LEARNING MODULE IN GENERAL CHEMISTRY (ORGANIC)

Atoms, Molecules and Ions

MODULE
III

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Objectives:

1. applied the rules of basic laws of matter

2. identified the proton, neutron, electron and the subatomic particles
3. recognized isotopes and ions

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What we know today, what we accept as
laws that govern chemical processes, what
we hold as sacred and irrefutable facts
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are the result of a long history of

work, sweat and toil of scientists who
attempted to find explanations for their
observations and various phenomena
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around them.

Thanks to generations who preceded

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us, a long string of names in the Who’s Who

of science: Antoine Lavoiser, Joseph Proust,
Lord Ernest Rutherford, Niels Bohr, Marie
Curie, John Dalton, and other names whose
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works provided the foundation for our

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understanding of matter. This understanding

allows us to harness matter to benefit
humankind and to open the possibilities yet
unknown.

Dalton’s atomic theory is a case in

point. The idea of atoms was first suggested
by Democritus, an ancient Greek who lived in
the 4th B.C. however, his idea of the atom
could not support chemical phenomena.
John Dalton (1766-1844), an English chemist
and physicists was able to relate chemical
changes to the level of individual atoms and state his atomic theory.

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 LEARNING MODULE IN GENERAL CHEMISTRY (ORGANIC)

Lesson 1. Dalton’s Atomic Theory

4. All elements are composed of tiny indivisible particles called atoms.

5. Atoms of the same element are identical. The atoms of any one element are
different from those of another element.
6. Atoms of different elements can combine with one another in simple whole

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number ratios to form compounds.
7. Chemical reactions occur when atoms are separated, joined or rearranged.
However,atoms of one element are not changed into atoms of another by a

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chemical reaction.

Big word! Atom! What is an atom

anyway? Try to take a small piece

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of element Copper (Cu) and break
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possible. No matter how small it
gets, the smallest piece will retsin
the properties of copper. Imagine
that you contonue breaking it
down further until it can no longer
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be divided and each particle still

retains the properties of the
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original substance. This smallest

indivisible particle is known as the
atom. The word atom comes from
the Greek word atomos which
means indivisible or uncut.
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We now know that some of Dalton’s assumptions were not correct. However, his
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atomic theory was a historic step in understanding of chemical behavior. In spite of its
shortcomings, Dalton’s theory inspired a generation of chemists to conduct experiments and
clarify and refine our understanding of the fundamental partciles of nature.
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Dalton’s ideas gave birth to the folowing

laws:
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1. Law of Conservation of Mass. This

was postulated by French chemist
Lavoisier in 1785. It states that when a
chemical reaction takes place, there will
be no detectable change in masses of the
substance. Expressed in equation form:

Mass of reactants = Mass of the products

Example:

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 LEARNING MODULE IN GENERAL CHEMISTRY (ORGANIC)

5 grams of Magnesium Chloride completely reacted with 20 grams of Ammonium Oxalate to

produce Magnesium Oxalate and Ammonium Chloride. What could be the mass of
Ammonium Chloride if the mass of Magnesium Oxalate is 15 grams?

Reactants Products
Magnesium Chloride 5 grams Magnesium Oxalate 15 grams
Ammonium Oxalate 20 grams Ammonium Chloride ?
Total = 25 grams Total = ?

The Law of Conservation of Mass would explain that if the original reactants totaled
25 grams, then the products would be unchanged. That means that the final products also
weigh 25 grams. Since we are given the weight of Magnesium Oxalate as 15 grams, then we
can determine that the weight of Ammonium Chloride is 25 grams less 15 grams or 10

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grams.

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2.
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Law of Definite Proportion. This law states that when elements combine to form a given
compound, they do so in fixed and invariable ratio by weight.
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 LEARNING MODULE IN GENERAL CHEMISTRY (ORGANIC)

Example:
a. Calcium Chloride has the molecular formula of CaCl2. This means that the elements:
Calcium and Chlorine make up the compound in a 1:2 ratio.

Under the right conditions, for as long as there are Calcium and Chlorine in a
quantity for a ratio, Calcium Chloride will form in a definite proportion of 1:2.

b. Potassium Nitrate has the molecular formula of KNO3. The elemens Potassium,
Nitrogen and Oxygen make up the compound in 1:1:3 ratio.

