boiling point or some combination of water and a chemical.
The Kelvin temperature scale is used by scientists because they wanted a
temperature scale where zero reflects the complete absence of thermal energy.
11
°C = Re × 1.25
°F = Re × 2.25 + 32
K = Re × 1.25 + 273.15
Ra = Re × 2.25 + 32 + 459.67
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
OR
SAMPLE UNITS CONVERSION
780,000cm
A. LARGER UNITS TO SMALLER UNITS
To convert from a larger unit to a smaller one, multiply.
EXAMPLES:
1. Convert 3.2 meters to cm
1 meter = 100 cm
3.2m = (3.2 x 100) = 320 cm
•
x
1m
1km
100 cm
1000m
x
= 7.8 km
Cancel the same units, then perform the necessary operation(s)
C. METRIC TO ENGLISH (US or IMPERIAL)
EXAMPLES:
1. Convert 115cm to inches
1 inch = 2.54 cm
115 cm x 1 / 2.54cm = 45.28 inches
2. Convert 7.8 km to mm
1 km = 1,000 m and 1m = 100 cm
1 km = (7.8 x 1000) = 7800 m
1 m = (7800 x 100) = 780,000 cm or 7.8 x 10^5 cm
2. Convert 115km to mile
1 mile = 1.6 km or 1.60934km
115 km x 1mile / 1.6km = 71.88 mi
115 km x 1mile / 1.60934km = 71.46 mi
B. SMALLER UNITS TO LARGER UNITS
To convert from a smaller unit to a larger one, divide.
EXAMPLES:
1. Convert 320cm to meter
D. ENGLISH TO METRIC
EXAMPLES:
1. Convert 100 yard to meter
1 yd = 3ft, 1ft = 12in, 1in = 2.54cm, 1m = 100cm
100 yd x 3ft / 1yd = 300 ft
300ft x 12in / 1ft = 3600 in
3600 in x 2.54 cm / 1in = 9144 cm
9144 cm x 1m / 100 cm = 91.44 m
Since in every 1m there are 100cm, so the ratio will be 1/100
320 cm x 1m / 100cm – cancel the same units
(320 x 1) /100 – get the product of 320 and 1 then divide by 100
3.2 – put the remaining unit which is m
= 3.2 m
2. Convert 780,000 cm to km
Since in every 1km there are 1,000 meters and in every 1 meter there are 100 centimeters, so the
ratios will be 1/1000 and 1/100
OR
780,000cm x 1m/100 cm - cancel the same units, then perform the necessary operation(s)
780,000 x 1 / 100 = 7,800 m
7,800 m x 1km/1,000 m – cancel the same units, then perform the necessary operation(s)
7,800 x 1 / 1,000 = 7.8 km
100yd
3ft
12in
xx x
1yd
•
2.54cm
x
1ft
1m
1in
x
100cm
Cancel the same units, then perform the necessary operation(s)
2. Convert 20ft to meter
1foot = 12in = 30.5cm (30.48cm) = 0.305m (or 0.3048cm)
20ft x 30.48 cm / 1ft x 1meter / 100 cm = 6.1 m OR
OR
780,000 cm x 1m/100 cm x 1km / 1000m – cancel the same units
780,000 x 1 / 100 x 1/1000 – follow the MDAS rule (solve from left to right)
780,000/100 x 1/1000 = 7800x1/1000 = 7800 / 1000 – put the remaining unit
= 7.8 km
20 feet
12
x
30.48 cm
1m
1 foot
100 cm
x
= 6.1 m
= 91.44 m
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
E. TIME
EXAMPLES:
1. Convert 15 minutes to second
1 minute = 60 seconds
4. Convert 315 K to Celsius
FORMULA: ° C = K – 273.15
° C = 315 – 273.15
= 41.85 ° C
60 s
15 min
x
5. Convert 315 K to Fahrenheit
FORMULA: ° F = K × 1.8 - 459.67
° F = 315 × 1.8 - 459.67
= 107.33 ° F
= 900 s
1 min
2. Convert 1,000,000 seconds to year
1 year = 52 weeks = 12 months = 365 ¼ days
1 day = 24 hours, 1 hour = 60 minutes, 1 minute = 60 seconds
1 min
1,000,000 s
x
1h
1 day
x
60 s
60 min
x
24 h
= 0.0317 calendar year or 0.0316 leap year
REFERENCES
1y
Lim, A. (2020, July 29). What is Chemistry? Retrieved from
https://www.livescience.com/45986-what-is-chemistry.html
x
365.25 days
Substances
and
Mixtures.
(n.d.)
Retrieved
from
https://courses.lumenlearning.com/introchem/chapter/substances-andmixtures/#:~:text=Matter%20can%20be%20broken%20down,type%20of%
20atom%20or%20molecule.
F. TEMPERATURE
EXAMPLES:
1. Convert 50 °C to Fahrenheit
FORMULA: ° F = 9/5 (° C) + 32 OR ° F = 1.8 (° C) + 32
° F = 1.8 (° C) + 32
= (1.8 x 50) + 32 = 90 + 32 = 122 ° F
2. Convert 50 ° C to Kelvin
FORMULA: K = ° C + 273.15
K = ° C + 273.15 = 50 + 273.15 = 323.15 K
Methods
of
Separating
Mixtures.
(n.d.).
https://byjus.com/chemistry/methods-of-separation/
3. Convert 122 F to Celsius
FORMULA: ° C = 5/9 (° F - 32)
°C=
5
5
(°F – 32) =
9
=
(122 – 32)
9
5
90
9
1
=
450
9
Retrieved
from
Separating
Mixtures.
(n.d.).
Retrieved
https://www.bbc.co.uk/bitesize/guides/zgvc4wx/revision/4
from
Solubility
of
Things.
(n.d.).
Retrieved
https://www.solubilityofthings.com/levels-of-solubility
from
Physical and Chemical Properties of Matter. (2020, August 17). Retrieved
from
https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Modules_an
= 50 ° C
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ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
d_Websites_(Inorganic_Chemistry)/Chemical_Reactions/Properties_of_Ma
tter
Important
RemInderS
Conduction, Convection, and Radiation. (n.d.). Retrieved from
https://www.greenteg.com/heat-flux-sensor/about-heat-flux/3-types-ofheattransfer/#:~:text=The%20three%20types%20of%20heat,seldomly%20occur
s%20on%20its%20own.
•
•
Helmenstine, Anne Marie, Ph.D. (2020, August 26). Measurement Definition
in Science. Retrieved from https://www.thoughtco.com/definition-ofmeasurement-605880
Tear this activity sheet and submit on the scheduled date along with the
other activity (ies) the instructor may have asked the students to do on a
separate paper.
If you are sending something you’ve done online such as MS presentation
(s), pictures, pdfs and alike as an attachment, then you may send them to
my email at germanvertudez1211gmail.com following this format:
(SECTION_LASTNAME_FIRSTNAME_ACTIVITYNAME
e.g.
IC1MA_BINABAN_PRINCESS_SCAVENGERS HUNT), or send a digital copy
from your flash drive together with this activity sheet.
Name: _______________________________________________________
Grade Level & Section: __________________________________________
Date Submitted: (to be filled in by the subject instructor): _______________
RATING:
MODULE 1:
Assessment
PART A: UNIT CONVERSION (45 PTS)
Refer to the unit conversion factor on Module 1 for answering. NO SOLUTION shown means
wrong. (2 points each)
1. 5 decades to months
2. 15 years to hours
3. 5 feet 5 inches in meters
4. 2000 kilometers in meters
5. 6.75 hours to minutes
6. 100 degrees Fahrenheit to Celsius
7. 0.006 degrees Celsius to Fahrenheit
8. 138 degrees Fahrenheit to Kelvin
9. 116 inches to feet
10. 1200 mL to liters
14
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Refer to the unit conversion factor on Module 1 for answering. NO SOLUTION shown means
wrong. (5 points each)
1. Nathaniel can lift 200kg with ease. How much is this in pounds?
2. A can of soda holds 355mL. How many mL would be in 2 cans of soda?
3. One cubic meter of Ms. Andes’ blood contains 5,000,000 WBC. There are about
4,900,000 cubic millimeters of blood in her body. Determine the approximate
number of WBC she has in her body. Express your final answer in correct scientific
notation.
4. A microscope is set so it makes an object appear 4x10^2 times larger than its actual
size. A virus has a diameter of 2X10^-7 meter. How large will the diameter of the
virus appear when it is viewed under the microscope? Express your final answer in
correct scientific notation.
5. A box contains 5X10^3 thumb tucks. The mass of each thumb tuck in the box is
8X10^-4 kilograms. What is the combined mass of the thumb tucks in the box?
Express your final answer in correct scientific notation.
4.
A gardener carefully places her outdoor thermometer in a shady location out of
direct sunlight, so that it doesn't give incorrectly high readings. What method of heat
transfer is she trying to avoid?
a. Induction
b. Radiation.
c. Conduction
d. Convection
5.
Which one of the following do all methods of heat transfer require?
a. The movement of particles
b. A difference in thermal energy.
c. A liquid or gaseous state
d. Direct physical contact
6.
Which of the following statements about radiation heat transfer is true?
a. A radiant heat source transfers heat by energizing the molecules of air around it
b. Radiation heat transfer does not involve particles.
c. Only glowing objects can be a radiant heat source
d. Radiant heat transfer explains why a spoon in a cup of hot tea soon feels warm
7.
Liquids
a. have a defined shape and volume
b. have a defined shape but undefined volume
c. have an undefined shape but defined volume.
d. have both undefined shape and volume
8.
Why do solids have a fixed shape?
a. The particles are fixed in place.
b. The particles are closely packed
c. The particles vibrate
d. None of these
9.
What is the name of the change when a liquid becomes a solid?
a. Melting
b. Solidification
c. Freezing
d. Both B and C.
PART C: MULTIPLE CHOICE (15 PTS)
Circle the letter that corresponds to your answers. (1 point each)
1.
Which one is a form of heat transfer?
a. Conduction
b. Refraction
c. Absorption
d. Reflection
2.
Which of the following statements about convection is true?
a. Convection always involves the circulation of a liquid or gas.
b. All types of currents are convection currents
c. Convection occurs between solids only at high temperatures
d. Convection can only occur during the process of boiling
3.
As a thunderhead builds, warm air at the Earth's surface rises and cold air high aloft
sinks downward. What heat transfer process is occurring here?
a. Conduction
b. Radiation
c. Convection.
d. Induction
15
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
10. What is sublimation?
a. When a solid turn into a gas.
b. When a gas turns into a solid
c. When a gas turns into a liquid
d. When gas turns into solid
d.
