GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
Chemistry is a branch of science that focuses on the structure,
composition, and properties of matter.
Matter
MATTER
is anything that occupies space and has mass (i.e., anything that
has density).It commonly exists in three phases: solid, liquid, and
gas.
I.
THE THREE PHASES OF MATTER:
n
Solid – is composed of particles that are tightly packed and
have a regular arrangement. It has definite shape and
volume.
o
Liquid – is composed of particles whose arrangement and
packing are somewhere between those in solid and gas. It
has definite volume but no specific shape (it takes the
shape of its container).
p
Gas – is composed of particles with no regular arrangement
and no appreciable packing. It has no definite shape or
volume (it takes on the shape and volume of its container).
A fluid is either a liquid or a gas.
Plasma is a special state of a very hot chemical element whose atoms are
completely stripped of their electrons. Plasma is recognized as a
separate state because it has properties that distinguish it even from
gas.
II.
PHASE CHANGES
Sublimation of
an
element
or
compound is a transition
from the solid to gas
phase without passing the
liquid stage.
Deposition
is a process in which gas
transforms into solid (also
known as desublimation).
Condensation is the change in matter of a substance to a denser phase, such
as a gas (or vapor) to a liquid.
Evaporation is the conversion of water from a liquid into a gas.
-1-
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C
CHAPTER
1 - Chemistry
C
GEA
G
AS
GE
ENERAL ENGINEE
ERING & APPLIED
D SCIENCES
III. CLASSIFICATION
C
O MATTER: (SCH
OF
HEMATIC DIAGR
RAM)
All sa
amples of matte
er can be classsified as either mixtures or pure substances.
MAATTER
Pure Subsstance
Element
n
Homogeneeous
Heteroogeneous
Pure
P
Substance is any variety of matter
m
that is homogeneous an
nd with constant
composition
c
by mass. It can be
b classified intto two ‐ the ellements and thee
compounds.
c
An eleement
pounds
Comp
o
Compound
Mixtures
is the simplest fo
orm of
matterr since it contains only
one kin
nd of atom.
are pure
p
substance
es that
contain two or more
ents
che
emically
eleme
combined in a definite
mass.
proportion
by
pounds
can
n
be
Comp
classiffied in variouss ways.
One way is to classify
some of them into
o bases,
acids, and salts.
QUICK FACTS
9A
At present, the
ere are 115
k
known
elements. Of these, 88
o
occur
naturally on
o earth and
t rest are synth
the
hetic.
9T
The most abund
dant element
i the universe iss hydrogen, in
in
t entire earth is
the
i iron, and in
t
the
earth’s crusst, bodies of
w
water,
and atm
mosphere, its
o
oxygen.
Mixture
M
is comp
posed of two or more distin
nct substances, which can bee
separated
s
by physical means. It can be classifieed into two – ho
omogeneous and
d
heterogeneous.
h
mogeneous
A hom
A heteerogeneous
mixture also caalled solution has a uniforrm composition
m
n
th
hroughout. Its components
c
can
nnot be distingu
uished from onee
an
nother since thee whole mixture has only one ph
hase.
m
mixture
is one with two or more distinct phases.. Mixtures of this
tyype can be furrther classified to suspensions, colloids, and
d
coourse mixtures.
-2-
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
IV.
Chemistry
COMMON PROPERTIES OF PURE SUBSTANCES
of any pure substance is its ratio of mass to unit volume.
The density
Specific gravity
is a dimensionless ratio of the density of the substance to
the density of a standard--typically another substance,
usually a liquid, in which the substance in question is
suspended. The usual standard of specific gravity is water,
in which case the specific gravity is numerically equivalent
to the density. But one can also define the specific gravity
of a substance in, for example, an organic liquid, such as
benzene.
The melting point
of any pure substance is the temperature at which, under
common atmospheric pressure, that substance changes its
state from solid to liquid. If the substance is a liquid at 25
degrees Celsius, this temperature is usually called the
freezing point.
The boiling point
of any pure substance is the temperature at which, under
common atmospheric pressure, that substance changes its
state from liquid to gas. If the substance is a gas at 25
degrees Celsius, this temperature is sometimes called the
condensation point.
The triple point
of any pure substance is that combination of temperature
and pressure at which all three phases of that substance
coexist simultaneously.
The specific heat
of any pure substance is the amount of heat required to
raise the temperature of a unit mass of that substance by
one degree on a given temperature scale.
The heat of fusion
of any pure substance is the amount of heat required to
change a unit mass of that substance, once brought to the
melting point, from solid to liquid.
The heat of
Vaporization
The critical point
of any pure substance is the amount of heat required to
change a unit mass of that substance, once brought to the
boiling point, from liquid to gas.
of any pure substance is a point on a three-dimensional
graph of temperature, pressure, and molar volume (ratio of
volume to amount-of-substance) beyond which that
substance can exist only as a gas.
-3-
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
V.
PROPERTIES OF MATTER:
Properties of matter can be classified as physical or chemical and intensive or
extensive.
n
Physical Properties – are properties which can be measured without
changing the identity and composition of a substance.
Examples:
{odor, color, density, boiling point, melting point, polarity, solubility, opacity,
viscosity, etc….}
™
Physical changes are those changes that do not result in the production of
a new substance.
Examples of physical changes are:
{Melting, freezing, condensing, breaking, crushing, cutting, & bending}
o
Chemical Properties – are properties that lead to changes in the identity
and composition of a substance.
Examples:
{Combustibility, Reaction with water, pH, etc….}
™
Chemical changes or chemical reactions are changes that result in the
production of another substance.
Examples of chemical changes are:
{Digestion, respiration, photosynthesis, burning, rusting, decomposition,
etc...}
p
Intensive properties - are those which do not depend on the size of the
sample involved. Some of the most common intensive properties are;
density, freezing point, color, melting point, reactivity, luster, malleability,
and conductivity
q
Extensive properties - are those that do depend on the size of the sample
involved. A large sample of carbon would take up a bigger area than a
small sample of carbon, so volume is an extensive property. Some of
the most common types of extensive properties are; length, volume,
mass and weight.