3. Law of Multiple Proprtions. When atoms combine to form a compound they always
combine in definite ratio and proportion expressed in small whole numbers. These same

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elementsmay also combine in a different proportion to yield a different compound.

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Example:
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a. Nitrogen and Oxygen can combine in a variety of proportions:
NO Nitric oxide = 14g Nitrogen + 16g of OxygCarbon Men
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NO2 Nitrigen dioxide = 14g Nitrogen + 32g Oxygen

b. Tin and Oxygen can combine in different proportions to form different compounds:
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SnO Tin (II) oxide = 119g Tin + 16g Oxygen

SnO2 Tin (IV) oxide = 119g Tin + 32g Oxygen

c. Carbon and Oxygen can combine in a 1:1 ratio or 1:2 ratio:

CO Carbon monoxide = 12g Carbon + 16g Oxygen
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CO2 Carbon dioxide = 12g Carbon + 32g Oxygen

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Lesson 2. Atomic Structure

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Electrons were discovered by scientists whose main interests were electricity rather
than chemistry. They studied the flow of electric current through gases of low pressure
contained in glass tubes with metal disks called electrodes at each end. When connceted to
a source of high voltage, the tube glowed. It was observed that one electrode became
positively charged wgile the other anode became negatively charged. The glowing beam,
which travels from the cathode to the other anode is called the cathode ray.

In 1897 Sir Joseph J. Thomsom (1856-1940) showed in his experiments that a

cathode ray is a collection of very small negatively charged particles, which he named
electrons. He reasoned that the electrons must be part of the atom of all elements.

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Shortly thereafter, scientists began to think about particles left over when a hydrogen
atoms loses an electron. Since atoms are electrically neutral, researchers reasoned that the
leftover particle should have a posotive charge. Experimental evidence for such particles,
later named as protons, was soon
found out.

In 1913 Baron Ernest

Rutherford 91871-1937) decided to
test the prevailing theory of atomic
structure. The theory at that time
was that the protons and electrons
were evenly distributed throughout
the volume of an atom.

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To test this theory, they directed a
beam of particles at a very thin

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sheet of gold. The particles they
chose for this experiment were
alpha particles, which are positively
charged Helium atoms that lack two
electrons.

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His team observed that a us
small fraction of the alpha particles
bounced back. From these results,
Rutherford proposed that the mass
of the atom and the positive charge
are concentrated at a small region.
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He called this region the nucleus.

He thought of the rest of the atom as
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more or less empty space: that the

electrons were in that area but were
so small that they did not interfere
with the movement of the alpha
particles.
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In 1932, English physicist Sir

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James Chadwick (1891-1974)

confirmed the existence of yet
another subatomic particle: the neutron
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In the decades following the discovery of the neutron, scientists discovered more
subatomic particles. Physicists refer to two families of particles: leptons, the electron, the
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mu-meson, taumesons and neutrinos: the second family is called the hadrons that include
the proton and neutron and quarcks.

Most chemical reactions can be explained by the atomic structure advanced by John
Dalton an for purposes of this study and discussion, the structure of the atom shall consider
the three subatomic particles: proton, electron and neutron.

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1. Electron (e-) – are the negatively charged sub-

atomic particles of an atom. They are located outside
the nucleus and occupy electron orbital.
2. Protons (p+) – the positively charged sub-atomic
particles of an atom. They are located in the nucleus.
3. Neutrons (n0) – are the neutral sub-atomic
particles of an atom and are also located in the
nucleus.

(The nucleus at the middle consists of protons and

neutrons. In constant motion circling the nucleus are

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electrons spinning at designated orbitals. )

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Let us take the element Gold (Au). In the Periodic table, there are 2 numbers that are
assigned for Gold: 1 is located above the symbol Au and the other is located below.

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The number above the element’s symbol is the atomic number. The atomic number
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is the whole number that increases as you read across each row of the periodic table from
left to right. This also corresponds to the number of protons. Since the element is a neutral
entity, the positive charge of the protons must be cancelled out by a similar number of
electrically charged particles. This would mean that the number of protons should also be
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the number of electrons in an element’s atom.