11. If you want to separate iron fillings from sand, you would use a _____.
a. Funnel
b. Filter
c. magnet.
d. colander or sieve
12. The process used to separate heterogeneous mixtures of solids and liquids is called
______.
a. filtration.
b. crystallization
c. distillation
d. Chromatography
13. A technique that separates a mixture based on the individual substance's tendency
to travel across a surface is called _____.
a. filtration
b. crystallization
c. distillation
d. chromatography.
14. During filtration, an insoluble solid collect on the filter paper. What is this called?
a. Solute
b. Solvent
c. Filtrate
d. Residue.
15. In the lab, a scientist accidentally dropped a ball bearing into a beaker of hot water
and wanted to get it out as soon as possible as the hot water will be used in an
experiment. What can he do?
a. Use paper chromatography
b. Use distillation to isolate the hot water
c. Use a magnet to attract the ball bearing and get it out of the hot water.
16
Evaporate the hot water to get the ball bearing
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
1
THIRD QUARTER
MODULE
INCLUSIVE DATE:
INCLUSIVE DATE:
Overview
When making a measurement, there is a limit to the accuracy of the
Significant Figure
and Scientific Notations
At the end of this module, you are expected to:
•
•
•
Write numbers in scientific notation;
Evaluate expressions in scientific notation; and
Perform the basic operations for significant figures and scientific
notation.
reported value. Both the reporter and the reader must follow the conventions of
significant figures when handling measured quantities. There are simple rules
which are used to tell how many significant figures are contained in a value.
Rules on Determining the Number of Significant Figures
1. All non-zero digits are significant
EXAMPLES: 123 (3 SFs), 6.123 (4 SFs)
KEY QUESTIONS:
• How does energy become responsible for the changing of one
matter state into another?
• Why is unit conversion important?
2. Zeroes between non-zero digits are significant
EXAMPLES: 106 (3 SFs), 140.01 (5 SFs)
3. All zeroes which come after non-zeroes and after the decimal place are
significant OR the trailing zeroes after / at the right side of the decimal point
are significant.
EXAMPLES: 3.00 (3 SFs), 1.0110 (5 SFs)
4. For number less than 1, all zeroes which come after the decimal point but are
before non-zeroes are not significant OR the leading zeroes (place holders) are
not significant
EXAMPLES: 0.003 (1 SF), 0.070 (2 SFs)
5. Trailing zeroes in a whole number with no decimal point shown are not
significant. (NOTE: When zeroes come before the decimal point and are after non-zeroes, it
is impossible to tell how many significant figures are present. Therefore, the number of SFs is
unknown)
EXAMPLES: 540 (2 SFs or unknown), 1000 (1 SF or unknown)
6. Trailing zeroes in a whole number with a decimal point shown are significant.
(NOTE: It is improper to report numbers in this fashion. Some ascribe to the system that a
decimal point implies all zeroes following the non-zeroes are significant)
EXAMPLES: 540. (3 SFs), 1000. (4 SFs)
7. For numbers in scientific notation, cases 1 through 3 are used to determine
the number of significant figures.
Fig 1. Parts of Scientific Notation
17
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
CASE 1: Multiplying or Dividing Significant Figures
RULES:
1. Multiply or divide the values as normal.
2. Determine the number of significant figures in the multiplicands.
3. Round the result off to the number of significant figures as the value with
the fewest OR least number of significant figures.
EXAMPLES:
1.36 × 8.9
= 12.104 = 12 (2 SFs)
79.502 ÷ 99.15
= 0.8018356026 = 0.8018 (4 SFs)
8.136 × 10^8 ÷ 6 × 10^-4 = 1356000000000 = 1 × 10^12 (1SF)
19.95 × 8.0
= 159.6 or 1.596 X10^2 = 1.6 × 10^2 (2SFs)
279 x 83
= 23.157 = 23,000 (2 SFs)
EXAMPLES:
The area of a triangle:
Area = 1/2 Base × Height
The SFs are solely determined by the base and height
8. Exact numbers have an infinite number of significant figures
EXAMPLE:
1 meter = 1.00 meter = 1.000000000 meter etc.
Scientific Notation
Scientific notation (also called exponential notation) is the way that scientists
easily handle very large numbers or very small numbers. For example, instead of
writing 0.0000000056, we write 5.6 x 10^-9. Scientific notation has three parts to it:
the coefficient, the base, and the exponent.
CASE 2: Adding or Subtracting Significant Figures
RULES:
1. Add or subtract the values as normal.
2. Determine the number of decimal places in the values.
3. Round the result off to the number of decimal places as the value with
the fewest or with the least number of decimal places.
EXAMPLES:
7.03199 + 6.954
= 13.98599 = 13.986 (5 SFs)
7.52 × 10^3+ 9.80 × 10^2 = 8500 = 8.50 × 10^3 (3 SFs)
1.00 - 8.6754
= -7.6754 = -7.68 (3 SFs)
27 - 0.440
= 26.56 = 27 (2 SFs)
300 – 47.465
= 252.535 = 300 (1 SF)
300. – 47.465
= 252.535 = 253 (3 SFs)
Fig 2. Parts of Scientific Notation
NOTES:
• The coefficient must be greater than 1 and
less than 10 and contain all the significant
(non-zero) digits in the number or
mantissa.
• The base is always 10.
• The exponent is the number of places the
decimal was moved to obtain the
coefficient.
• Sign is indicated if necessary.
RULES:
1. All numbers must be ≥ 1 and ≤ (or 1 – 9, or the first significant digit in a number)
2. All exponents on the X10^n must be whole numbers.
3. The exponent on the X10^n must be equal to the number of places you move the
decimal point to satisfy Rule #1.
CASE 3: Formula Containing Constants
RULES:
1. When using a formula which contains irrational physical constants such
as pi, the value contains as many significant figures as you enter in the
formula.
2. When using a formula which contains integers, the integers are assumed
to have infinite significance, and do not limit the number of significant
figures in the result
CONVERTING STANDARD FORM TO SCIENTIFIC NOTATION
• If the number is big or greater than 1 or when your start moving the decimal
point from RIGHT TO LEFT, the exponent must be positive (+)
18
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
5,000,000,000,000
- take the first significant digit which is 5
- 5 X 10^12 since it took 12 moves or there are 12
numbers after the first significant digit which is 5.
• If the number is small or less than one or when you start moving the decimal
point from LEFT TO RIGHT, the exponent must be negative (-).
0.000000000005
- take the first significant digit which is 5
- 5 X 10^-12 since it took 12 moves or there are 12
numbers before the first significant digit which is 5.
Exponent PLUS 1 (-4 + 1 = 3), that’s the number of zero digits we need to
place before the given significant digits or after the decimal point, giving us
the answer of 0.0003456
4. Only have as many numbers as you have significant figures. Take all the
mantissa (magnitude) or the significant digits to the right of the decimal
point.
1,213,400,000
= 1.2134 X 10^9 (CORRECT)
= 1.21 X10^9 (WRONG)
≈ 1.21 X10^9 (ACCEPTABLE, if directed)
CONVERTING SCIENTIFIC NOTATION TO STANDARD FORM (or REAL NUMBERS)
NOTE: The equals sign or equality sign (=) is a mathematical symbol used to indicate equality in
METHOD 1:
Multiply the decimal number by 10 raised to the power indicated.
EXAMPLES:
1. 3.456 x 10^4 = 3.456 x 10,000 = 34560
2. 3.456 x 10^-4 = 3.456 x (1/10000) = 0.0003456 OR
3.456 x 10^-4 = 3.456 x .0001 = 0.0003456
some well-defined sense. On the other hand, approximately equal to sign (≈) is used to show
inaccuracy or just an estimate.
Operations on Scientific Notation
1. Multiplying and Dividing Scientific Notation
METHOD 2:
Multiply the decimal number by 10 raised to the power indicated.
EXAMPLES:
3.456 x 10^4 = 34560
1. Exponent MINUS the number of mantissa or significant digits after the
decimal point, 4 – 3 = 1, it means that you need to add 1 zero after the last
significant figure which in this example is 6 OR
RULES FOR MULTIPLICATION:
a. Multiply the coefficients
b. Add the exponents
c. Express or round off the final answer in scientific notation based on the given
that has the LEAST number of significant figures
EXAMPLE:
(5.60×10^12) (7.102×10^4)
5.6×7.102 X 10 ^ (12+4)
39.8 X 10 ^ 16 – move the decimal point once from RIGHT TO LEFT (plus)
3.98 X 10 ^ 16+1
3.98 X 10 ^ 17
1. Move the decimal point three times from LEFT TO RIGHT (minus)
3.456 10^4-3 = 3456 x 10^1
2. Move it one more time from LEFT to RIGHT (minus); since you take one move
you need now to add a zero digit after the last digit in the number
3456 x 10^1-1
34560 x 10^0 = 34560 x 1 = 34560 (any number with a zero exponent is equal to 1)
RULES FOR DIVISION:
a. Divide the coefficients
b. Subtract the exponent in the denominator from the exponent of the
numerator
3.456 x 10^-4 = 0.0003456
1. Always add 1 to the negative exponent
19
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
c. Express or round off the final answer in scientific notation based on the given
that has the LEAST number of significant figures
EXAMPLE:
(3.04×10^5) ÷ (9.89×10^2)
5.928 X 10^3 – it’s already in the scientific notation
5.9 X10^3
RULES FOR MIXED OPERATIONS
FOR SIGNIFICANT FIGURE AND SCIENTIFIC NOTATION
3.04
X 10^ (5-2)
9.89
0.307 X 10^3 – move the decimal point once from LEFT TO RIGHT (minus)
3.07 X 10^ (3-1)
3.07 X 10^2
Follow GEMDAS – Grouping (parentheses, braces, brackets), E (exponent), M
(multiplication), D (division), A (addition), and S (subtraction) – solve it from left to right
fashion.
EXAMPLES:
1. 1.123 + 2.1 – 3.12
3.223 – 3.12
3.224 = 0.103 or 0.1 – since 2.1 has one decimal place (or least precise)
2. 1.123 – 2.1 + 3.12
- 0. 977 +3.12
2.143 or 2.1 – since 2.1 has one decimal place (or least precise)
2. Adding and Subtracting Scientific Notation
RULES FOR ADDITION AND SUBTRACTION:
a. Adjust the powers of 10 in the 2 numbers so that they have the same index.
(Tip: It is easier to adjust the smaller index to equal the larger index).
b. Add or subtract the numbers.
c. Express or round off the final answer in scientific notation based on the given
that has the least number of decimal places. Don’t forget to include all the
mantissa or significant digits after the decimal point, and put an overbar for
repeating decimals.