-4-
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
THE ATOM
Atom is the basic building block of matter. It is the smallest particle of element.
An ELEMENT is a fundamental type of matter in which all of the atoms in the
material are the same.
I.
FUNDAMENTAL CHEMICAL LAWS
n
The Law of Conservation of Mass
“Mass is neither created nor destroyed”
o
The Law of Definite Proportion (originally called “Proust’s Law)
A given compound always contains exactly the same proportion of
elements by mass.
p
The Law of Multiple Proportions
When two elements from a series of compounds, the ratios of the
masses of the second element that combine with 1 gram of the first
element can always be reduced to small whole numbers.
II.
THE ATOMIC MODEL OF MATTER
The Dalton Model: (John Dalton, 1766-1844)
™
™
™
™
Each element is composed of indivisible particles called atoms.
Atoms of the same element are identical. Atoms of different
elements are different.
Compounds are formed when atoms of more than one element
combine. In any compound, the combination of atoms is in a fixed
ratio of small whole numbers.
In a chemical reaction, atoms are not created, destroyed, or
changed into other type of atoms.
The Thomson Model: (J.J. Thomson, 1856-1940)
™
An atom consists of a (positively charge) jellylike mass with
(negative) electrons scattered throughout it as far as possible.
The Rutherford Model: (Ernest Rutherford, 1871-1937)
™ Most of the (volume of the) atom is empty space.
™ Most of the mass of the atom is concentrated in a dense, positively
charged nucleus.
™ Electrons are present in the space surrounding the nucleus.
-5-
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
III. THE SUBATOMIC PARTICLES
Electrons
are the negatively charged particles
of an atom.
Protons
are the positively charged particles
of an atom
Neutrons
are the neutrally charged particles
of an atom.
CHARGE AND MASS OF SUBATOMIC PARTICLES
charge
+ 1.6002 x 10-19 C
Proton
0
Neutron
- 1.6002 x 10-19 C
electron
mass
1.6726 x 10-27 kg
1.6749 x 10-27 kg
9.1094 x 10-31 kg
location
nucleus
Nucleus
outside
IV. ATOMIC NUMBER AND MASS NUMBER
The atomic number of an element is the number of protons that is contained in
the nucleus of each of its atoms.
Mass number or atomic weight is the sum of the number of protons and
neutrons in the nucleus of the atom.
Formula:
Number of Neutrons = Mass number − Atomic Number
V. QUANTUM NUMBERS
Electrons within atoms are characterized by four quantum numbers:
The principal quantum number, n, determines the energy state of an electron.
It can have integer values of 1, 2, 3, up to n.
o The subshell number , l, defines the orbital shape. Its values start at (n-1),
and become smaller by integer values, ending at zero.
p An orbital number, m, which specifies the spatial orientation of an orbital. It
has integer values going from +1 through 0 to -1.
q Finally a spin quantum number, s, which can have values of +1/2 or -1/2, and
does not depend upon the values of n, l, or m. The electron within an atom
behaves as though it spins on its own axis.
Pauli Exclusion Principle – states that no two electrons in a atom can
have the same 4 quantum numbers. This means that each electron must have its
own unique set of 4 quantum numbers. Two electrons in an atom may have the
same values of n, l, and m, but the fourth quantum number, s, the spin number,
must be different.
n
-6-
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
ISOTOPES
Isotopes are atoms with the same atomic number (that is, same number of
protons) but different mass numbers (that is, different number of neutrons).
For example, chlorine occurs in nature in the form of two isotopes,
37
17
35
17
Cl and
Cl . The composition of these isotopes can be described as follows:
Isotope
35
17
Cl
37
17
Cl
No. of Protons
17
No. of Electrons
17
17
No. of Neutrons
18
17
20
Atoms of different atomic numbers but of the same mass number are called
isobars.
ISOTOPE DESIGNATION
The symbol for one particular type of chlorine atom is written as
Mass number
37
Atomic number
17
Sample Problem
Cl
Silver has two isotopes. One has 60 neutrons while the other has 62. The
atomic number of silver is 47.Write the symbols for these two isotopes of
silver.
Solution:
For the first isotope of silver:
mass number = no. of protons + no. of neutrons
= 47 + 60
= 107
" Symbol: 107
Ag
47
For the second isotope of silver:
mass number = no. of protons + no. of neutrons
= 47 + 62
" Symbol: 109
= 109
47 Ag
-7-
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
MOLECULES
A MOLECULE is the smallest identifiable sample of a substance.
All materials are made up of molecules and in turn are made up of atoms.
Monatomic molecules contain one atom
Diatomic molecules contain two atoms
Triatomic molecules contain three atoms
Polyatomic molecules describes any molecules that contain more than one
atom
Every molecule can be represented by a molecular formula, in which the
symbol of the element is succeeded by a subscript that indicates the number
of atoms present. The presence of one atom is indicated by writing the
symbol without the subscript 1.
ƒ
ƒ
ƒ
ƒ
Example: (Molecular Formulas of Some Common Molecules)
Diatomic Molecules
Polyatomic Molecules
Elements
Compounds
Elements
Compounds
H2, N2, O2, Fe2
CO, HCl
O3, P4, C60
H2O, CO2, CH4
™
NAMING MOLECULAR COMPOUNDS
Molecular compounds are named using Greek prefixes that indicate how
many atoms of an element are present in a molecule of a compound.
Examples:
CO
SO2
N2O3
H2O
- carbon monoxide
- sulfur dioxide
- dinitrogen trioxide
- dihydrogen monoxide
Note that the prefix “mono” is never used for the first element, as in
carbon monoxide rather than monocarbon monoxide and the ending –a
of the prefix is omitted when the next letter is a vowel, as in tetroxide
rather than tetraoxide.
Greek Prefixes Used for Naming Molecular Compounds
Prefix
Number of
Atoms
1
2
3
4
5
6
mono
di
tri
tetra
penta
hexa
-8-
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
IONS
An ion is an atom or molecule, which has lost or gained one or more
electrons, making it positively or negatively charged.