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The number below the element’s symbol corresponds to the atomic mass expressed
as atomic mass unit (amu). Just how ingenious for scientists to have devised a means to
measure the mass of the infinitesimally small atom. They arbitrarily assigned a mass to one
atom and determined the masses of other atoms realtive to it. By international agreement,
the atomic mass standard is the pure isotope Carbon-12, which is assigned a mass of
exactly 12 atomic mass units (12u). based on this standard, an atomic mass unit (amu) is
exactly one-twelfth the mass of a carbon-12 atom.

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The term atomic mass replaced the older term atomic weight which is still used by
the International Union of Pure and Apllied Chemistry (IUPAC). These two terms are used
interchangeably.

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Since the mass of the atom is concentrated in the nucleus (which is made up of
protons and neutrons), given that protons must equal the number of electrons, the number of
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neutrons can thus be mathematically determined by subtracting the atomic mass to the
atomic number.
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Below is a tabulation of the first 10 elements from the atomic mass and atomic
number, the number of protons, neutrons and electrons are determined.

Name of Symbol Atomic Atomic Protons Electron Neutrons

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Element Mass (a) Number (b) (b) (a) – (b)

(b)
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Hydrogen H 1 1 1 1 0
Helium He 4 2 2 2 2
Lithium Li 7 3 3 3 4
Beryllium Be 9 4 4 4 5
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Boron B 11 5 5 5 6
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Carbon C 12 6 6 6 6
Nitrogen N 14 7 7 7 7
Oxygen O 16 8 8 8 8
Flourine F 19 9 9 9 10
Neon Ne 20 10 10 10 10

Lesson 3. Molecules, Ions, Isotopes and

Nuclides

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A. MOLECULES

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Molecules are the smallest particle in a chemical element or compound that exhibit
the chemical properties of that element or compound. Molecules are made up of atoms
that are held together by chemicsl bonds. Atoms may combine with atoms with the same
element to form molecules of an element. Atoms of one element may also combine with
atoms of another element to form molecules of a compound.

B. IONS

In their elemental state, elements have

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the same number of protons and electrons.
Thus, the positive charges cancel out the
negative charges and elements are neutral.
Sometimes atoms gain or lose electrons.

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When they do so they either gain a negatice
charge or a positive charge because the
number of electrons do not equal the number
of protons in the atom or molecule. When
atoms GAIN electrons, the result is that there

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are more electrons than protons. This would us
result to a net negative charge. When atoms
LOSE electrons, the result is that there are less electrons than protons and a net positive
charge of the ion. A positively charged ion is called a CATION while a negative charged
ion is called an ANION.
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Let us take a look at Sodium. Sodium has the chemical symbol of NA. Its atomic
number is 11 which means that Sodium has:
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11 protons 11 protons
11 electrons (lose 1 electron) 10 electrons
Net charge 0 Net charge +1
Na+ is called a CATION
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Chlorine on the other hand has the chemical symbol of Cl. Its atomic number is 17 which
means that Chlorine has :

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 LEARNING MODULE IN GENERAL CHEMISTRY (ORGANIC)

17 protons 17 protons
17 electrons (lose 1 electron) 18 electrons
Net charge 0 Net charge -1
Cl- is called an ANION

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ISOTOPESIsotopes have different atomic masses (mass number). Isotopes of an
element have nuclei with the same number of protons (the same atomic number) but
different numbers of neutrons. Isotopes have different mass numbers, which give the
total number of nucleons (protons

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+ neutrons).

Example, the most common

isotope of Hydrogen has no neutrons
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at all but there is also a Hydrogen
isotope called Deuterium which has
one neutron; and Tritium which has
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two neutrons.
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To symbolize the composition of an isotope, two numbers are written to the left of the
chemical symbol. The mass number is written as a superscript (above) and the atomic
number is written as a subscript (below).
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Example:
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Theree isotopes of Carbon are Carbon-12, Carbon-13. And carbon-14. Write the chemical
symbol for each.