EXAMPLES:
2 × 10^3 + 3.6 × 10^4
0.2 × 10^ (3+1) + 3.6 × 10^4 – make one move from RIGHT TO LEFT (plus) in the first
3. 1.123 X 2.1 – 3.12
2.3583 – 3.12
-0.7617 or -0.76 since 2.1 has 2 SFs
4. 1.123 x 2.1 ÷ 3.12
2.3586 ÷ 3.12
0.7558653846 or 0.76 since 2.1 has 2 SFs
given with smaller index or exponent.
5. 1.123 ÷ 2.1 x 3.12
0.5347619048 x 3.12
1.668457143 or 1.7 since 2.1 has 2SFs
0.2 × 10^4 + 3.6 × 10^4
0.2 + 3.6 X 10^4 – add the numbers
3.8 X 10^4 or 4 X10^4
6. 1.123 ÷ 2.1 x 3.12 – 1.123 + 2.1
0.5347619048 x 3.12 – 1.123 + 2.1
1.668457143 – 1.123 + 2.1
0.545471429 + 2.1
2.645457143 or 2.6 since 2.1 has 2SFs
6.2 X10^3 – 2.72 X10^2
6.2 X10^3 – 0.272 X10^ (2+1) - make one move from RIGHT TO LEFT (plus) in the first
given with smaller index or exponent.
6.2 X10^3 – 0.272 X10^3
6.2 – 0.272 X10^3 – subtract the numbers
20
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
7. 1.123^2 + 2.1 – 3.12
1.261129 + 2.1 – 3.12
3.361129 – 3.12
0.241129 or 0.24 since 2.1 has 2SFs
8.
REFERENCES
Significant Figures. (n.d.). Retrieved from
https://www2.southeastern.edu/Academics/Faculty/wparkinson/help/significa
nt_figures/
1.123 x (2.1) ^2 – 3.12
3.42
1.123 x 4.41 – 3.12
3.42
4.95243 – 3.12
3.42
1.83243
3.42
0.5357982456 or 0.54 since 2.1 has 2SFs
Scientific Notation. (2014, December 14). Retrieved from
https://www.ck12.org/book/ck-12-algebra-i-second-edition/section/8.4/
Vertudez, German. Gibby-Eskwela. (2020, August 10). Significant Figures
(video). YouTube. Retrieved from
https://www.youtube.com/watch?v=7jIMJfa3Dyc
21
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Important
RemInderS
•
•
PART B: SCIENTIFIC NOTATION (20 PTS)
Express the following scientific notation into its standard form, vice-versa or express it to its
correct scientific notation if needed (2 points each)
1. 200000
2. 0.000006
3. 0.63
4. 60
5. 0.0091
6. 2X10^-7
7. 2.16X10^-4
8. 4.89X10^4
9. 33 X10^-3
10. 71X10^3
Tear this activity sheet and submit on the scheduled date along with the
other activity (ies) the instructor may have asked the students to do on a
separate paper.
If you are sending something you’ve done online such as MS presentation
(s), pictures, pdfs and alike as an attachment, then you may send them to
my email at germanvertudez1211gmail.com following this format:
(SECTION_LASTNAME_FIRSTNAME_ACTIVITYNAME
e.g.
IC1MA_BINABAN_PRINCESS_SCAVENGERS HUNT), or send a digital copy
from your flash drive together with this activity sheet.
PART C: COMPUTATION (30 PTS)
Name: _______________________________________________________
Grade Level & Section: __________________________________________
Date Submitted: (to be filled in by the subject instructor): _______________
Compute the following items as required. (2 points each)
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
RATING:
MODULE 2:
Assessment
PART A: SIGNIFICANT FIGURE (10 PTS)
Determine the number of sig figs in each item. (1 point each)
1. 42050
2. 7080
3. 30,050.
4. 0.08060
5. 750.064080
6. 17
7. 101
8. 0.0001305
9. 500
10. 500.
22
4.231 + 3.51
7.2361 + 8.42
400 – 47.465
421 – 3.59
420. + 3.51
8.4 X 5
279 X 83
3.46 X 2.1
464.6895 / 12.145
3.6 X 4.6439 + 5.831
1.65X10^4 + 9.71X10^4
3.71X10^3 + 7.316X10^3
1.65X10^4 + 9.71X10^3
(2.6X10^7) (4.1X10^-3)
2.5X10^4 / 5 X10^2
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
THIRD QUARTER
MODULE
The Greek philosopher Democritus introduced the idea of the atom.
However, the idea was essentially forgotten for more than 2000 years. In 1800,
John Dalton re-introduced the atom. In Greek, the prefix "a" means "not" and the
word "tomos" means cut. Our word atom therefore comes from atomos, a Greek
word meaning uncuttable or indivisible.
Atoms and Its Related
Concepts
At the end of this module, you are expected to:
INCLUSIVE DATE:
•
•
•
Explain the different atomic models;
Determine the atomic mass and atomic number; and
Discuss the different parts of an atom.
Table 1: A TIMELINE ON THE BRIEF HISTORY ON MATTER AND ATOMS
YEAR
KEY QUESTIONS:
450 BC
• Why are atomic models important?
• How did technological advancements pave the way for
modifying the atomic models?
400 BC
380 – 320 BC
1799
Overview
1808
Atoms
are
extremely
1869
small – with diameters about many
billionths of an inch – which make
it impossible to be seen with the
naked eyes. However, with the
invention of the scanning tunneling
microscope (STM) and other
sophisticated instruments, atoms
may now be seen.
1890s
1895
1897
1904
1908 – 1917
The nucleus (center) of the atom contains the protons (positively charged)
and the neutrons (no charge). The outermost regions of the atom are called
electron shells and contain the electrons (negatively charged).
1910 – 1911
23
EVENT
Empedocles asserted that all things are composed of four
primal elements: earth, air, fire, and water.
Democritus proposed that all matter is made up of very
small particles called atom, which cannot be divided into
smaller units.
Aristotle proposed that all matter was continuous and be
further divided infinitely into smaller pieces.
Joseph Proust proposed the Law of Definite Proportions
John Dalton formulate the atomic theory and proposed
the Law of Multiple Proportions.
Dmitry Mendeleev arranged the known elements in a
periodic table based on the atomic mass.
Antoine Becquerel and Marie Curie observed that
radioactivity causes some atoms to break down
spontaneously.
Wilhelm Roentgen discovered X-rays
Joseph John Thomson discovered electrons
J.J. Thomson suggested the Plum Pudding Model that
shows electrons move in concentric orbits around a
nucleus
Robert Millikan found that the charge of an electron is
equal to – 1.6022 X10^-19 coulombs.
Ernest Rutherford observed that atoms are mostly empty
space.
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
1913
1919
1932
Niels Bohr proposed an atomic model that shows
electrons move in concentric orbits around a nucleus.
Henry Gwyn Jeffreys Moseley used X-rays spectra to
study atomic structure.
E. Rutherford discovered protons
James Chadwick discovered neutrons.
2. E. RUTHERFORD’S MODEL (1911)
(Discovery of Proton)
DALTON’S ATOMIC THEORY (1808)
1. Each chemical element is composed of extremely small particles that are
indivisible and cannot be seen by the naked eye, called atoms.
2. All atoms of an element are alike in mass and other properties, but the atoms
of one element differ from all other elements. For example, gold and silver
have different atomic masses and different properties.
3. For each compound, different elements combine in a simple numerical ratio.
4. In a chemical reaction, atoms are neither created nor destroyed. They simply
combine, separate, or rearrange.
Fig 3. Rutherford’s Atomic Model
ATOMIC MODELS
Physicist Ernest Rutherford envisioned the atom as a miniature solar
system, with electrons orbiting around a massive nucleus, and as mostly empty
space, with the nucleus occupying only a very small part of the atom. The neutron
had not been discovered when Rutherford proposed his model, which had a
nucleus consisting only of protons. In 1909, Ernest Rutherford (1871-1937)
performed a series of experiments by shotting a beam of alpha particles (positively
charged) at a thin piece of gold foil. Rutherford observed that the majority of the
alpha particles went through the foil; however, some particles were slightly
deflected, a small number were greatly deflected, and another small number were
thrown back in nearly the direction from which they had come. In other words,
Rutherford's model shows that an atom is mostly empty space, with electrons
orbiting a fixed, positively charged nucleus in set, predictable paths. In other
words, this model states that there were positive particles within the nucleus, but
failed to define what these particles are. Rutherford discovered these particles in
1919, when he conducted an experiment that scattered alpha particles against
1. THE PLUM PUDDING MODEL (1898 – 1903)
(Discovery of Electron)
After J.J. Thompson discovered the electron in the cathode rays, he proposed
the plum pudding model of an atom, which states that the electrons float in
positively-charged material. This model was named after the plum-pudding
dessert. However, it was George Stoney who first gave the term electrons to the
cathode rays. The plum pudding model is defined by electrons surrounded by a
volume of positive charge, like negatively-charged “plums” embedded in a
positively-charged “pudding” (hence the name).
Fig 2. Plum Pudding Model
24
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
nitrogen atoms. When the alpha particles and nitrogen atoms collided, protons
were released.
In 1933, James Chadwick (1891-1974) discovered a new type of radiation that
consisted of neutral particles. It was discovered that these neutral atoms come
from the nucleus of the atom. This last discovery completed the atomic model.
•
lowest energy level. The orbits n=1, 2, 3, 4… are assigned as K, L, M, N…. shells and
when an electron attain the lowest energy level, it is said to be in the ground state.
The electrons in an atom move from a lower energy level to a higher energy level by
gaining the required energy and an electron moves from a higher energy level to lower
energy level by losing energy.
4. SOMMERFELD’S MODEL (1916)
3. BOHR’S ATOMIC MODEL (1913)
In 1913 Bohr proposed his quantized shell model of the atom to explain how
electrons can have stable orbits around the nucleus. Bohr modified the Rutherford
model by requiring that the electrons move in orbits of fixed size and energy. The
energy of an electron depends on the size of the orbit and is lower for smaller
orbits. Radiation can occur only when the electron jumps from one orbit to
another. The atom will be completely stable in the state with the smallest orbit,
since there is no orbit of lower energy into which the electron can jump.
Fig 5. Sommerfeld’s Model of an Atom
The German Physicist Arnold Sommerfeld postulated that electrons can
follow elliptical, rather than a purely circular, paths and therefore introducing a
new (azimuthal) quantum number to account for angular momentum. This model
explains the fine spectrum of Hydrogen atom. The important postulates of
Sommerfeld atomic model are:
Fig 4. Bohr’s Model of an Atom
•
•
•
•
•
•
Bohr’s model consists of a small nucleus (positively charged) surrounded by negative
electrons moving around the nucleus in orbits. Bohr found that an electron located
away from the nucleus has more energy, and electrons close to the nucleus have less
energy.