9 A positively charged ion is called cation.
9 A negatively charged ion is called anion.
9 A monatomic ion is an ion consisting of a single atom.
9 A radical ion is an ion that contains an odd number of electrons and are
mostly very reactive and unstable.
9 A dianion is an ion which has two negative charges on it.
9 A polyatomic ion is an electrically charged particle that consists of two or
more atoms linked together in much the same way as in neutral
molecule.
Oxyanions are polyatomic ions containing oxygen, such as carbonate and
sulfate.
™
NAMING IONIC COMPOUNDS
In naming ionic compound:
9 Combine the names of the two elements that make up the
compound.
9 The ion that has a positive charge will be place first followed by the
ion with the negative charge.
9 Add the suffix –ide at the end of the name
EXAMPLES: NaBr is sodium bromide
NaCl is sodium chloride
AlF3 is aluminum fluoride
SOME COMMON ELEMENTS WITH NEGATIVE OXIDATION NUMBERS AND THEIR NAME ENDING.
Symbol
Name ending
-
H
F
Br
I
Cl
2S
2O
3P
3N
4C
hydride
fluoride
bromide
iodide
chloride
sulfide
oxide
phospide
nitride
carbide
For ions with more than one positive charge (like many of the metallic elements),
write a Roman numeral in parentheses, which corresponds to its positive charge
For example, iron can exist as the Fe2+ or Fe3+ ion and, therefore can form two
oxides, namely FeO and Fe2O3. These compounds can be named as iron(II) oxide
and iron(III) oxide, respectively.
-9-
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C
CHAPTER
1 - Chemistry
C
GEA
G
AS
GE
ENERAL ENGINEE
ERING & APPLIED
D SCIENCES
IONS:
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GENE
ERAL ENGINE
EERING AND
D APPLIED SCIENCES
S
CHAPT
TER 1
iA
Chemisttry
THE PERI
IODIC
C TAB
BLE
iiA
iiiA ivA
iIiB ivB
vB
viB viiB
viiiB
iB
vAA
viiiA
viA viiA
iiB
T
Distinct Areas
A
in the Pe
eriodic Table are:
The Three
n
Main
M
Group Ele
ements
o
Transition
T
Grou
up Elements
p
Inner Transition
n Group Eleme
ents
o horizontal row of the periiodic table.
A perriod contains the elements in one
A gro
oup contains the
e elements in one
o column of the
t periodic tab
ble
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
CHEMICAL FORMULAS
A chemical formula is a concise way of expressing information about the atoms
that constitute a particular chemical compound. A chemical formula is also a
short way of showing how a chemical reaction occurs.
™
WRITING CHEMICAL FORMULA FOR IONIC COMPOUNDS: The Criss – Cross Method
General Representation:
A x +By − → A yBx
n
Write the symbols of the component elements A and B with their charges
(superscripts) x+ and y-, respectively. Write the positive ion first and the
negative ion last.
o
Crisscross the superscripts and write them as subscripts. Disregard the
signs of the charges when they become subscripts.
EXAMPLE:
3+
Write the compound formula when aluminum ion Al combines with sulfur ion
2S .
Solution:
Al3−S2 − → Al2S3
EXAMPLE:
2+
2Write the formula of the compound that contains Ca ions and O ions
Solution:
Ca 2 + O 2 − → CaO
Note: Always reduce the subscripts to the lowest possible ratio in the final formula.
EXAMPLE:
Write the chemical formula of the combination of barium ion with phosphate
ion.
Solution:
2+
For barium ion, the symbol is Ba
3For phosphate ion, the symbol is PO4
Ba 2 + (PO4 )
3−
→ Ba3 (PO4 )2
- 12 -
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
CHEMICAL BONDING
A chemical bond is a force which holds a group of atoms together so as to form
an electrically neutral aggregate.
n IONIC BOND
Ionic bond is formed by the electrostatic attraction of ions of opposite
charge formed by electron transfer. It involves a metal and a non-metal.
o
COVALENT BOND
A covalent bond is form of chemical bonding in which electrons are
shared between two atoms. It involves two non-metals.
is the amount of energy involved in the formation and
breaking of a bond.
pertains to single bond, double bond, triple bond and those
intermediate between single and double bonds, etc.
is the distance between the nuclei of the atoms forming the
bond.
CHEMICAL
CALCULATIONS
Bond Energy
Bond order
Bond Length
n ATOMIC MASS
The atomic mass (or atomic weight) of an element is the average of the
element’s isotopic masses.
CALCULATING AVERAGE MASS OF AN ELEMENT
Atomic mass = ( m1p1 + m2p 2 + m3p3 + ...)
Where
mn = mass of isotopes 1,2,3...
p n = percent abundance of isotopes 1,2,3...
Example
Calculate the average atomic mass of magnesium which has three
25
isotopes consisting of the following: 79% of 24
12 Mg , 10% of 12 Mg , and
11% of
26
12
Mg .
Solution:
average atomic mass = ( 0.79 )( 24 ) + ( 0.10 )( 25) + ( 0.11)( 26 )
= 24.32
Thus, the average mass of magnesium is 24.32 amu.
- 13 -
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
o FORMULA MASS
The formula mass (or formula weight) is the sum of the masses of all atoms in
a given formula.
CALCULATING formula mass
Find the formula mass of the compound CuSO4.
Solution:
Multiply the atomic mass of each element to the number of atoms of the
same element present in the given compound:
= 63.55 amu
Cu ;
63.55 amu × 1
= 32.07 amu
S ;
32.07 amu × 1
= 64.00 amu
O ;
16.00 amu × 4
Then, add the resulting atomic masses:
formula mass = 63.55 + 32.07 + 64.00
= 159.62 amu
p MOLE CONCEPT & MOLAR MASS
A mole is the amount of pure substance containing the same number of
chemical units, as there are atoms in exactly 12 grams of carbon-12.
