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Chemistry Online Textbook
Chapter 1: Atomic Structure…………………....2
Chapter 2: Isotopes……………………...……….7
Chapter 3: The Periodic Table….......................13
Chapter 4: Bonding – Ionic and Covalent…….20
Chapter 5: Chemical Reactions……………..…24
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Chapter 1: Atomic Structure
Inside the Atom
In the early 1900s, scientists began to identify the particles that make up atoms
(subatomic particles). Below is a picture of the basic parts of the atom.
Every atom has a core called a nucleus, where the majority (99.9%, to be exact) of
an atom's mass is held. Although the nucleus contains the majority of the mass of
the atom, the nucleus is very small compared to the size of the whole atom,
because most of the atom is empty space surrounding the nucleus. The nucleus is
made up of two smaller particles called protons and neutrons. The third type of
particles that make up the atom, electrons, orbit around the nucleus.
Protons
• Protons are positively charged particles which form part of the nucleus of an
atom.
• Every atom of a particular element contains the same number of protons. In fact,
the number of protons is unique to each element. Each element has a unique
atomic number, or a unique number of protons in its nucleus. Proton number
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never changes for any given element. For example, oxygen has an atomic number
of 8. That tells us that oxygen always has 8 protons.
Neutrons
• Neutrons are the other particle that make up the nucleus of an atom.
• Neutrons have about the same mass as protons.
• Unlike protons and electrons, however, neutrons carry no electrical charge.
Therefore, neutrons are "neutral."
• Atoms of a given element do not always contain the same number of neutrons.
Atoms of an element that have a different number of neutrons in the nucleus are
called isotopes of that element.
Electrons
• Electrons are negatively charged particles that orbit around the outside of the
nucleus.
• The mass of an electron is about 1/2000th of the mass of a proton or a neutron.
• The sharing or exchange of electrons between atoms forms chemical bonds,
which is how new molecules and compounds are formed.
Atomic Number
• An atom's atomic number tells you how many protons are in that atom's
nucleus. For example, oxygen has an atomic number of 8, meaning that there are 8
protons in the nucleus of an atom of oxygen. Copper's atomic number is 29,
meaning that there are 29 protons in the nucleus of an atom of copper. Later, you'll
see how the periodic table conveniently tells you each element's atomic number.
Atomic Mass
• Because atoms are so small, their masses cannot be measured in grams or
milligrams. Instead, scientists have created the atomic mass unit (amu) to
measure mass of subatomic particles.
• The mass of a proton or a neutron is about 1 amu. The mass of an electron,
however, is about 1/2000 amu.
• To find the atomic mass of an atom, add the number of protons and
neutrons in the nucleus.
• Example: If an atom has 3 protons, 4 neutrons, and 3 electrons, the atomic mass
is 7 amu, because you do not count the very small mass of the atom's electrons
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(1/2000 amu). Only add the number of protons and neutrons (each has a mass of 1
amu) in the nucleus.
Think Inside the Box
On the periodic table, each box represents a different
element, and each box contains vital information about the
element, including its name, symbol, atomic number, and
atomic mass. Different periodic tables may put the numbers
in different places, but the example box and instructions
below should help you find your way around.
Atomic Number
The whole number (no decimals) is the atomic number. Every element has its
own unique atomic number. The atomic number tells how many protons are in one
atom of that element. Since no two elements have the same atomic number, no two
elements have the same number of protons.
Symbol
The large letter is the element's symbol, and usually just below that is the element's
name. Each element has its own unique symbol and name. It is often very useful to
memorize symbols and names for elements, especially the more commonly used
elements.
Atomic Mass
The number with the extra decimal places is the element's atomic mass. The
atomic mass is the mass in atomic mass units for all possible isotopes of that
element. The atomic mass essentially gives you an estimate of how massive one
atom of that element is.
Ions
We've talked about ions before. Now it's time to get down to basics. Ions are
atoms with either extra electrons or missing electrons. An atom with no extra or
missing electrons is called a “neutral atom”. In a neutral atom, the number of
electrons is the same as the number of protons (the atomic number).