In an atom, electrons (negatively charged) revolve around the positively charged
nucleus in a definite circular path called orbits or shells.
Each orbit or shell has a fixed energy and these circular orbits are known as orbital
shells.
The energy levels are represented by an integer (n=1, 2, 3…) known as the quantum
number. This range of quantum number starts from nucleus side with n=1 having the
•
•
25
The orbits may be both circular and elliptical.
When path is elliptical, then there are two axis – major axis & minor axis. When
length of major & minor axis becomes equal then orbit is circular.
The angular momentum of electron moving in an elliptical orbit is (kh/2π). Where
k is an integer except zero. Value of k = 1,2,3,4. (n/k)= length of major axis / length
of minor axis. With increase in value of k, ellipticity of the orbit decreases. When
n= k, then orbit is circular.
Sommerfeld suggested that orbits are made up of sub energy levels. These are s,
p, d, f. These sub shells possess slightly different energies
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
•
•
atomic number. In contrast, the number of neutrons for a given element can vary.
Forms of the same atom that differ only in their number of neutrons are called
isotopes such as carbon, potassium, and uranium - have multiple naturally
occurring isotopes.
When an electron jumps from one orbit to another orbit, the difference of energy
(ΔE) depends upon sub energy levels.
It explains the splitting of individual spectral lines of hydrogen & thus fine
spectrum. It could not predict the exact number of lines which are actually present
in the fine spectrum.
ATOMIC NUMBER (Z)
Z = P = e (when an atom of an element is neutral)
ATOMIC MASS AND ATOMIC NUMBER
Today, we know that the atomic number gives the number of protons (positive
charges) in the nucleus. This was the discovery made by Henry Gwyn-Jefferies
Moseley. He found that certain lines in the X-ray spectrum of each element moved
the same amount each time you increased the atomic number by one. Atomic
number (Z), also called proton number, refers to the number of proton number in
the nucleus of each atom of an element. In a neutral atom, proton is equal to
electron number
ATOMIC MASS (A)
A=P+N
Table 2: ATOMIC NUMBER AND ATOMIC MASS
SYMBOL
ELEMENT NAME
A
I
Iodine
127
Cr+2
Chromium > Chromous ion
51
O-2
Oxygen > oxide
16
Z
53
24
8
P
53
24
8
E
53
22
10
N
74
28
8
Atoms have different properties based on the arrangement and number
of their basic particles. Before we delve into how the 118 elements in the Periodic
Table were formed, let’s familiarize first the following radioactive particles
involved in nuclear reactions: a) fusion – two nuclei combine and b) fission – a
nucleus split into two smaller nuclei. These two reactions or processes involving
the emission of energetic particles of an atom is called radioactivity. Below are the
examples of radioactive particles present in nuclear reactions:
MOST COMMON TYPES OF NUCLEAR REACTIONS
1. ALPHA DECAY - high speed particle consisting of 2 protons and 2 neutrons. 2
protons and 2 neutrons lost. Atomic number down by 2, atomic mass down by
4.
2. BETA DECAY – high speed electron. 1 neutron turns into a proton. Atomic
number up by 1
Fig 2. Atomic Number (Z) and Atomic Mass (A) of Sodium
Atoms of each element contain a characteristic number of protons. In fact,
the number of protons determines what atom we are looking at (e.g., all atoms
with six protons are carbon atoms); the number of protons in an atom is called the
26
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Important
RemInderS
3. GAMMA RADIATION – high energy stream of photons (a bundle of
electromagnetic energy. It is the basic unit that makes up all light). Due to a
high energy nucleus, energy is given off and nucleus becomes stable.
4. POSITRON EMISSION – positively charged electron in which 1 proton turns into
a neutron. Atomic number down by 1.
5. ELECTRON CAPTURE – drawing of an electron into an atom’s nucleus.
6. BOMBARDMENT OF ALPHA PARTICLE: addition of alpha particle
7. PROTON – positive charge
8. NEUTRON - no charge or uncharged
•
•
REFERENCES
Name: _______________________________________________________
Grade Level & Section: __________________________________________
Date Submitted: (to be filled in by the subject instructor): _______________
Atomic Theory. (2020, August 22). Retrieved from
https://chem.libretexts.org/Bookshelves/Physical_and_Theoreti
cal_Chemistry_Textbook_Maps/Supplemental_Modules_(Physi
cal_and_Theoretical_Chemistry)/Atomic_Theory/Atomic_Theor
y
MODULE 3:
Rutherford Model. (2020, November 09). Retrieved from
https://www.britannica.com/science/Rutherford-model
Bohr’s Model of an Atom. (n.d.).
https://byjus.com/chemistry/bohrs-model/
Retrieved
Tear this activity sheet and submit on the scheduled date along with the
other activity (ies) the instructor may have asked the students to do on a
separate paper.
If you are sending something you’ve done online such as MS presentation
(s), pictures, pdfs and alike as an attachment, then you may send them to
my email at germanvertudez1211gmail.com following this format:
(SECTION_LASTNAME_FIRSTNAME_ACTIVITYNAME
e.g.
IC1MA_BINABAN_PRINCESS_SCAVENGERS HUNT), or send a digital copy
from your flash drive together with this activity sheet.
Assessment
RATING:
PART A: ATOMIC MASS AND ATOMIC NUMBER (35 PTS)
Supply the table with needed information.
NAME
Copper
Tin
Iodine
Uranium
Potassium
Lithium
Oxygen
Gold
Sulfur
from
27
SYMBOL
Z
29
53
92
8
79
A
119
127
P
29
50
E
N
35
146
39
7
16
197
32
3
4
16
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
a.
b.
PART B: MULTIPLE CHOICE (25 PTS)
Circle the letter that corresponds to your answers. (1 point each)
1.
2.
A proton has approximately the same mass as ____.
a. a neutron.
c. an alpha particle
b. a beta particle
d. an electron
4.
Which symbols represent atoms that are isotopes?
a. C-14 and N-14
c. O-16 and O-18.
b. I-131 and I-131
d. Rn-222 and Ra-222
5.
Which atom contains exactly 15 protons?
a. P-32.
b. O-15
An ion with 5 protons, 6 neutrons, and a charge of 3+ has an atomic number of
_____.
a. 5.
c. 6
b. 8
d. 11
7.
What is the mass number of an atom which contains 28 protons, 28 electrons, and
34 neutrons?
a. 28
c. 56
b. 62.
d. 90
c. Thomson.
d. Millikan
11. Mass number is equal to the_________.
a. number of protons + number of electrons
b. number of protons + number of neutrons.
c. number of neutrons + number of electrons
d. number of electrons
12. The development of the concept that elements have isotopes helps explain why
atomic _______.
a. masses are not whole numbers.
b. masses differ from atomic numbers
c. nuclei are charged
d. nuclei are neutral
c. S-32
d. N-15
6.
Electron was discovered by___________.
a. Chadwick
b. Rutherford
c. electrons
d. Neutrons
10. Carbon-12 atom has_______________.
a. 6 electrons, 12 protons, 6 neutrons
b. 12 electrons, 6 protons, 6 neutrons
c. 6 electrons, 6 protons, 6 neutrons.
d. 18 electrons, 6 protons and 6 neutrons
When alpha particles are used to bombard gold foil, most of the alpha particles pass
through undeflected. This result indicates that most of the volume of a gold atom
consists of ____.
a. deuterons (H-2)
c. neutrons
b. protons
d. unoccupied space.
3.
8.
9.
Compared to the charge and mass of a proton, an electron has ______.
a. the same charge and a smaller mass
b. the same charge and the same mass
c. an opposite charge and a smaller mass.
d. an opposite charge and the same mass
Protons
nucleus.
13. In an experiment by Rutherford, the deflection of alpha particles backward when
shot at gold foil indicated _____.
a. all the positive charge and most of the mass of the atom was concentrated in a
small volume.
b. all of the mass and charge of the atom were contained in the same region
c. the alpha particle was very light and positively charged
d. all matter was continuous and impenetrable
14. When Rutherford bombarded gold foil with positively charged alpha particles, most
of the particles went through but some were deflected back. Rutherford concluded
that atoms _____.
a. have negative charges
c. contain neutral particles
Almost the entire mass of an atom is concentrated in the _____.
28
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
b.
are solid spheres
d. have positive nuclei.
b.
15. Which one of the following would contain the other four?
a. compound.
d. molecule
b. Element
e. atom
c. nucleus
other elements
d. mixtures
23. A characteristic of the nucleus of an atom is that ________.
a. it has positive charge equal to the atomic number.
b. its mass is small compared to the mass of the atom
c. it contains all the electrons of the atom
d. the electrons and the protons balance each other
16. Positive ions are formed from neutral atoms by the loss of _____.
a. Neutrons
c. protons
b. Electrons.
d. energy
24. Atom is to element, as molecule is to ________.
a. Mixture
c. atom
b. Compound.
d. solution
17. The particles which have equal but opposite electrical charges are ________.
a. electrons and neutrons
c. protons and hydrogen nuclei
b. protons and electrons.
d. neutrons and protons
25. Scientists visualize the way atoms behave by using ________.
a. Pictures
c. genes
b. Microscopes
d. models.
18. Which of the following statement is correct?
a. A proton is a positively-charged particle in the nucleus.
b. A proton is a negatively-charged particle in the nucleus
c. Neutrons and protons are charged particles in the nucleus
d. None of these
19. The neutral atoms in a given sample of an element could have different ________.
a. number of protons
c. mass numbers.
b. atomic numbers
d. number of electrons
20. Which of the following is a characteristic of an atom?
a. atoms are alike
b. An atom is the smallest particle of matter
c. An atom is the smallest particle of an element with the properties of the
element.
d. All atoms have the same mass, which equals l a.m.u
21. The number of protons in the nucleus of an atom determines the ________.
a. mass number
c. atomic number.
b. atomic mass
d. isotopic mass
22. Elements can be combined chemically to form _______.
a. Compounds.
c. oxygen
29
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
THIRD QUARTER
Here is a summary of the types of orbitals and how many electrons each can
contain:
Elements in the Periodic
Table
Table 1: ORBITALS AND ELECTRON CAPACITY
At the end of this module, you are expected to:
•
•
INCLUSIVE DATE:
•
•
Classify and group the elements in the periodic table,
Write the electron configuration, orbital diagram and noble
gas notation,
Explain the basis of arrangement of elements in the PTE, and
Draw the Lewis dot structure.