EXAMPLE:
Mass of 1 atom of silver = 108 amu
Mass of 1 mole of silver = 108 g
Molar mass of silver
= 108 g/mol
AVOGADRO’S NUMBER
One mole refers to Avogadro’s number of particles of anything:
NA = 6.02 × 1023
q MOLE – MASS CONVERSIONS
The formula for calculating among mass, gram-formula mass (also known as
molar mass), and the number of moles:
m
Μ
n = number of moles
M = mass of the substance in grams
MM= molar mass in grams per mole
n=
Where:
EXAMPLE:
How many moles of molecules are contained in 67.25 grams of NH3. The
molar mass of NH3 is 17.03 grams/mole.
Solution:
67.25 g
n=
= 3.949 mol of NH3
17.03 g / mole
- 14 -
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
r
Chemistry
MOLE – NUMBERS OF PARTICLES CONVERSION
Conversion factor:
1 mole = 6.02 × 1023 particles
EXAMPLE:
How many particles are in 2.00 moles of SO2?
Solution:
⎛ 6.02 × 1023 molecules SO2 ⎞
N moles SO2 = 2.00 mol SO2 ⎜
⎟
1 mol SO2
⎝
⎠
24
= 1.20 × 10 molecules SO 2
s
EQUIVALENT WEIGHT
The equivalent is the amount of substance that supplies one gram-mole
(that is, 6.022x1023) of reacting units.
EW =
where
t
MW
Δoxidation number
EW = equivalent weight
MW = molecular weight
EMPIRICAL FORMULA
An empirical formula is a formula that gives the simplest whole-number
ratio of atoms in a compound.
Steps for Determining an Empirical Formula
n
Start with the number of grams of each element, given in the
problem.
If percentages are given, assume that the total mass is 100 grams so that
the mass of each element = the percent given.
o
Convert the mass of each element to moles using the molar mass
from the periodic table.
p
Divide each mole value by the smallest number of moles calculated.
- 15 -
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
q
Round to the nearest whole number. This is the mole ratio of the
elements and is represented by subscripts in the empirical formula.
™
If the number is too far to round (x.1 ~ x.9), then multiply each
solution by the same factor to get the lowest whole number
multiple.
e.g. If one solution is 1.5, then multiply each solution in the
problem by 2 to get 3.
e.g. If one solution is 1.25, then multiply each solution in the
problem by 4 to get 5.
EXAMPLE:
Chemical analysis of methyl ether showed it to be composed of
52.17%C,13.05%H, and 34.78%O. Determine its empirical formula.
Solution:
Converting percent to grams by assuming 100g of the compound, we
determine the number of moles in each mass using atomic masses of
each as unit conversion factors:
⎛ 1 molC ⎞
52.17 gC ⎜
⎟ = 4.344 mol C
⎝ 12.01 g C ⎠
⎛ 1 mol H ⎞
13.05 gH ⎜
⎟ = 12.9 mol H
⎝ 1.01 g H ⎠
⎛ 1 mol O ⎞
34.78 g O ⎜
⎟ = 2.174 mol O
⎝ 16.00 g O ⎠
Divide each mole value by the smallest number of moles calculated and
round off to the nearest whole number.
⎛ 1 molC ⎞
52.17 gC ⎜
⎟=
⎝ 12.01 g C ⎠
⎛ 1 mol H ⎞
13.05 gH ⎜
⎟=
⎝ 1.01 g H ⎠
4.344
mol C → 1.998 ≈o
2.174
12.9
mol H → 5.93 ≈s
2.174
⎛ 1 mol O ⎞ 2.174
34.78 g O ⎜
mol O →n
⎟=
⎝ 16.00 g O ⎠ 2.174
Thus, the empirical formula of methyl ether is C2H6O
- 16 -
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
MOLECULAR FORMULA
Once the empirical formula is found, the molecular formula for a
compound can be determined if the molar mass of the compound is
known.
Steps for Determining molecular Formula
u
1.
2.
3.
4.
Find the empirical formula
Find the mass of the empirical unit.
Divide the molecular mass of the compound by the mass of the empirical formula.
Multiply all the atoms (subscripts) of the empirical formula by this ratio to find the
molecular formula.
SAMPLE PROBLEM:
By chemical analysis, a compound was found to be composed of 75.46% carbon,
4.43% hydrogen, and 20.10% oxygen. Its molecular mass was found to be
approximately 318 g/mol. What is the molecular formula for this compound?
Solution:
Find the empirical formula.
Get the mass of each element by assuming a certain overall mass for the sample
(100 g is a good mass to assume when working with percentages).
(0.7546) (100 g) = 75.46 g C
(0.0443) (100 g) = 4.43 g H
(0.2010) (100 g) = 20.10 g O
Convert the mass of each element to moles.
(75.46 g C) (1 mol/ 12.00 g C) = 6.289 mol C
(4.43 g H) (1 mol/ 1.008 g H) = 4.39 mol H
(20.10 g O) (1 mol/ 16.00 g O) = 1.256 mol O
Find the ratio of the moles of each element.
(1.256 mol O)/ (1.256) = 1 mol O
(6.289 mol C)/ (1.256) = 5.007 mol C
(4.39 mol H)/ (1.256) = 3.50 mol H
Use the mole ratio to write the empirical formula.
Since the ratio 3.5 is too far to round off, multiply the mole ratios by two to get
whole number. The empirical formula becomes:
C10H7O2
Find the mass of the empirical unit.
10(12.00) + 7(1.008) + 2(16.00) = 159.06 g/mol
Divide the molecular mass of the compound by the mass of the empirical formula.
(318 g/mol) / (159.06 g/mol) =1.999 ≈ 2 empirical units per molecular unit
Multiply all the atoms (subscripts) of the empirical formula by this ratio and write
the molecular formula. C20H14O4
- 17 -
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
CHEMICAL EQUATIONS
A chemical equation is a shorthand representation of a chemical reaction using
chemical formulas indicating the reactants and products.
Reactant refers to the original material in the reaction.
Products refer to the new substances formed because of the reaction.