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So, why would an atom ever have extra or missing electrons? The electrons in
atoms are arranged in layers or shells. Atoms like to have complete layers, so if
they’re missing an electron or two in their outer layer, they may grab extra
electrons from some other atom. Or if they have an extra electron or two in their
outer layer, they may give those electrons to another atom trying to complete its
outer layer. Each layer or shell can only hold a certain number of electrons. The
first shell can only hold 2 electrons, the second holds 8, the third holds 8, the fourth
holds 18, the fifth holds 18 and the 6th holds 32 and the 7th can hold 32.
What do you do if you are a sodium (Na) atom? You
have eleven electrons. Two are in the first shell, 8 are
in the second and one extra on the outside shell. The
outer shell is called the valence shell. This outer shell
wants to have a full set of electrons. You need to find
another element who will take that electron away from
you. Bring in chlorine (Cl) which is only missing one
electron from its outer shell (or valence shell).
Chlorine will take that electron away and leave
sodium with 10 electrons inside of two filled shells.
Sodium is now a happy atom and has become a
sodium ion (Na+) with one less electron than its
atomic number. Since it now has more protons than electrons, sodium is a positive
ion.
Chlorine is also a “happy” atom since it needed one more electron to fill its outer
layer. However, this does give chlorine one more electron than its atomic number,
so like sodium, chlorine has also become an ion (Cl -) except chlorine is a negative
ion since it has more electrons than protons.
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Chapter 1 Questions: Atomic Structure
Inside the Atom
1. What are the three basic particles which make up an atom?
2. Which subatomic particles make up the nucleus of the atom?
3. Which subatomic particle has a negative charge?
4. Which subatomic particle has much less mass than the others?
5. Think about it:
a. The mass of an atom is mostly due to which subatomic particle(s)?
b. If you change the number of protons in an atom, what happens?
c. If you change the number of electrons in an atom, what happens?
Think Inside the Box
6. What does the large letter (or sometimes 2 or 3 letters) in each box of the
periodic table mean or stand for?
7. What are the two numbers found in each box of the periodic table?
8. How can you find the number of protons for an element on the periodic table?
9. How do you calculate the mass of an atom?
Ions
10. Describe what an ion is.
11. What is a valence electron shell?
12. What is the “happy atom” idea about?
13. Which subatomic particle is involved with ions?
14. Copy the following table onto your paper and fill it in completely.
Element
Symbol
# protons
# electrons
# neutrons
Magnesium
Sulfur
Argon
Oxygen ion
(-2 charge)
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Chapter 2:
Isotopes
Introduction
If you ever tasted ice cream or had a popsicle you know that you can have different
flavors. However even though ice cream can be vanilla, chocolate, or strawberry
and popsicles have flavors like cherry or grape, all flavors of ice cream are still ice
cream and all flavors of popsicles are popsicles. This is how we answer what are
isotopes. Isotopes are different flavors of the same atomic element.
Isotopes are basically the same element with a different
number of neutrons. The nucleus of every atom is made up of
neutrons and protons. While neutrons don’t have a charge they
do important work such as helping to bind the positive charged
protons together via the strong force. Isotopes can have
different numbers of neutrons but the basic elemental structure
remains the same. For example deuterium will have one more
neutron than normal hydrogen but it will still be considered
hydrogen. The only thing that will change is the atomic
weight.
Atoms are composed of a cloud of electrons surrounding a dense nucleus that is
100,000 times smaller. The nucleus is comprised of protons and neutrons, each
with mass roughly 2000 times that of the electron. The number of protons -- the
''atomic number'' -- determines the element; for example, a Strontium nucleus
always has 38 protons, and a Rubidium nucleus always has 37. There is an equal
number of electrons surrounding the nucleus to keep the atom electrically neutral,
and these electrons determine the chemical properties of the element -- such as,
forming the bonds that enable molecules like Strontium Chloride SrCl2 to form
from individual Strontium and Chlorine atoms.