KEY QUESTIONS:
• Why do we assign quantum numbers to electrons?
• Why are the rules followed in making electron notations?
• What is the significance of observing the specific arrangement
of elements in the periodic table?
Overview
Electron
configurations
are the summary of where the
electrons are around a nucleus. For
example,
the
electron
configuration of Potassium atom as
shown in the left side depicts the
arrangement
of
electrons
distributed among the orbitals and
subshells.
Fig 2. How to Write Electron Configuration
30
The symbols used for writing the
electron configuration start with the
shell number (n) followed by the type of
orbital and finally the superscript
indicates how many electrons are in the
orbital. Looking at the periodic table, you
can see that Oxygen has 8 electrons. So,
oxygen's electron configuration would
be 1s^2 2s^2 2p^4.
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
RULES OF ORBITAL DIAGRAM
1. Pauli Exclusion Principle
•
states that, in an atom or molecule, no
two electrons can have the same four
electronic quantum numbers. The atom
has 2 bound electrons and they occupy
the outermost shell with opposite spins.
•
The Aufbau principle states that
electrons fill lower-energy atomic
orbitals before filling higher-energy ones
(Aufbau is German for "building-up").
2. Aufbau Principle
3. Hund’s Rule
•
Fig 3. SPDF Blocks and Parts of Periodic Table of Elements
31
Every orbital in a
subshell
is
singly
occupied with one
electron before any one
orbital
is
doubly
occupied,
and
all
electrons
in
singly
occupied orbitals have
the same spin.
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
There are two main exceptions to electron configuration: chromium and
copper. In these cases, a completely full or half full d sub-level is more stable than
a partially filled d sub-level, so an electron from the 4s orbital is excited and rises
to a 3d orbital.
WAYS ON WRITING ELECTRON CONFIGURATION
1. SPDF NOTATION
The configuration notation provides an easy way for scientists to write and
communicate how electrons are arranged around the nucleus of an atom.
Calcium: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2
LEWIS DOT STRUCTURE
2. ORBITAL DIAGRAM
Orbital diagrams are pictorial descriptions of the electrons in an atom.
CALCIUM:
A Lewis Structure is a very simplified representation of the valence shell
electrons in a molecule. It is used to show how the electrons are arranged around
individual atoms in a molecule. Electrons are shown as "dots" or for bonding electrons
as a line between the two atoms. The goal is to obtain the "best" electron
configuration, i.e., the octet rule and formal charges need to be satisfied.
3. NOBLE GAS NOTATION
A noble gas configuration of an atom consists of the elemental symbol of the
last noble gas prior to that atom, followed by the configuration of the
remaining electrons.
CALCIUM: [AR] 4s^2
EXCEPTIONS
In the d block, specifically the groups containing Chromium and Copper,
there is an exception in how they are filled. Here are the actual configurations:
Fig 4. Lewis Dot Structure
32
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Lewis structures (also known as
Lewis dot structures or electron dot
structures) are diagrams that represent the
valence electrons of atoms within a
molecule. Calcium belongs in Group 2A or
Alkaline Earth Metal therefore, it has 2
valence electrons. Valence electrons are
used for chemical bonding.
STEPS IN DRAWING LEWIS DOT STRUCTURE FOR
32 valence electrons
- 8 valence electrons for the central atom
24 remaining valence electrons to distribute to the terminal atoms
3. DISTRIBUTE THE REMAINING VALENCE
ELECTRONS
Add the remaining electrons to the
terminal atoms until their octets are
filled.
Remaining 24e were already used
Used all 32 electrons
SiBr4
4. CHECK THE OCTET RULE OF ATOM
Every atom must have 8e around it. Except for Helium (H), it only needs 2e to
be stable following the duplet rule.
1. COUNT ALL THE VALENCE ELECTRON
Silicon is in Group 4A with 1 atom
Bromine is in Group 7A with 4 atoms
4X1 + 7x4 = 32 valence electrons
5. MAKE MULTIPLE BONDS (if required)
When all electrons are used and some atom does not complete octet
4x1 + 6x1 + 7x2 = 24e
Used all 24e
Move electron from Oxygen
to make double bond with
Carbon to complete the
Octet Rule.
2. ARRANGE AND CONNECT ATOMS
Distribute the valence electrons for each atom. Atoms with least
electronegativity is the central atom
NOTE:
33
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
KEY IDEAS
•
•
•
•
•
•
•
•
•
•
In the modern periodic table, the elements are listed in order of increasing
atomic number or proton number. Dmitri Mendeleev is the Father of
Modern Periodic Table of Elements (118 where 94 of them were naturallyoccurring)
The rows of the periodic table are called periods or series.
The columns of the periodic table are called groups or families.
The three broad categories of elements are metals, nonmetals, and
metalloids. Most elements are metals. Nonmetals are located on the
righthand side of the periodic table. Metalloids have properties of both
metals and nonmetals.
Molecules are aggregate of at least 2 atoms held together by chemical forces
and are uncharged.
1. Monoatomic: Na+ and Cl-1
2. Diatomic: N2, O2, H2, F2, I2, Cl2, Br2
3. Polyatomic: O3, H2O, NH4, CO
An ion is a charged atom or molecule: cation (+) and anion (-)
Periodic Law, a law stating that the elements, when listed in order of their
atomic numbers (originally, atomic weights), fall into recurring groups, so
that elements with similar properties occur at regular intervals.
Periodicity refers to trends or recurring variations in element properties with
increasing atomic number.
Electronegativity is an atoms ability to pull electrons towards itself.
Electronegativity generally increases as you move from left to right across a
period and decreases as you move down a group. The most electronegative
element is Fluorine.
•
The Electron Affinity of an element is the amount of energy gained or
released with the addition of an electron. The electronegativity and Electron
Affinity increases in the same pattern in the periodic table. Left to right and
bottom to top.
•
There are three main types of chemical formulas:
1. Molecular formulas show the number of each
type of atom in a molecule
2. Empirical formulas show the simplest wholenumber ratio of atoms in a compound
3. Structural formulas show how the atoms in a
molecule are bonded to each other.
TRENDS AND DEVELOPMENT
1. JOHANNE DOBEREINER (1829)
Model of Triads, he observed that elements with similar physical and
chemical properties fall into groups of three. One of these triads included
chlorine, bromine, and iodine; another consisted of calcium, strontium, and
barium. It states that the atomic number of the intermediate element is the
approximate average of atomic weight or density of the two elements: For
example: Li, Na, K
Z: 3Li, 11Na, 19K = 3+19 / 2 = 11
A: 7Li, 23Na, 39K = 7+39 / 2 = 23
2. JOHN NEWLANDS (1863)
Law of Octaves, states that if the chemical elements are arranged according
to increasing atomic weight, those with similar physical and chemical
properties occur after each interval of seven elements.
Ionization energy is the amount of energy required to remove an electron
from an atom. The more electronegative the element, the higher the
ionization energy. The ionization energy decreases from top to bottom in
groups, and increases from left to right across a period. Thus, helium has the
largest first ionization energy, while francium has one of the lowest.
3. DMITRI MENDELEEV (1869)
34
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
He corrected the atomic weight of Be, Ir, and U, and correctly predicted the
periodic position of unknown elements such as Sc, Ge, and Ga.
REFERENCES
4. LOTHAR MEYER (1869)
Arranged the PTE at increasing atomic mass, and periodicity or the physical
and chemical properties of elements.
Lewis
structures.
(2020,
August
16)
Retrieved
from
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistr
y_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Che
mistry)/Physical_Properties_of_Matter/Atomic_and_Molecular_Properties/L
ewis_Structures
5. HENRY MOSELEY (1913)
Moseley postulated that each successive element has a nuclear charge
exactly one unit greater than its predecessor. This determined the actual
atomic number through his works on x-rays by exposing them to alpha (+)
rays confirming that the elements were indeed arranged according to
increasing atomic number.
Lewis
Structures.
(n.d.)
Retrieved
from
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/lewis.html#:
~:text=Step%201%3A%20Determine%20the%20total,valence%20electrons%
20as%20nonbonding%20electrons.
6. GLENN T. SEABORG (1944)
He co-discovered the 10 new elements (transuranium elements, or
elements with atomic number greater than 92). He put 14 elements below
lanthanide series and named Sb after him while he was still alive.
Electron
Configurations.
(n.d.)
Retrieved
https://www.chem.fsu.edu/chemlab/chm1045/e_config.html
from
Santiago, K.S.., Silverio, A.A., (2016). Exploring Life through Science Series:
Senior High School Physical Science. Elements in the Periodic Table. pp 26-45.
Quezon City, Phoenix Publishing House Inc.
35
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
9. He proposed the Plum Pudding atomic model
10. He determined the actual Z through his works on x-ray by exposing the
elements to alpha rays.
11. It states that no two electrons in an atom can have the same four
quantum numbers.
12. It says that the lower energy level fill before the higher energy level in
order of increasing atomic number.
13. It suggests that one electron goes into each orbital before they start
to double up.
14. The chart that organizes all of the elements.
15. The columns in the periodic table.
16. The rows in the periodic table.
17. Ability to attract electrons toward itself when bonded.
18. It states that the elements with similar and physical properties occur
after each interval of 7 elements.
19. It states that the atomic weight of the middle element is nearly the
same as average of the atomic weights of other two elements.
20. It is the number of naturally occurring elements found in the Periodic
Table.
21. It is the other name for atomic number.
22. This is the other name given for the lanthanides and actinides series.
23. It happens to the atomic number as you move from left to right and
top to bottom of the Periodic Table.
24. These are elements that have same metallic and nonmetallic
properties.
25. He has correctly predicted the existence of several unknown elements
like Ga, Ge, Sc.
26. These are atoms of the same element that may have different number
of neutrons.
27. There are three parts of an atom: protons, neutron, and electrons.
__________ have a positive charge, _________ have a negative
charge, and ___________ possess no net charge. __________ are the
smallest parts of the atom.
28. The __________ or proton number (symbol Z) of a chemical element
is the number of protons found in the nucleus of an atom while the
______________ (symbol A), also called atomic mass number or
nucleon number, is the total number of protons and neutrons
(together known as nucleons) in an atomic nucleus.
Important
RemInderS
•
•
Tear this activity sheet and submit on the scheduled date along with the
other activity (ies) the instructor may have asked the students to do on a
separate paper.
If you are sending something you’ve done online such as MS presentation
(s), pictures, pdfs and alike as an attachment, then you may send them to
my email at germanvertudez1211gmail.com following this format:
(SECTION_LASTNAME_FIRSTNAME_ACTIVITYNAME
e.g.
IC1MA_BINABAN_PRINCESS_SCAVENGERS HUNT), or send a digital copy
from your flash drive together with this activity sheet.