™
SYMBOLS USED IN CHEMICAL EQUATIONS
Symbols used as shorthand information in a chemical reaction:
9 ( ↑ ) - is used to indicate evolution of a gas
9
( ↓ ) - is used to indicate formation of a precipitate
9
⎛ Δ⎞
⎜ → ⎟ - is used to indicate the application of heat in the reaction
⎝ ⎠
Symbols for the physical states of the substance:
9 (s) - solid
9 (L) - liquid
9 (g) - gas
9 (aq) - for a substance dissolved in water
™
WRITING CHEMICAL EQUATIONS
In writing chemical equation
9 Chemical formulas of the compounds are used instead of their
names
9 Formulas of the reactants are written on the left side of the
equation while the products are on the right side.
9 The (+) sign is used in place of the word and.
9 In between the reactants and products, an arrow ( →) is used to
mean form, produce, or yield.
Example:
Write a chemical equation for the following chemical reaction:
Methane (CH4) gas reacts with oxygen gas to produce carbon dioxide
gas, liquid water, and heat.
Solution:
CH4 (g) + O2 (g) → CO2 (g) + H2O(L) + heat
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
CLASSIFICATION OF
CHEMICAL REACTIONS
A chemical reaction is a process in which a substance or a combination of
substances undergo a change in appearance or properties, and further
transform into a different substance or a combination of new substances
n
DIRECT COMBINATION OR SYNTHESIS
A direct combination or synthesis reaction
involves the combination of two or more
reactants to form one product. The reactants can
be elements or compounds.
ACTIVITY SERIES
METALS
NONMETALS
M
General Equation:
M
A + B → AB
o
L
DECOMPOSITION
A decomposition reaction involves the
breakdown of a single reactant into two or more
products.
General Equation:
M- most
L - least
AB → A + B
p
SINGLE – REPLACEMENT REACTIONS
In a single – replacement reaction, an
uncombined element replaces another element
that is part of a compound. As a result, the
replaced element becomes uncombined.
General Equation:
A + BC → AC + B
q
Double – Replacement Reactions
In a double-replacement reaction, two elements
in different compounds replace each other.
General Equation:
AB + CD → AD + CB
L
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
SOLUTIONS
I.
UNITS OF CONCENTRATION
n
MOLE FRACTION
The number of moles of solute divided by the number of moles of solvent
and all solutes.
Suppose a solution contains two components, A and B. The mole fraction
(X) of components A or B is denoted as follows:
XA =
nA
nA + nB
or
XB =
nB
nA + nB
Where:
N = number of moles of each component present
o
NORMALITY
The number of gram equivalent weights of solute per liter. A solution is “
normal” if there is exactly one gram equivalent weight per liter.
nORMALITY =
p
Equivalent weight in grams
Vsolution in liters
MOLARITY
Molarity (M) is defined as the number of moles of solute dissolved in 1
liter of solution. In other words, molarity is a ratio between number of
moles of solute and the number of liters of solution.
M =
nsolute
Vsolutio ( L )
Where:
M = molarity in molar
n = number of moles of solute
V = volume of solution in liters
q
FORMALITY
The number of gram formula weights (i.e., molecular weights in grams)
per liter of solution.
FORMALITY =
Formula weight in grams
Vsolution in liters
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
r
Chemistry
MOLALITY
Molality (m) is defined as the number of moles of solute dissolved in 1 kg
of solvent. In other words molality is the ratio between the number of
moles of solute and the mass of the solvent expressed in kilograms.
mass solute
nsolute
MM solute
=
m=
kg solvent
kg solvent
Where:
m = molality in molal
n = number of moles of solute
MM= molar mass of solute
s
PERCENT OF VOLUME
Percent of volume refers to the number of millilitres of solute dissolved in
100 ml of solution.
% volume =
t
volume solute
x100
volume solution
DILUTION
Dilution is the process of adding solvent (usually water) to a
concentrated solution to achieve a solution of the desired concentration.
When we dilute a solution, we do not change the number of moles of
solute present, we simply add more solvent. Thus,
Moles of solute after dilution=Moles of solute before dilution
nafter = nbefore
( MV )after = ( MV )before
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
ACIDS
Acid is any compound that dissociates in water into H+ ions.
Acids with 1,2, and 3 ionizable hydrogen atoms are called monoprotic,
diprotic, and triprotic acids, respectively.
PROPERTIES OF ACIDS:
1.
2.
3.
4.
5.
6.
7.
Acid conducts electricity in aqueous solutions
Acids have a sour taste
Acids turn blue litmus paper to red
Acids have pH between 0 and 7
Acids neutralizes bases
Acids react with active metals to form hydrogen
Acids react with oxides and hydroxides to form salts and water
pH Equation:
⎛ 1 ⎞
pH = log ⎜ + ⎟
⎜⎡ ⎤⎟
⎝ ⎣H ⎦ ⎠
Where:
⎡⎣ H + ⎤⎦ ⎫⎪
⎬ = ionic concentration in moles of ions per liter
⎡⎣OH − ⎤⎦ ⎪
⎭
For a partially ionized compound, X, in a solution of known molarity, M, the
ionic concentration is:
[ X] = ( fraction ionized ) × M
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
BASES
Base is any compound that dissociates in water into OH- ions.
Bases with 1, 2, and 3 replaceable hydroxide ions are called monohydroxic,
dihydroxic, and trihydroxic bases, respectively.
Properties of Basis:
1. Bases conduct electricity in aqueous solutions
2. Bases have bitter taste
3. Bases turn red litmus paper to blue
4. Bases have pH between 7 and 14
5. Bases neutralize acids, forming salts and water
pOH Equation:
⎛ 1 ⎞
⎟
pOH = log ⎜
⎜ ⎡ −⎤ ⎟
⎝ ⎣OH ⎦ ⎠
pH and pOH Relationship:
pH + pOH = 14
Neutralization:
Acids and Bases neutralize each other to form water.
H+ + OH− → H2O
A neutral solution has a pH of 7.
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
TEST - 1
1.
An instrument that separates particles of different isotopic composition
and measure their individual relative masses.
A. mass spectrometer
B. barometer
C. hygrometer
D. mass spectometer
2.