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The different isotopes of an element have
different number of neutrons; for
example, while most (82.58 %) of
Strontium atoms have 50 neutrons, some
have only 46, 48, or 49. For only certain
numbers of neutrons, atoms are ''stable''
and do not radioactively decay into
another element. For Strontium, the stable
isotopes have atomic mass 84 [38 protons
and 46 neutrons], 86 [38 protons and 48
neutrons], 87, and 88. For other numbers
of neutrons, the atoms are unstable and radioactively decay. For example,
Strontium-82 is radioactive and decays to Rubidium-82 with a half-life of 25 days.
Chemical and molecular properties
A neutral atom has the same number of electrons as protons. Thus, different
isotopes of a given element all have the same number of protons and share a
similar electronic structure. Because the chemical behavior of an atom is largely
determined by its electronic structure, different isotopes exhibit nearly identical
chemical behavior.
Numbers of isotopes per element
Of the 80 elements with a stable isotope, the largest number of stable isotopes
observed for any element is ten (for the element tin). Xenon is the only element
that has nine stable isotopes. No element has eight stable isotopes. Four elements
have seven stable isotopes, nine have six stable isotopes, nine have five stable
isotopes, nine have four stable isotopes, five have three stable isotopes, 16 have
two stable isotope and 26 elements have only a single stable isotope. In total, there
are 255 isotopes that have not been observed to decay. For the 80 elements that
have one or more stable isotopes, the average number of stable isotopes is 255/80 =
3.2 isotopes per element.
As discussed above, only 80 elements have any stable isotopes, and 26 of these
have only one stable isotope. Thus, about two thirds of stable elements occur
naturally on Earth in multiple stable isotopes, with the largest number of stable
isotopes for an element being ten, for tin. There are about 94 elements found
naturally on Earth, though some are detected only in very tiny amounts. Scientists
estimate that the elements that occur naturally on Earth occur as 339 isotopes in
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total. Only 255 of these naturally occurring isotopes are stable in the sense of never
having been observed to decay as of the present time.
All the known stable isotopes occur naturally on Earth; the other naturally
occurring-isotopes are radioactive but occur on Earth due to their relatively long
half-lives, or else due to other means of ongoing natural production.
Calculating the Mass of Elements
If you look at a periodic table you may notice that the atomic mass of an element is
rarely a whole number. That’s because of the isotopes. Atomic masses are
calculated by figuring out how many atoms of each isotope exist in the universe.
There are a lot of carbon-12, a couple carbon-13, and a few carbon-14 atoms.
Averaging the masses, you get a number that is a little bit higher than 12 (the
atomic mass for carbon is 12.011). Not knowing which carbon isotope(s) a given
sample might contain, you should always use the average (atomic) mass for
carbon.
Importance of Isotopes
Isotopes have a special set of properties called nuclear properties. For example,
certain isotopes are more prone to radioactive decay making them ideal materials
to use in nuclear reactions. The most famous element with isotopes is uranium.
Uranium is one of the key fuels used in nuclear weapons and reactors. However,
only a handful of isotopes of this element can be used.
Scientists also use some radioactive isotopes as a
way to find out the date of ancient objects. For this
process, the isotope Carbon 14 is used. This is a
special isotope of carbon normally found in living
things or items made from organic materials
(materials that contain Carbon). Carbon dating uses
the concept of half-life. A half-life is the time it
takes for half the amount of a radioactive isotope to
decay to half its original amount. Since this time is
fixed, it can be used to calculate the age of organic
material (material that contains carbon).
Another nuclear property of isotopes involves stable isotopes. These are isotopes
that are normally not prone to detectable radioactive decay. It just means that these
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isotopes seem to have half-lives that are longer than the age of the earth, 4.6 billion
years. At the moment, there is no way to find out exactly what their true
radioactive properties are.
Isotopes are also used in industry and medicine. Certain isotopes of elements may
be good for manufacturing and others are being used for advanced imaging of
internal organs. Elements with very short half-lives are normally used since they
produce a very low level of radiation in a reasonable amount of time.