Answer:
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Since carbon has an atomic number of 6, all Carbon atoms have 6 protons
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While discovered less than 100 years ago, isotopes are now used in a wide variety of
scientific applications that touch the lives of almost every citizen. These include:
Radiopharmaceuticals used for medical imaging and diagnosis of a wide range of ailments;
for cancer treatment and other therapeutic applications; for smoke detectors used in home
and offices; batteries that power NASA satelites in the far reaches of our solor system; for
control rods that prevent nuclear power reactors from melting down; to enable new sources
of energy such as nuclear fusion; and many other applications in energy production,
industrial diagnostic methods, archeology, geology, ecology, astronomy and physics.

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D. NUCLIDE

Nuclide is any particular atomic nucleus with a specific

atomic number Z and mass number A. it is equivalently an

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atomic nucleus with a specific number of protons and
neutrons. Collectively, all the isotopes of all the elements
form the set of nuclides. The terms isotope and nuclide are
often used interchangeably. Isotope is best used when
referring to several different nuclides of the same element;

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nuclide is more generic ansd is used when referring one to
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one nucleus or several nuclei of different elements. For
example, it is more correct to say that an element such as
Fluorine consists of one stable nuclide rsthehr than that it has
one stable isotope.
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Atomic number = the number of electrons and protons of the atom.

The atomic mass = atomic number + the number of neutrons.
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Therefore, the number of neutrons= (atomic mass-atomic number)

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SELF LEARNING ACTIVITY

MODULE III. Lesson 1

Direction: Solve the following problems. Apply the Dalton’s atomic laws.

1. CaCO₃, decomposes when heated to CaO and CO₂. if 100g of CaCO₃ can decompose, 56
g of CaO is produced. What mass of CO₂ is given off by the decomposition?

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2. Suppose 5.00g of Ca (OH)₂ and 10.00g of NH₄Cl are mixed in a test tube and heated until
no more ammonia is given off. The remaining material in the test tube has a mass of 1.27g.
what is the total mass of ammonia and water vapor produced?

3. A reaction of 17.0g of ammonia with 26.6 of chlorine gas gives off 3.5g of nitrogen. What
mass of ammonium chloride is formed in the reaction?

4. Mg (s) + ZnCl₂ (aq) MgCl₂ (aq) + Zn (s)

45.0g 297.6g 176.5g ___________

5. If 46.5 g of reactants are used in the following reaction, what will be the mass of the

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products?

Zn (s) + 2HCl (aq) ZnCl₂ (aq) + H₂(g)

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SELF LEARNING ACTIVITY

MODULE III. Lesson 2

A. Research Topics:

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1. There are subatomic particles other than the proton, e;ectron and neutron. Choose
ONE subatomic particle. Research on how it was discovered. What are its
properties?
2. Elements are also used by our bodies for various chemical processes.
A.What are these elements and what are their specific role in bodily processes?
B. What will result in the absence/ lack of these elements in or bodies?

Note: Use this color reference for making

your molecular model.
H B C N

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O F Si P

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S Cl Br I

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B. Complete the table below.

Symbol Atomic # Mass # # of # of # of Element

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Protons Electrons Neutrons
1. Fr 87 136
2. Rg 280 111
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3. Yb 70 103
4. Te 127.6 52
5. Md 101 157
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SELF LEARNING ACTIVITY

MODULE III. Lesson 3

A.

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1. What is an ion?
_________________________________________________________________
________________________________.
2. What does the number next to the ions signify?
____________________________________________________
3. Complete the table below using your Periodic table.

Important reminders: Name of Ion symbol Number of Number of Number of

element Protons Electrons Electrons
*In a neutral atom the
number of protons equals lost or
the number of electrons. gained
Ex. Flourine F- 9 (9+1=10) Gained 1
*An atom can NEVER gain
1. Hydrogen H+

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or lose protons.
2. Aluminum Al+3
*The number of protons
equals the atomic number.
3. Potassium K+
4. Sulfur S+2

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5. Iodine I-

B.

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1. Here are 3 isotopes of an element 6
12
C 6
13
C 6 14 C
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The element is : _____________
The number 6 refers to the ____________
The numbers 12, 13 and 14 refers to the _______________
How many protons and neutrons are in the 1st isotope?___________
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How many protons and neutrons are in the 2nd isotope?___________

How many protons and neutrons are in the 3rd isotope?___________
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2.Complete the table below using your Periodic table.

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Cr Chromium -58 Chromium -63
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# of Protons
# of Neutrons
# of Electrons
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