Name: _______________________________________________________
Grade Level & Section: __________________________________________
Date Submitted: (to be filled in by the subject instructor): _______________
RATING:
MODULE 4:
Assessment
PART A: IDENTIFICATION (30 PTS)
Identify which area of Biology the statement describes. Write your answer before the
number. (1 point each)
1. He discovered neutron in an atom.
2. He said that objects are composed of indivisible units he called
atomos.
3. He discovered electron.
4. He is the Father of Modern Periodic Table.
5. He suggested that elements are arranged according to the Model of
Triads.
6. He proposed periodicity which states that elements are arranged
according to their physical and chemical properties.
7. He proposed the Law of Octaves.
8. He’s the chemist whose an element was named after while still alive.
36
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
29. Elements in the Periodic Table are arranged according to their
__________ atomic number.
30. It states that the atomic weight of the intermediate element is the
approximate average of the sum of the two elements.
c) A positive and a negative charge
d) A positive charge
2.
Neutrons are in the nucleus of the atom. A neutron has
a) A positive charge
b) No charge
c) A negative charge
d) Twice as much positive charge as a proton.
3.
An electron is in a region outside the nucleus. An electron
a) is larger than a proton and has no charge
b) Has less mass than a proton and has a negative charge
c) Is smaller than a proton and has no charge
d) Has a positive charge.
4.
A hydrogen atom is made up of one proton and one electron. The
electron stay near each other because
a) Positive and negative charges repel
b) Positive and positive charges repel
c) Positive and negative charges attract
d) Two negatives make a positive
5.
The atomic number of an atom is
a) The mass of the atom
b) The number of protons added to the number of neutrons
c) The number of protons
d) Negatively charged
PART D: LEWIS DOT STRUCTURE
6.
Draw the Lewis dot structure for the following.
1. Lithium and Nitrogen
2. Oxygen and Chlorine
3. Carbon and Sulfur
4. Potassium and Iodine
5. Potassium and Oxygen
The atoms of the same element can have different isotopes. An isotope of an atom
a) Is an atom with a different number of protons
b) Is an atom with a different number of neutrons
c) Is an atom with a different number of electrons
d) Has a different atomic number
7.
The atomic mass of an element is
a) The average mass of all the isotopes of the element
b) A measure of the density of that element
c) The mass of the most common isotope of that element
d) The number of protons and electrons in the atoms of the element
8.
An element and an atom are different but related because
PART B: TABLE COMPLETION (20 PTS)
Supply the table with the needed information.
Element
Name
Symbol
Iodine
Chromium
Atomic
Number
Number
of Protons
53
Number of
Electrons
53
24
Cr +2
Iron
24
26
Li
Potassium
Atomic
Mass
19
52
56
5
24
Number
of
Neutrons
78
56
22
30
3
1
PART C: ELECTRON CONFIGURATION
Write the ground state electron configuration of the following neutral elements in: Orbital
Diagrams, SPDF notation, and Noble gas notation.
1. Beryllium
2. Nitrogen
3. Potassium
4. Manganese
5. Chlorine
PART D: MULTIPLE CHOICE (20 PTS)
Circle the letter that corresponds to your answers. (1 point each)
1.
Protons are located in the nucleus of the atom. A proton has
a) No charge
b) A negative charge
37
proton and
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
a) A particular element is made up of many different types of atoms
b) A molecule is the same as an atom
c) An element is made up of all the same type of atom
d) An element is smaller than an atom
9.
c) by increasing atomic number and similar properties
d) in alphabetical order
15. The elements in the present periodic table are arranged according to their:
a) atomic masses
b) mass number
c) atomic number
d) atomic weight
The periodic table shows that a carbon atom has six protons. This means that a carbon
atom also has
a) Six electrons
b) Six neutrons
c) More protons than electrons
d) An atomic mass that equals six
16. A vertical column in the periodic table is called a:
a) Valence
b) branch
c) group
d) period
10. The atomic number of Nitrogen is 7. The atomic mass is 14.01. This means that
a) All nitrogen atoms have exactly 7 neutrons.
b) A small percentage of nitrogen atoms have fewer than 7 neutrons
c) A small percentage of nitrogen atoms have more than 7 neutrons
d) Some nitrogen atoms have fewer than 7 electrons
17. A horizontal row in the periodic table is called a:
a) Shells
b) branch
c) group
d) period
18. Elements in the same group have:
a) similar symbols
b) the same number of neutrons
c) the same number of valence electrons
d) the same number of electrons
11. Electrons are in regions around the nucleus called energy levels. The first energy level
a) Is furthest from the nucleus
b) Is closest to the nucleus
c) Holds the most electrons
d) Needs more than two electrons to fill it up
19. Elements in the same period have:
a) the same number of neutrons
b) gradually changing properties
c) similar symbols
d) identical chemical properties
12. Neon has 10 protons and 10 electrons. The electrons fill the energy levels in Neon
like this:
a) 2 in the first, 2 in the second, and 6 in the third
b) 4 in the first, 4 in the second, and 2 in the third
c) 2 in the first, 4 in the second, and 4 in the third
d) 2 in the first, and 8 in the second
20. How many valence electrons does boron have?
a) One
b) two
c) three
d) four
13. The atoms in a column of the periodic table all have
a) The same abbreviation
b) The same number of energy levels
c) The same number of electrons
d) The same number of electrons in the outer energy level
14. Dmitri Mendeleev organized the chemical elements:
a) by number of electrons
b) by increasing atomic weight and similar properties
38
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
THIRD QUARTER
Forces of Attraction and
Chemical Bonds
At the end of this module, you are expected to:
•
•
INCLUSIVE DATE:
•
Determine if a molecule is polar or nonpolar
Describe the types of intermolecular and intramolecular
forces, and
Explain how the uses of some materials depend on their
properties.
have similar affinity for electrons and neither has a tendency to donate them,
they share electrons in order to achieve octet configuration and become more
stable.
a. A nonpolar covalent bond is formed between same atoms or
atoms with very similar electronegativities—the difference in
electronegativity between bonded atoms is less than 0.5.
b. A polar covalent bond is formed when atoms of slightly different
electronegativities share electrons. The difference in
electronegativity between bonded atoms is between 0.5 and 1.9.
Hydrogen chloride, HCl; O - H bonds in water, H2O and hydrogen
fluoride (HF), are all examples of polar covalent bonds.
KEY QUESTIONS:
• Why are most atoms held by chemical bonds?
• Did the rule “like dissolves like” hold true for all solutions? If
not, what could be the possible reasons for the observed
deviations?
Overview
There are two kinds of forces, or attractions, that operate in a molecule—
intramolecular (forces that hold atoms together within a molecule) and intermolecular
(forces that exist between molecules).
TYPES OF INTRAMOLECULAR FORCES OF ATTRACTION:
1. IONIC BOND: This bond is formed by the complete transfer of valence
electron(s) between atoms. It is a type of chemical bond that generates two
oppositely charged ions. In ionic bonds, the metal loses electrons to become
a positively charged cation, whereas the nonmetal accepts those electrons to
become a negatively charged anion.
2. COVALENT BOND: This bond is formed between atoms that have similar
electronegativities—the affinity or desire for electrons. Because both atoms
Fig 1. Ionic Bond, Nonpolar Covalent Bond, and Polar Covalent Bond (from top to bottom)
39
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
The electronegativity difference of two atoms determines their bond type.
If the electronegativity difference is more than 1.7, the bond will have an ionic
character. If the electronegativity difference is between 0.4 and 1.7, the bond will
have a polar covalent character.
as possible. The main geometries without lone pair electrons are: linear, trigonal,
tetrahedral, trigonal bipyramidal, and octahedral.
Fig 2. Bond Polarity
Lewis structure helps determine the three-dimensional arrangement of
atoms in a molecule. Such arrangement is referred to as molecular geometry.
The valence shell electron pair repulsion (VSEPR) model focuses on the
bonding and nonbonding electron pairs present in the outermost (valence) shell
of an atom that connects with two or more other atoms. Fundamentally, the
VSEPR model theorizes that regions of negative electric charge will repel each
other, causing them (and the chemical bonds that they form) to stay as far apart
Fig 3. Common Molecular Geometries
40
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
explains the exceptionally high boiling points and melting points of
compounds like water and hydrogen fluoride. Hydrogen bonding plays an
important role in biology; for example, hydrogen bonds are responsible for
holding nucleotide bases together in DNA and RNA.
3. LONDON DISPERSION FORCES, UNDER THE CATEGORY OF VAN DER WAAL
FORCES: These are the weakest of the intermolecular forces and exist between
all types of molecules, whether ionic or covalent—polar or nonpolar. The more
electrons a molecule has, the stronger the London dispersion forces are.
3. METALLIC BONDING: This
type of covalent bonding specifically
occurs between atoms of metals, in
which the valence electrons are free to
move through the lattice. This bond is
formed via the attraction of the mobile
electrons—referred to as sea of
electrons—and the fixed positively
charged metal ions.
TYPES OF INTERMOLECULAR FORCES THAT EXIST BETWEEN MOLECULES
1. DIPOLE-DIPOLE INTERACTIONS: These forces occur when the partially
positively charged part of a molecule interacts with the partially negatively
charged part of the neighboring molecule. The prerequisite for this type of
attraction to exist is partially charged ions—for example, the case of polar
covalent bonds such as hydrogen chloride (HCl). Dipole-dipole interactions are
the strongest intermolecular force of attraction.
2. HYDROGEN BONDING: This is a special kind of dipole-dipole interaction that
occurs specifically between a hydrogen atom bonded to either an oxygen,
nitrogen, or fluorine atom. The partially positive end of hydrogen is attracted
to the partially negative end of the oxygen, nitrogen, or fluorine of another
molecule. Hydrogen bonding is a relatively strong force of attraction between
molecules, and considerable energy is required to break hydrogen bonds. This
Fig 4. Types of Intermolecular Forces (from top to bottom)
The strength of a substance’s IMFA determines many of its physical properties, inc
(Br2) that boils very easily have extremely weak IMFA, and water boils at higher point poss
41
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Fig 5. Comparison between Ionic andREFERENCES
Molecular (Covalent) Bonds
42
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Intramolecular and intermolecular forces. (2021). Retrieved from
https://www.khanacademy.org/science/class-11-chemistryindia/xfbb6cb8fc2bd00c8:in-in-states-of-matter/xfbb6cb8fc2bd00c8:in-inintermolecular-forces/a/intramolecular-and-intermolecular-forces
Important
RemInderS
•
Molecular
Geometry.
(n.d.).