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Chapter 2 Questions: Isotopes
Introduction
1. How are isotopes like flavors of ice cream?
2. How does a deuterium, which is an isotope of hydrogen differ from a regular
hydrogen atom?
3. How does Strontium-82 differ from Strontium-84?
Chemical and molecular properties
4. How do the chemical and molecular properties if differing isotopes of the
same element compare? Why?
Numbers of isotopes per element
5. What element has the largest number of stable isotopes?
6. How many elements have only one stable isotope?
7. What is the average number of stable isotopes per element in the periodic
table?
8. What does it mean when an isotope is considered “stable?”
Calculating the Mass of Elements
9. Why is the average atomic mass of Carbon calculated to be 12.011?
Importance of Isotopes
10. What element is used to fuel nuclear reactors?
11.What is Carbon 14 used for? Why is Carbon 14 used for this purpose and not
another isotope?
12.What is a half life?
13.Why is it impossible to determine the half-lives of stable isotopes?
14.What are two additional uses for isotopes? What kind of isotopes do these
uses utilize? Why?
15. Fill in the chart below
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Element
Symbol
# protons
# electrons
# neutrons
Gold
Neon
Calcium
Chlorine ion
Potassium ion
Strontium 90
Carbon - 14
Helium - 6
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Chapter 3: The Periodic Table
How to Read the Periodic Table
The Periodic table is designed to help you predict what an element's physical and
chemical properties are. You can also predict what elements will bond with each
other.
First, let's look at the columns and rows of the periodic table.
Groups or Families
The vertical columns of the periodic table (there are 18) are called groups or
families. Elements in the same group or family have similar but not identical
characteristics. You will learn more about the 18 groups in a later section. You can
know properties of a certain element by knowing which group it belongs to.
Periods
The horizontal rows of the periodic table are called periods. Elements in a period
are not alike in properties. As a rule, the first element in a period is usually an
active solid, and the last element in a period is always an inactive gas. Atomic size
decreases from left to right across a period, but atomic mass increases from left to
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right across a period. Atoms on the left of the period, therefore, are usually larger
and more lightweight than the smaller, heavier atoms on the right of the period.
Metals, Nonmetals & Metalloids
Most periodic tables contain a stair step line which allows you to identify which
elements are metals, nonmetals, and metalloids. Following are descriptions of each
of the three types of materials.
Metals
Most elements are metals. 88 elements to the left of the stair step line are metals or
metal like elements.
Physical Properties of Metals:
• Luster (shininess)
• Good conductors of heat and electricity
• High density (heavy for their size)
• High melting point
• Ductile (most metals can be drawn out into thin wires)
• Malleable (most metals can be hammered into thin sheets)
Chemical Properties of Metals:
• Easily lose electrons
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• Corrode easily. Corrosion is a gradual wearing away. (Example: silver
tarnishing and iron rusting)
Nonmetals
Nonmetals are found to the right of the stair step line. Their characteristics are
opposite those of metals.
Physical Properties of Nonmetals:
• No luster (dull appearance)
• Poor conductor of heat and electricity
• Brittle (breaks easily)
• Not ductile
• Not malleable
• Low density
• Low melting point
Chemical Properties of Nonmetals:
• Tend to gain electrons
Since metals tend to lose electrons and nonmetals tend to gain electrons, metals
and nonmetals like to form compounds with each other. These compounds are
called ionic compounds. When two or more nonmetals bond with each other, they
form a covalent compound.
Metalloids
Elements on both sides of the zigzag line have properties of both metals and
nonmetals. These elements are called metalloids.
Physical Properties of Metalloids:
• Solids
• Can be shiny or dull
• Ductile
• Malleable
• Conduct heat and electricity better than nonmetals but not as well as metals
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Trends of the Periodic Table
Note: These are general periodic trends of the elements. There are many
exceptions to these general rules.
Atomic Radius - Atomic radius is simply the radius of the atom, an indication of
the atom's volume.
Period - atomic radius decreases as you go from left to right across a period.
Why? Stronger attractive forces in atoms (as you go from left to right)
between the opposite charges in the nucleus and electron cloud cause
the atom to be 'sucked' together a little tighter.