These are compounds containing water molecules loosely bound to the
other components.
A. isotope
B. hydrates
C. ion
D. mixture
3.
If a more active element replaces a less active one in a compound, the
reaction is:
A. combustion reactions
B. replacement reactions
C. metathesis
D. neutralization
4.
If a single reactant is transformed by heat or electricity into two or more
products, the type of reaction is
A.
B.
C.
D.
5.
decomposition
combination
displacement
double displacement
The numerical value for standard pressure of any gas is
A. 76 mm Hg
B. 760 cm Hg
C. 760 mm Hg
D. 7.6 cm Hg
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
6.
Chemistry
Any process extracting the metal in a fused state is called
A. Calcination
B. Roasting
C. Smelting
D. Froth flotation process
7.
Which of the following does not change with change in temperature?
A.
B.
C.
D.
8.
A device used to measure density.
A.
B.
C.
D.
9.
volume
mass
pressure
density
manometer
hydrometer
spectrometer
densimeter
The statement “mass is neither created nor destroyed in a chemical
reaction” is known as:
A.
B.
C.
D.
The law of conservation of mass
The law of constant composition
The law of multiple proportions
The law of chemical reaction
10. What kind of chemical bond will form in binary compounds where the
electronegativity difference between atoms is greater than 2.0
A.
B.
C.
D.
Ionic Bond
Covalent bond
Metallic bond
Chemical bond
11. What kind of chemical bond will form in binary compounds where the
electronegativity difference between atoms is less than 1.5?
A.
B.
C.
D.
Covalent bond
Ionic bond
Metallic bond
Chemical bond
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
12. The element which has a mass of about 9 amu is beryllium (Be), atomic
number 4. What is the charge on the Be atom?
A. +4
B. +8
C. -4
D. neutral
13. The horizontal arrangement of elements of increasing atomic number in
a periodic table is called:
A. period
B. group
C. family
D. row
14. Refers to atoms or ions, which have the same electronic configuration.
A.
B.
C.
D.
isoelectronic
isotope
isotomic
isometric
15. Which group of the periodic table is known as the alkali metals?
A. Group I
B. Group IV
C. Group III
D. Group VII
16. Which group is known as the halogens?
A. Group V
B. Group II
C. Group VII
D. Group III
17. The mass that enters into a chemical reaction remains the unchanged as
a result of the reaction. In precise form: mass is neither created nor
destroyed. This is known as:
A.
B.
C.
D.
the law of conservation of mass
the law of definite proportion
the law of multiple proportion
law of conservation of energy
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
18. When the same elements can form two different compounds, the ratio of
masses of one of the elements in the two compounds is a small whole
number relative to a given mass of the other element. This is known as:
A.
B.
C.
D.
the law of constant composition.
the law of conservation of mass
the law of multiple proportion
law of conservation of energy
19. The ratio of the density of the test liquid to the density of a reference
liquid is called:
A.
B.
C.
D.
specific gravity
relative gravity
specific weight
relative weight
20. How many electrons are there in a covalent bond?
A.
B.
C.
D.
3
2
4
8
21. The SI unit of temperature is
A.
B.
C.
D.
Fahreheit
Kelvin
Celsius
Rankine
22. The elements that a compound is composed of are present in fixed and
precise proportion by mass. This is known as:
A.
B.
C.
D.
the law of constant composition
the law of conservation of mass
the law of multiple proportion
law of conservation of energy
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
23. The mass to the nearest atomic-mass unit of an atom which contains 9
protons and 10 neutrons.
A.
B.
C.
D.
10 amu
19 amu
15 amu
21 amu
24. The number of protons in the nucleus of an atom is called
A.
B.
C.
D.
atomic number
percent abundance
atomic weight
oxidation number
25. The measure of the resistance of an object to a changed in its state of
motion is called
A. momentum
B. mass
C. inertia
D. velocity
26. Refers to the agreement of a particular value with the true value.
A. precision
B. error
C. tolerance
D. accuracy
27. Refers to the degree of agreement among several measurements of the
same quantity.
A. accuracy
B. precision
C. error
D. margin
28. A property of matter that is often used by chemist as an “identification
tag” for a substance.
A.
B.
C.
D.
mass
molarity
density
volume
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
29. Protons and neutrons can be broken down further into elementary
particles called
A. quarks
B. ions
C. isotope
D. warks
30. The principle of the constant composition of compounds, originally called
“Proust’s Law” is now known as
A.
B.
C.
D.
The law of multiple proportion
The law of definite proportion
The law of conservation of mass
The law of compounds
31. The mass of an alpha α particle is how many times more than that of the
electron?
A.
B.
C.
D.
1837 times
7300 times
1829 times
1567 times
32. Atoms with the same number of protons but different number of neutrons
are called
A. ions
B. quarks
C. isotopes
D. compounds
33. The forces that hold atoms together are called
A. mechanical bond
B. formula bond
C. atomic bond
D. chemical bond
34. An atom or group of atoms that has a net positive or negative charge is
called
A. ion
B. isotope
C. positron
D. polymer
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
35. A positive ion is called
A.
B.
C.
D.
positron
anion
cation
quark
36. A negative ion is called
A. positron
B. anion
C. cation
D. quark
37. The force of attraction between oppositely charged ions is called
A.
B.
C.
D.
ionic bonding
covalent bonding
polar bonding
metallic bonding
38. The horizontal arrangement of elements of increasing atomic number in
a periodic table is called
A. group
B. period
C. series
D. row
39. The vertical arrangement of elements in the periodic table is called
A.
B.
C.
D.
period
group
series
column
40. If the number of gas molecule is doubled in a certain volume of gas the
pressure
A.
B.
C.
D.
is decreased to half
is doubled
is increased to four times
remains unchanged
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
41. A symbolic representation to emphasize the valence shell of an atom is
called
A. argand diagram
B. canstellation diagram
C. electron dot diagram
D. structural formula
42. For which elements do the number of electrons in their outer or valence
shell correspond to their group number?
A.
B.
C.