Retrieved
from
https://courses.lumenlearning.com/boundless-chemistry/chapter/moleculargeometry/
•
Santiago, K.S.., Silverio, A.A., (2016). Exploring Life through Science Series:
Senior High School Physical Science. What Ties Us? pp 48-75. Quezon City,
Phoenix Publishing House Inc.
Tear this activity sheet and submit on the scheduled date along with the
other activity (ies) the instructor may have asked the students to do on a
separate paper.
If you are sending something you’ve done online such as MS presentation
(s), pictures, pdfs and alike as an attachment, then you may send them to
my email at germanvertudez1211gmail.com following this format:
(SECTION___LASTNAME___FIRSTNAME___ACTIVITYNAME (for example:
IC1MA___BINABAN___PRINCESS___SCAVENGERS HUNT), or send a digital
copy from your flash drive together with this activity sheet.
Name: _______________________________________________________
Grade Level & Section: __________________________________________
Date Submitted: (to be filled in by the subject instructor): _______________
RATING:
MODULE 5:
Assessment
PART A: IDENTIFICATION (10 PTS)
Identify which area of Biology the statement describes. Write your answer before the
number. (1 point each)
1.
2.
3.
4.
5.
6.
43
It is a state or a condition of an atom or a molecule having positive and also
negative charges, especially in case of magnetic or an electrical pole.
The force that holds together the atoms making up a molecule or compound.
The force that acts between stable molecules or between functional groups of
macromolecules.
This is a chemical bond wherein there is a transfer of an electron from one atom to
another.
This bond allows the sharing of electrons.
These are attractive forces between the positive end of one polar molecule and the
negative end of another polar molecule.
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
7.
This is a special type of dipole-dipole attraction between molecules, not a covalent
bond to a hydrogen atom.
8. This is a temporary attractive force that results when the electrons in two adjacent
atoms occupy positions that make the atoms form temporary dipoles
9. It is a type of chemical bond that is formed when electrons are shared equally
between two atoms.
10. This a covalent bond in which the atoms are distributed asymmetrically.
Identify the IMFA present between molecules of each compound. Rank (from 1 – 3) the
compounds in terms of increasing strength of IMFA (15 points)
COMPOUND
Water
PART B: TABLE COMPLETION (35 POINTS)
H2O
LEWIS STRUCTURE
(3pts)
BOILING
POINT
IMFA Between
Molecules
(1pt)
RANK (1pt)
100 °C
Complete the table with the information as needed (20 points)
MOLECULE
LEWIS STRUCTURE
(2pts)
MOLECULAR
GEOMETRY (1pt)
POLARITY
(polar or nonpolar, 1pt)
Ethanol
C2H6O
N2
78.37 °C
SO2
Hexane
C6H14
68 °C
BF3
PART C: MULTIPLE CHOICE TEST (15 PTS)
Circle the letter that corresponds to your answer (1 point each)
1. Chemical bonding results in __________.
a. increased stability.
c. lower bond lengths
b. decreased stability
d. lower density
HCN
CH2F2
2. The shielding effect is the tendency of inner energy level electrons to block
the attraction of the nucleus for __________.
a. Neutrons
c. other atoms
44
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
b. Protons
d. valence electrons.
3. The CO molecule is __________.
a. linear and polar.
c. tetrahedral and nonpolar
b. linear and nonpolar
d. trigonal planar and polar
d. The interactions among the molecules in molecular solids are generally
stronger than those among the particles that define either covalent or
ionic crystal lattices.
10. Which one of the following classifications is incorrect?
a. H2O(s), molecular solid
c. C4H10(s), molecular solid
b. SiC(s), covalent solid
d. S(s), metallic solid.
4. A solute is most likely to be highly soluble in a solvent if the solute is _____
and the solvent is ______.
a. ionic or polar, non-polar
c. ionic or polar, polar.
b. non-polar, polar
d. non-polar, ionic
11. Which is classified as nonpolar covalent?
a. the H-I bond in HI
c. the H-S bond in H2S
b. the N-Cl bond in NCl3.
d. the N-H bond in NH3
5. The boiling point of CH4 is much lower than that of HF. This is because:
a. of ion-dipole interactions in CH4. c. HF is more polarizable.
b. of hydrogen bonding in HF.
D. CH4 is polar.
12. Which one of the compounds below is most likely to be ionic?
a. ScCl3.
c. NO2
b. CCl4
d. ClO2
6. The intermolecular forces of attraction in the above substances is described
by which of the following:
a. ion-dipole forces
b. dispersion (or London) forces.
c. dipole-dipole forces (permanent dipoles)
d. gravitational forces
13. Choose the molecule that is incorrectly matched with the electronic geometry
about the central atom.
a. CF4 – tetrahedral
c. BeBr2 - linear
b. H2O – tetrahedral
d. PF3 – pyramidal.
7. What type of intermolecular forces are due to the attraction between
temporary dipoles and their induced temporary dipoles?
a. metallic bond
c. London dispersion.
b. hydrogen bond
d. ionic bond
14. Which of the following four molecules are polar: PH3 OF2 HF SO3?
a. all except SO3.
c. only HF
b. only HF and OF2
d. all of these
15. Which of the following is an ionic compound?
a. CS2
c. MgCl2.
b. OF2
d. SO2
8. What type of interparticle forces holds liquid N2 together?
a. ionic bonding
c. London forces.
b. dipole-dipole interaction
d. covalent bonding
9. Which statement is false?
a. Molecular solids generally have lower melting points than covalent solids
b. Metallic solids exhibit a wide range of melting points because metallic
bonds cover a wide range of bond strength
c. Most molecular solids melt at lower temperatures than metallic solids
45
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Table 1: DIFFERENCE BETWEEN ORGANIC AND INORGANIC COMPOUNDS
THIRD QUARTER
Naming and Chemical
Compounds
At the end of this module, you are expected to:
•
INCLUSIVE DATE:
•
Convert between the chemical formula of an ionic compound and its
name, and
Identify the molecular or covalent formula of a compound given
either its name or structural formula.
KEY QUESTIONS:
• How are chemical compounds named?
• Why do we need rules in naming and writing compounds?
• How important are elements and compounds in our daily life?
Overview
Chemical
compounds are any
substance(s)
composed
of
identical molecules
consisting of atoms
of two or more
chemical elements.
ORGANIC COMPOUNDS
INORGANIC COMPOUNDS
Organic compounds are characterized
by the presence of carbon atoms in
them
Organic compounds consisting of
hydrogen, oxygen, carbon, and their
other derivatives
Organic compounds are said to be
more volatile and also highly
inflammable
These compounds exist in the form of
solids, gases, and liquids.
These are insoluble in water
Most inorganic compounds do not
have carbon atoms in them (some
exceptions do exist)
They do not possess hydrogen or
oxygen and their derivatives
These compounds have the carbonhydrogen bonds
Organic compounds are mainly found
in most of the living things
Organic compounds form covalent
bonds
In most of the aqueous solutions, these
are poor conductors of heat and
electricity
Examples of organic compounds
include fats, nucleic acids, sugars,
enzymes, proteins and hydrocarbon
fuels
These have relatively low melting
points and boiling points.
These are biological and more complex
in nature
Organic compounds cannot make salts
The rate of reaction is slow in organic
compounds
The primary difference that lies between these organic compounds and
inorganic compounds is that organic compounds always have a carbon atom while
most of the inorganic compounds do not contain the carbon atom in them.
46
These compounds are not inflammable
and are non-volatile in nature
These exist as solids
These are soluble in water and also
non-soluble in some of the organic
solutions
These do not have the carbonhydrogen bonds
These compounds are found in nonliving things
Inorganic compounds form ionic bonds
between the atoms of molecules
In aqueous solutions, these are known
to be good conductors of heat and
electricity
The example for inorganic compounds
includes non-metals, salts, metals,
acids, bases, substances which are
made from single elements
These have low melting and boiling
points compared to organic
compounds
These are of mineral and not much
complexity in nature
Inorganic compounds can make salts
Inorganic compounds have a high rate
of reaction
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
CLASSIFICATIONS OF CHEMICAL COMPOUNDS:
NAMING RULES FOR IONIC COMPOUNDS:
1. Always write the cation first then the anion: Magnesium (Mg^+2) and Oxygen
(O^-2)
2. Write the compound using subscript (crisscross the oxidation number or
charge) so that the net charge is zero: Mg2O2 (molecular formula) to MgO
(empirical formula)
3. Change the ending of the second element to – ide: MgO (magnesium + oxygen
= magnesium oxide)
4. If the cation is transitional metal (groups 3-12), write the charge as Roman
Numeral following the rules of Stock System by Alfred Stock:
Lead (Pb^+2, +4) and Oxygen (O^-2)
Lead (II) or Pb (II) and Lead (IV) or Pb (IV)
5. Change the ending of the second element to – ide.
Lead (Pb^+2) + Oxygen (O^-2) = PbO
STOCK SYSTEM: Lead (II) Oxide
CLASSICAL NOMENCLATURE: Plumbous oxide
Lead (Pb^+4) + Oxygen (O^-2) = Pb2O4 or PbO2
STOCK SYSTEM: Lead (IV) Oxide
CLASSICAL NOMENCLATURE: Plumbic oxide
1. Based on Specific Elements Present
- For example, oxides contain one or more oxygen atoms, hydrides contain
one or more hydrogen atoms, and halides contain one or more halogen
(Group 17) atoms.
6.
NOTE: In classical nomenclature by Guy deMoreau (1787), Latin root is
use for the cation: OUS (lower charge) and IC (higher charge(s)).
Fig 1. Comparison between Ionic and Molecular (Covalent) Bonds
2. Based on Types of Bonds
a. IONIC COMPOUNDS
- contain ions and are held together by the attractive forces among the
oppositely charged ions. For example, common salt (sodium chloride,
NaCl).
Fig 2. Common Ways of Naming Compounds
47
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
b. MOLECULAR (COVALENT) COMPOUNDS
- contain discrete molecules, which are held together by sharing
electrons (covalent bonding). For example, H2O, CH4, HF
3. Based on Reactivity
- It is the types of chemical reactions that the compounds are likely to
undergo.
a. ACIDS are compounds that produce H+ ions (protons) when dissolved
in water to produce aqueous solutions.
b. BASES, on the other hand, are proton acceptors.
c. OXIDATION-REDUCTION REACTIONS constitute another important
class of chemical reactions.
NAMING RULES FOR ACIDS AND BASES:
Acids are named by the anion they form when dissolved in water. Depending on what
anion the hydrogen is attached to, acids will have different names.
NAMING ACIDS:
1.
When the anion ends in –ide, the acid name begins with the prefix hydro-. The
root of the anion name goes in the blank (chlor for chloride), followed by the suffix
–ic.