Group - atomic radius increases as you go down a group.
Why? There is a significant jump in the size of the nucleus (protons
+ neutrons) each time you move from period to period down a
group. Additionally, new energy levels of electron clouds are
added to the atom as you move from period to period down a
group, making the each atom significantly more massive, both is
mass and volume.
Electronegativity - Electronegativity is an atom's 'desire' to grab another atom's
electrons.
Period - electronegativity increases as you go from left to right across a
period.
Why? Elements on the left of the period table have 1 -2 valence
electrons and would rather give those few valence electrons away
(to achieve the octet in a lower energy level) than grab another
atom's electrons. As a result, they have low electronegativity.
Elements on the right side of the period table only need a few
electrons to complete the octet, so they have strong desire to grab
another atom's electrons.
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Group - electronegativity decreases as you go down a group.
Why? Elements near the top of the period table have few electrons
to begin with; every electron is a big deal. They have a stronger
desire to acquire more electrons. Elements near the bottom of the
chart have so many electrons that loosing or acquiring an electron
is not as big a deal. This is due to the shielding affect where
electrons in lower energy levels shield the positive charge of the
nucleus from outer electrons resulting in those outer electrons not
being as tightly bound to the atom.
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Chapter 3 Questions: The Periodic Table
How to Read the Periodic Table
1.
2.
3.
4.
5.
What are the vertical columns of the Periodic Table called?
What are the horizontal rows of the Periodic Table called?
How many groups/families are there on the Periodic Table?
How many periods are there on the Periodic Table?
Describe a trend in the Periodic Table as you go across the Periodic Table
from left to right.
6. Name which family each of the following elements can be found in:
a. Calcium
b. Chlorine
c. Lithium
d. neon
Metals, Nonmetals & Metalloids
7. What is the most common type of element on the Periodic Table? (metals,
nonmetals, metalloids)
8. What are the physical properties of metals?
9. What are the physical properties of nonmetals?
10.What kind of element would best make a coffee cup that would not burn
your hand?
11.What kind of element conducts electricity the best?
12.What is the difference between an ionic and a covalent bond?
13.Describe the physical properties of metalloids.
14.List two other elements that have properties similar to potassium?
15.What is the lightest element on the periodic table with properties similar to
bromine?
16.For each of the following elements, list if the element is a metal, nonmetal or
metalloid:
a. Oxygen
b. Calcium
c. Sulfur
d. Silicon
17.List two examples of a nonmetal.
18.Name three of the properties of metals
19.List any element(s) which are metalloids found in the fourth period.
20.List any element(s) which are in the third period, and nonmetal.
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Trends in the Periodic Table
21.What is electronegativity?
22.Describe the trend on the periodic table (both going across and down) for
electronegativity.
23.Describe the trend on the periodic table for atomic radius.
24.Arrange the following three elements in order from most electronegative to
least
electronegative: copper, nitrogen, molybdenum.
25.Arrange the following three elements in order from smallest to largest:
francium,
aluminum, zirconium
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Chapter 4: Bonding – Ionic and Covalent
Ionic Bonds
You must first learn why atoms bond together. We use a concept called "Happy
Atoms." We figure most atoms want to be happy, just like you. The idea behind
Happy Atoms is that atomic shells like to be full. That's it. If you are an atom and
you have a shell, you want your shell to be full. Some atoms have too many
electrons (one or two extra). These atoms like to give up their electrons. Some
atoms are really close to having a full shell. Those atoms go around looking for
other atoms who want to give up an electron.
Let's take a look at some examples.
We should start with the atoms with atomic numbers between 1 and 18. There is a
2-8-8 rule for these elements. The first shell is filled with 2 electrons, the second is
filled with 8 electrons, and the third is filled with 8. You can see that sodium (Na)
and magnesium (Mg) have a couple of extra electrons. They, like all atoms, want
to be happy. They have two possibilities: (1) They can try to get eight electrons to
fill up their third shell. Or (2) they give up a few electrons and have a filled second
shell. For them it’s easier to give up a few electrons.