D.
transition group
noble gas
representative or main group
metals
43. What is the maximum number of electrons that can fit into a “p” orbital?
A. 2
B. 4
C. 6
D. 8
44. Atoms or ions which have the same electronic configuration.
A. isoelectronic
B. isometric
C. iso-ionic
D. isotope
45. Which group of the periodic table is known as the alkali metals?
A.
B.
C.
D.
Group 1
Group 2
Group 3
Group 4
46. Which group of the periodic table is known as the halogens?
A.
B.
C.
D.
Group 4
Group 5
Group 6
Group 7
47. Which group of the periodic table is known as the alkaline earths ?
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CHAPTER 1 - Chemistry
GEAS
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A.
B.
C.
D.
Group 1
Group 2
Group 3
Group 4
48. Which group of the periodic table is known as the noble gases?
A.
B.
C.
D.
Group 5
Group 6
Group 7
Group 8
49. The components resulting from the reactions are called
A. products
B. reductants
C. reactants
D. oxidants
50. What kind of chemical bond will form in binary compounds where the
electronegativity difference between atoms is less than 1.5?
A. Covalent Bond
B. Ionic Bond
C. Super bond
D. Electrovalent bond
51. Compounds which contain only carbon and hydrogen are called
A. polymorphs
B. hydrocarbons
C. polycarbon
D. plastics
52. Which of the following designation means that the amount of solute is
expressed in physical mass units, i.e., grams, and the amount of solution
(not solvent) is expressed in volume units, i.e., milliliters.
A. v/v
B. w/v
C. w/w
D. v/w
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
53. What kind of chemical bond will form in binary compounds where the
electronegativity difference between atoms is greater than 2.0
A. Covalent Bond
B. Ionic Bond
C. Super bond
D. Electrovalent bond
54. The property of liquid which describes their resistance to flow is called
A.
B.
C.
D.
viscosity
elasticity
glueyness
stickiness
55. The anions and cations which are unaffected by the reaction in solution
are called
A.
B.
C.
D.
neutral ions
spectator ions
noble ions
observer ions
56. A reaction in which heat is produced is called
A. exothermic
B. isothermic
C. endothermic
D. pyrothermic
57. A reaction in which heat is absorbed is called
A. exothermic
B. isothermic
C. endothermic
D. pyrothermic
58. In oxidation-reduction or redox reactions, the component supplying the
electrons is called the
A.
B.
C.
D.
reductant
reducing agent
oxidant
acceptor
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
59. In oxidation-reduction or redox reactions, the component receiving the
electrons is called
A. reductant
B. reducing agent
C. oxidant
D. acceptor
60. A covalent bond formed by the combination of collinear p orbitals is also
called a
A. alpha bond
B. sigma bond
C. delta bond
D. gamma bond
61. The tendency of atoms to attract electrons into their valence shells to
form anions is described by the concept of
A. electronegativity
B. electron mobility
C. electron affinity
D. electron ability
62. The tendency of an atom to attract electrons shared in a covalent bond is
called
A. electronegativity
B. electron mobility
C. electron affinity
D. electron ability
63. A covalent bond between atoms of identical electronegativity is called
A.
B.
C.
D.
polar
bipolar
nonpolar
monopolar
64. A formula which describes only the numbers of each element in the
molecule is called
A.
B.
C.
D.
structural formula
molecular formula
empirical formula
ionic formula
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
65. The formula that describes how atoms are joined together is called the
A.
B.
C.
D.
structural formula
molecular formula
empirical formula
ionic formula
66. The pairs of electrons not shared in the covalent bond are called
A.
B.
C.
D.
bonded electrons
free electrons
valence electrons
nonbonded electrons or lone pairs
67. The word atom comes from the Greek word, atomos meaning
A.
B.
C.
D.
unique
cannot be cut
single
cannot be destroyed
68. The central part of an atom is called the
A. nucleus
B. core
C. hub
D. heart
69. The mass of a proton is about how many times the mass of an electron?
A. 1639
B. 1837
C. 1387
D. 1587
70. The particles in the nucleus, namely the neutrons and the protons, are
collectively referred to as
A. positrons
B. electrods
C. nucleons
D. isotope
71. The number of orbiting electrons is normally ________the number of
protons in the nucleus of an atom.
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
A.
B.
C.
D.
more than
less than
equal
half
72. The net electric charge on any atom is
A. zero
B. +1
C. -1
D. dependent on the number of protons in the nucleus
73. Under normal state, an atom is said to be
A.
B.
C.
D.
positively charged
negatively charged
electrically neural
positively or negatively charged
74. The word electron comes from the Greek word “elektron” which means
A. cannot be cut
B. amber
C. unique
D. negative
75. The word proton comes from the Greek word “proteios” meaning
A. of first importance
B. with positive charge
C. unique
D. cannot be cut
76. A solid which has no crystalline structure is called
A. Non-crystalline
B. Amorphous
C. Fused
D. Immiscible
77. A chemical substance which readily evaporates and readily diffuses at
ordinary room temperature and pressure conditions is called
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Chemistry
A. Inflammable
B. Inert
C. Volatile
D. Corrosive
78. Which of the following falls under fluids?
A.
B.
C.
D.
Solid
Gas
Liquid
Both liquid and gas
79. The gases that rarely take part in a chemical reaction are called
A.
B.
C.
D.
Miscible gases
Volatile gases
Noble gases
Permanent gases
80. Which type of ions, metals form when enter into a chemical reaction?
A.
B.
C.
D.
Negative ions
Positive ions
Either positive or negative ions
They do not form any ions
81. The chemical name for baking soda is
A.
B.
C.
D.
Sodium bicarbonate
Sodium sulphate
Sodium chloride
Sodium carbonate
82. The subatomic particle with a negative charge and mass of 9.1 X 10
is
A. Proton
B. Neutron
C. Electron
D. Positron
-31
kg
83. The subatomic particle with a positive charge and mass of 9.1 X 10-27 kg
is
A. Proton
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
B. Neutron
C. Electron
D. Positron
84. Neutron was discovered by
A. J.J. Thompson
B. Chadwick
C. Rutherford
D. Einstein
85. Electron was discovered by
A.