-
Fig 3. Molecular Bonding Comparison
2.
When the anion ends in –ate, the name of the acid is the root of the anion
followed by the suffix –ic. There is no prefix.
H2SO4 is sulfuric acid (not sulfic) because SO4^2- is the sulfate ion.
3.
When the anion ends in –ite, the name of the acid is the root of the anion followed
by the suffix –ous. Again, there is no prefix.
HNO2 is nitrous acid because NO2- is the nitrite ion.
NAMING RULES FOR MOLECULAR COMPOUNDS:
1. Change the ending of the second element to – ide.
2. Use Greek prefixes to name the first and second elements indicating the
number of their atoms present:
1 – mono
5 – penta
9 – nona
2 – di
6 – hexa
10 – deca
3 – tri
7 – hepta
11 – hendeca
4 – tetra
8 – octa
12 – dodeca
HCl is hydrochloric acid because Cl- is the chloride ion. HCN is hydrocyanic acid because
CN- is the cyanide ion.
NAMING BASES:
1.
3. If there is only one of the first element, you can drop the prefix. For
example, CO is carbon monoxide, not monocarbon monoxide.
4. If there are two vowels in a row that sound the same once the prefix is
added (they “conflict”), the extra vowel on the end of the prefix is
removed. For example, one oxygen would be monooxide, but instead
it’s monoxide. The extra o is dropped.
Most strong bases contain hydroxide, a polyatomic ion.
Strong bases are named following the rules for naming
ionic compounds. For example, NaOH is sodium hydroxide,
KOH is potassium hydroxide, and Ca (OH)2 is calcium
hydroxide.
-
Weak bases made of ionic compounds are also named
using the ionic naming system. For example, NH4OH is
ammonium hydroxide.
48
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
2.
Weak bases are also sometimes molecular compounds or organic compounds because
they have covalent bonds. Therefore, they are named following the rules for molecular
or organic compounds.
-
REFERENCES
For example, methyl amine (CH3NH2) is a weak base. Some weak bases have “common”
names. For example, NH3 is called ammonia; its name isn’t derived from any naming
system.
Zumdahl, S. (n.d.). Chemical Compounds. In Britannica.com dictionary. Retrieved
from https://www.britannica.com/science/chemical-compound
Difference Between Organic and Inorganic Compounds. (n.d.). Retrieved from
https://byjus.com/chemistry/difference-between-organic-and-inorganiccompounds/#:~:text=Organic%20and%20inorganic%20compounds%20form%2
0one%20of%20the%20primary%20basis%20for%20chemistry.&text=The%20pr
imary%20difference%20that%20lies,the%20carbon%20atom%20in%20them.
Binary molecular (covalent) compounds. (n.d.). Retrieved from
https://www.britannica.com/science/chemical-compound/Binary-molecularcovalent-compounds
Stoichiometry
of
Chemical
Reactions.
(n.d.).
Retrieved
from
https://courses.lumenlearning.com/chemistryformajors/chapter/writing-andbalancing-chemical-equations/#:~:text=Show%20Selected%20Solutions,1.,law%20of%20conservation%20of%20matter.
Atoms,
Molecules,
and
Ions.
(n.d.).
Retrieved
from
https://courses.lumenlearning.com/boundless-chemistry/chapter/chemicalformulas/
Atoms,
Molecules,
and
Ions.
(n.d.).
Retrieved
from
https://courses.lumenlearning.com/boundless-chemistry/chapter/namingcompounds/
Acids - Naming and Formulas. (2021, January 14). Retrieved from
https://flexbooks.ck12.org/cbook/ck-12-chemistry-flexbook2.0/section/7.12/primary/lesson/naming-acids-chem
Fig 5. Most Common Acids and Bases
49
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
8.
This naming system of chemical compounds states that the oxidation states of some
or all of the elements in a compound are indicated in parentheses by Roman
numerals.
9. These compounds will contain a carbon atom (and often a hydrogen atom, to form
hydrocarbons)
10. Any substance in which two or more chemical elements (usually other than carbon)
are combined, nearly always in definite proportions.
Important
RemInderS
•
•
Tear this activity sheet and submit on the scheduled date along with the
other activity (ies) the instructor may have asked the students to do on a
separate paper.
If you are sending something you’ve done online such as MS presentation
(s), pictures, pdfs and alike as an attachment, then you may send them to
my email at germanvertudez1211gmail.com following this format:
(SECTION___LASTNAME___FIRSTNAME___ACTIVITYNAME (for example:
IC1MA___BINABAN___PRINCESS___SCAVENGERS HUNT), or send a digital
copy from your flash drive together with this activity sheet.
PART B: TABLE COMPLETION (25 PTS)
Put a check mark to which the description is being referred to (15 points)
PROPERTIES
Crystalline solid at room temperature
Liquid, gas, or solid at room
temperature
Low melting point
High melting point
Conduct electricity as a liquid
Most have low solubility in water
Most have high solubility in water
Conduct electricity when dissolved in
water.
Bonding between metals and
nonmetals
Bonding between two or more
nonmetals
Sodium chloride (NaCl)
Hydrogen gas (H2)
Sharing of electrons
Transferring of electrons
Name: _______________________________________________________
Grade Level & Section: __________________________________________
Date Submitted: (to be filled in by the subject instructor): _______________
RATING:
MODULE 6:
Assessment
PART A: IDENTIFICATION (10 PTS)
Identify which area of Biology the statement describes. Write your answer before the
number. (1 point each)
1.
2.
3.
4.
5.
6.
7.
It is a chemical substance can be broken down into two or more simpler substances.
It is the complete transfer of valence electron(s) between atoms.
It is a covalent bond between two atoms where the electrons forming the bond are
unequally distributed.
It is formed when electrons are shared equally between two atoms.
____________ formulas show the simplest whole-number ratio of atoms in a
compound.
____________ formulas show the number of each type of atom in a molecule.
____________formulas show how the atoms in a molecule are bonded to each
other.
50
IONIC BONDING
COVALENT
BONDING
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Use your periodic table in identifying the elements present in each formula. Categorize them
as either metal or nonmetal then determine the type of bond each compound has. Number
1 and 2 are already provided (1 point for all correct answer(s) in a row) – 10 points.
FORMULA
METALS
NONMETALS
Sodium
Hydrogen Fluoride
or (mono)hydrogen
monofluoride)
Chlorine
HF
NaCl
C4H10
Al2O3
CBr4
Na2S
Sr3N2
H2S
MgCl2
NO2
C2H4
BaF2
Write the chemical formula/e for each of the following compounds (1 point each)
1. Nickel (II) chloride
2. Cuprous nitrate
3. Ammonium sulfate
4. Magnesium nitride
5. Mercury (I) sulfide
6. Carbon monoxide
7. Iron (II) oxide
8. Diphosphorus pentoxide
9. Hydrogen sulfate
10. Aluminum hydroxide
TYPE OF BOND
Covalent
Ionic
Combine each pair of ions to get the chemical formula, then name the compound (2 points
each)
COMPOUND
FORMULA
INDIVIDUAL IONS
PART C: NAMING AND WRITING CHEMICAL COMPOUNDS (50 PTS)
Write either the traditional or Stock System for each of the following compounds (1 point
each)
1. Na2S
2. NH4Cl
3.
4.
5.
6.
7.
8.
9.
10.
CuF
CuF2
PbSO4
Hg(NO3)2
Al2O3
N2O4
H2S
HClO2
51
Mg (+2)
F (-)
Ni (+2)
S (-2)
Ca (+2)
Br (-)
Al (+3)
P (-3)
COMPOUND
NAME
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
Co (+2)
NO2 (-)
Al (+3)
Na (+)
Cl (-)
K (+)
CrO4 (-2)
Fe (+3)
S (-2)
Fe (+3)
NO3 (-)
K (+)
CO3 (-2)
Ni (+2)
Br (-)
Cu (+2)
SO4 (-2)
Ba (+2)
PART C: MULTIPLE CHOICE TEST (15 PTS)
Circle the letter that corresponds to your answer (1 point each)
O (-2)
Ag (+)
SO4 (-2)
1. Which one of the formulas for ionic compounds below is incorrect?
a. SrCl2
b. Cs2S
c. AlCl3
d. Al3P2.
2. When naming a transition metal ion that can have more than one common ionic
charge, the numerical value of the charge is indicated by a ____.
a. Prefix
b. Suffix
c. Roman numeral following the name.
d. superscript after the name
3. The nonmetals in Groups 6A and 7A ____.
a. lose electrons when they form ions
b.
have a numerical charge that is found by subtracting 8 from the group
number
c. all have ions with a – 1 charge.
d. end in -ate
Cl (-)
4. Which of the following is NOT a cation?
a. iron(III) ion
b. sulfate.
52
c. Ca
d. mercurous ion
ASIAN INSTITUTE OF COMPUTER STUDIES (AICS)
5. An -ate or -ite at the end of a compound name usually indicates that the
compound contains ____.
a. fewer electrons than protons c. only two elements
b. neutral molecules
d. a polyatomic anion.
12. When naming acids, the prefix hydro- is used when the name of the acid anion
ends in ____.
a. -ide.
c. -ate
b. -ite
d. -ic
6. Which of the following is true about the composition of ionic compounds?
a. They are composed of anions and cations.
b. They are composed of anions only
c. They are composed of cations only
d. They are formed from two or more nonmetallic elements
13. ____ What is the name of H2SO3?
a. hyposulfuric acid
c. sulfuric acid
b. hydrosulfuric acid
d. sulfurous acid.
14. When the name of an anion that is part of an acid ends in -ite, the acid name
includes the suffix ____.
a. -ous.
c. -ate
b. -ic
d. -ite
7. Which element, when combined with fluorine, would most likely form an ionic
compound?
a. Lithium.
c. phosphorus
b. Carbon
d. chlorine
15. What is the correct name for Sn (PO ) ?
a. tritin diphosphate
c. tin(III) phosphate
b. tin(II) phosphate.
d. tin(IV) phosphate
8. Which of the following compounds contains the lead(II) ion?
a. a. PbO.
c. Pb2O
b. b. PbCl4
d. Pb2S
9. What is the correct formula for potassium sulfite?
a. KHSO3
c. K2SO3.
b. KHSO4
d. K2SO4
10. What type of compound is CuSO4?
a. monotomic ionic
c. polyatomic ionic.
b. polyatomic covalent
d. binary molecular
11. Molecular compounds are usually ____.
a. composed of two or more transition elements
b. composed of positive and negative ions
c. composed of two or more nonmetallic elements.
d. exceptions to the law of definite proportions
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