What a coincidence! Many other atoms are interested in gaining a few extra
electrons.
Oxygen (O) and fluorine (F) are two good examples. Each of those elements is
looking for a couple of electrons to make a filled shell. They have one filled shell
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with two electrons but their second shell wants to have eight. There are a couple of
ways they can get the electrons. (1) They can share electrons, making a covalent
bond. Or (2) they can just borrow them, and make an ionic bond.
So we've got a sodium (Na) atom that has an extra electron. We've also got a
fluorine (F) atom that is looking for one.
They wind up working together and both wind up happy! Sodium (Na) gives up its
extra electron. The sodium (Na) has a full second shell and the fluorine (F) has a
full second shell. Two happy atoms! That's one way things are able to bond
together. They can give up or share electrons. The two elements have created an
electrovalent bond.
Covalent Bonds
A covalent compound is a compound in which the atoms that are bonded share
electrons rather than transfer electrons from one to the other. While ionic
compounds are usually formed when metals bond to nonmetals, covalent
compounds are formed when two nonmetals bond to each other.
The big question that students frequently have is, "Why do elements share
electrons? After all, wouldn't atoms rather grab electrons outright? That's what
happens when ionic compounds are formed."
The reason that nonmetals have to share electrons with each other has to do with
electronegativity. Recall that electronegativity is a measure of how much an
element pulls electrons away from other elements it is bonded to. Metals generally
have very low electronegativities (they don't much want to grab electrons) while
nonmetals have high electronegativities (they really want to grab electrons). The
reason for this trend is the octet rule, which says that all elements want to have the
same number of electrons as the nearest noble gas, because noble gases are
unusually stable. When metals bond to nonmetals, ionic compounds are formed
because the metal atoms don't want electrons and easily give them to nonmetals
that do want electrons.
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It's a different story when two nonmetals bond with each other. Instead of having
one element give electrons to another, we run into a case where we have two
elements that have roughly the same electronegativity. As a result, neither element
can steal electrons from the other. As a result, if either of them are going to be like
the nearest noble gas, they'll have to share electrons rather than transfer them. Keep
in mind, also, even if one non-metal could steal electrons from the other nonmetal, this would only allow one of the two non-metals to complete its shell of
electrons. The other would still have an incorrect number of electrons for a
complete shell.
Is it possible to predict whether bonds are covalent or not? A good rule of thumb is
that bonds between non-metals (remember that hydrogen is considered a nonmetal) are usually covalent bonds. For example, the carbon dioxide (CO2)
molecules you exhale are bonded together covalently.
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Chapter 4 Questions Bonding - Ionic and
Covalent
Ionic Bonds
1. What does it mean when we say an atom is “happy”?
2. What does an atom want to do if it’s missing just one electron (it has seven
electrons in its outer shell)?
3. What does an atom want to do if it has an “extra” electron (it has just one
electron in its outer shell)?
4. What is the atom called after it gives up or grabs an extra electron (hint: this is
why it’s called an ionic bond)?
5. When one atom gives up an electron and the other atom takes an electron, why
are the two ions attracted to each other?
Covalent Bonds
6. What do atoms in a covalent bond do rather than transferring electrons?
7. What two types of atoms tend to bond together in a covalent bond (metal/metal,
metal/nonmetal, nonmetal/nonmetal)?
8. What two types of atoms tend to bond together in an ionic bond?
9. Why would an ionic bond not work for two nonmetals? (Use an example like
two oxygen atoms and explain what would happen)
10. Which family of elements tends to not form any bonds (ionic or covalent)? Why
doesn’t this family of elements form bonds?
11. Which subatomic particle is involved in the bonding of atoms to form
molecules?
12. Briefly describe the difference between covalent and ionic bonds.
13. Which type of bond would each of the following pairs of atoms form: ionic or
covalent:
a. C, O
b. Ca, O
c. N, N
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Chapter 5: Chemical Reactions
To react or not to react? THAT is the question!
Chemical changes are a result of chemical reactions. All chemical reactions
involve a change in substances and a change in energy. Neither matter nor energy
is created or destroyed in a chemical reaction–only changed. There are so many
chemical reactions that it is helpful to classify them into 4 general types which
include the following:
SYNTHESIS REACTION
In a synthesis reaction two or more simple substances combine to form a more
complex substance. Two or more reactants yielding one product is another way to
identify a synthesis reaction.
For example, simple hydrogen gas combined with simple oxygen gas can produce
a more complex substance: water!
The chemical equation for this synthesis reaction looks like:
To visualize a synthesis reaction look at the following cartoon:
In the cartoon, the skinny bird (reactant) and the worm (reactant) combine to make
one product, a fat bird.
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DECOMPOSITION REACTION
In a decomposition reaction a more complex substance breaks down into its more
simple parts. One reactant yields 2 or more products. Basically, synthesis and
decomposition reactions are opposites.
For example, water can be broken down into hydrogen gas and oxygen gas.
The chemical equation for this decomposition reaction looks like:
To visualize a decomposition reaction look at the following cartoon:
In this cartoon the egg (the reactant), which contained the turtle at one time, now
has opened and the turtle (product) and egg shell (product) are now two separate
substances.
SINGLE DISPLACEMENT REACTION
In a single displacement reaction a single uncombined element replaces another in
a compound. Two reactants yield two products. For example when zinc combines
with hydrochloric acid, the zinc replaces hydrogen.
The chemical equation for this single displacement reaction looks like:
25
To visualize a single displacement reaction look at the following cartoon:
Notice, the guy in the grey shirt steals the date of the other guy. So, a part of one of
the reactants trades places and is in a different place among the products.
DOUBLE DISPLACEMENT REACTION
In a double displacement reaction parts of two compounds switch places to form
two new compounds. Two reactants yield two products. For example when silver
nitrate combines with sodium chloride, two new compounds--silver chloride and
sodium nitrate are formed because the sodium and silver switched places.
The chemical equation for this double displacement reaction looks like:
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To visualize a double displacement reaction look at the following cartoon:
In this cartoon, the two fellows wearing the hats trade hats with each other. So, in
the chemical reaction, a part of one reactant trades places with a part of another
reactant. In other words, one part of each reactant got switched with each other.
ENERGY OF CHEMICAL REACTIONS
Chemical reactions always involve a change in energy. Remember: in a chemical
reaction, atoms are rearranged from their original molecules to form new
molecules. To do this, it is necessary to break the bonds between atoms in the old
molecules and form new bonds between atoms to make new molecules. Breaking
bonds takes energy. When new bonds form, energy is given off.
If more energy is needed for breaking the old bonds than is given back when the
new bonds form, the chemical reaction will absorb or take in heat or energy. This
is called an endothermic reaction.
When the new bonds formed give back more energy than what was required to
break the old bonds, the chemical reaction will give off or release energy, often as
heat. This is called an exothermic reaction. The heat given off often causes the
product(s) to feel hot. Any reaction that involves combustion (burning) is an
exothermic chemical reaction.
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Chapter 5 Questions – Chemical Reactions
To React or Not To React
1. List four main types of chemical reactions.
2. Briefly describe what happens in each of those reaction types.
3. Using a Venn diagram, compare and contrast a synthesis and decomposition
reaction.
4. Using a Venn diagram, compare and contrast a single displacement and double
displacement reaction.
5. What is the main difference between an endothermic reaction and an exothermic
reaction?
6. Indicate the following reactions as endothermic or exothermic:
a. burning wood in a fireplace
b. baking a cake
c. boiling water (evaporation of water)
d. cooking an egg
e. making an ice cube
7. Write the following reaction down on your paper and balance them. Also,
indicate what type of reaction they are:
a. ____ Fe + ____ H2O  ____ Fe3O4 + ____ H2
Type of Reaction:
b. ____ P + ____ O2  ____ P4O10
Type of Reaction:
c. ____ KClO3 ____ KCl + ____ O2
Type of Reaction:
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