B.
C.
D.
J.J. Thompson
Chadwick
Bohr
Einstein
86. The isotopes of an element differ in the number of
A. Electrons
B. Neutrons
C. Protons and neutrons
D.
Both
87. The isotope of hydrogen is
A. Protium
C. Deuterium
B. Tritium
D. All of the above
88. The isotope of hydrogen with only one neutron is called
A.
B.
C.
D.
Protium
Deuterium
Tritium
monotium
89. The isotope of hydrogen with two neutrons is called
A. Protium
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
B. Deuterium
C. Tritium
D. politium
90. The atoms having different atomic numbers but the same mass number
are called
A.
B.
C.
D.
Isotones
Isotopes
Homologues
Isobars
91. The atoms which have the same number of neutrons but different mass
numbers are called
A. Isotones
B. Isotopes
C. Homologues
D. Isobars
92. Rutherford model of an atom had failed to explain
A. the location of electrons in an atom
B. the position of protons and neutrons
C. the distribution of electrons around the nucleus
D. both (A) and (B)
93. The concept that electrons revolve around the nucleus in specific paths
called orbits or energy levels was proposed by
A. Rutherford
B. Niels Bohr
C. J.J. Thompson
D. Chadwick
94. The number of atoms (6.023x1023) present in 12 grams of carbon-12
is called
A. Avogadro’s constant
B. Planck’s constant
C. Reinhold’s constant
D. Bohr’s constant
95. The chemical formula which shows the relative number of atoms
the elements present in a compound is called
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of all
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
A. Molecular formula
B. Empirical formula
C. Structural formula
D. Compound formula
96. The chemical formula which shows the exact numbers of atoms
the elements present in a compound is called
of all
A. Molecular formula
B. Empirical formula
C. Structural formula
D. Compound formula
97. When an atom loses an electron, it forms
A. Cation
B. Either cation or anion
C. Anion
D. Neither cation nor anion
98. When an electron gains an electron, it forms
A. Cation
B. Anion
C. Either cation or anion
D. Neither cation nor anion
99. The maximum number of electrons, the first energy level can
accommodate in an atom is
A. Two
B. Sixteen
C. Eight
D. Thirty two
100. The X-rays were discovered by
A.
B.
C.
D.
Madam Curie
Pierre Curie
Henry Becquerel
W.C. Roentgen
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GENE
ERAL ENGINE
EERING AND
D APPLIED SCIENCES
S
CHAPT
TER 1
Chemisttry
So
olve
ed Pr
roble
ems
Inn
1.
Chhem
mistr
try
The
T
solubility of
o the sucrose
e, a chemical name
n
for suga
ar, is 490/100g
g
water
w
at 100°C
C. The solution is prepared byy mixing 175g sugar in 0.045
5
kg
k of water at 100°C.
1
Determ
mine the maxim
mum amount of sugar that can
n
be
b dissolved in the water at 100°C.
Solution:
S
⎞ ⎛ 490g sugarr ⎞
⎛
1000g
mass of su
ugar = ⎜ 0.045 kkg ×
H 2O ⎟ ⎜⎜
⎟ ⎝ 100gH 2O ⎟⎟⎠
⎜
1 kg
⎠
⎝
mass of su
ugar = 220.5g sugar
2.
Calculate
C
the mass of a tita
anium atom. Titanium
T
eleme
ent has atomicc
mass
m
of 47.9 amu.
Solution:
S
1 atom Ti ×
Ti
1mol T
6.022 × 10
23
atomsTi
×
47.9g Ti
1mol Ti
2
7.95 × 10−23
g Ti
3.
Determine
D
the number of molecules in 20g of
o C9H8O4.
Solution:
S
molar masss C9H8O 4 = 9 (12g) + 8 (1g) + 4 (16g)
molar masss C9H8O 4 = 1880gC9H8O4
20 gC9H8O 4 ×
1 molC9H8O4
180 gC9H8O 4
×
6.022 × 1023 molecules1molC
m
9H8O 4
1 molC9H8O4
6.69 × 1022 molecules C9H8O4
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
4.
What is the percentage mass of oxygen in Fe2O3? (Fe = 55.9amu, O =
16amu.)
Solution:
1mol Fe2O3 = 3mol O
= 2mol Fe
%O =
3 (16gO )
2 ( 55.9gFe ) + 3 (16gO )
× 100%
%O = 30%
5.
What is the atomic mass of oxygen which consists of three isotopes with
atomic masses 16amu, 17amu, and 18 amu, with abundances 99.76%,
0.04%, 02% respectively.
Solution:
Atomic Mass:
( 0.9976 )(16amu ) + ( 0.0004 )(17amu ) + ( 0.002 )(18amu )
16amu
6.
How many moles are there in 55g of CF2Cl2, a chlorofluorocarbon that
damages the ozone layer in the atmosphere? (C = 12amu, F = 19amu,
Cl = 35.45 amu).
Solution:
molar mass CF2Cl2 ⇒ 12g + 2 (19g) + 2 ( 35.45g)
molar mass CF2Cl2 ⇒ 120.9g = 1mol
# moles CF2Cl2 :
55g CF2Cl2 ×
1mol CF2Cl2
120.9gCF2Cl2
0.45 mol CF2Cl2
8.
A container has 83mL of nitric acid solution which is labelled 7.2M HNO3.
Determine the number of moles of HNO3 in the container?
Solution:
# mols of solute
Molarity =
L solution
mols HNO3
7.2M =
83 × 10−3 L
mols HNO3 = 0.6mol HNO3
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
9.
Chemistry
Calculate the mass in grams of sulphur that can be obtained from 25 g of
C6H10O2S. C=12amu, H=1amu, O=16amu, S=32amu.
Solution:
molar mass C6H10O2S ⇒ 6 (12g) + 10 (1g) + 2 (16g) + 32g
molar mass C6H10O2S ⇒ 146g = 1mol
mass(g) S :
25g C6H10O2S ×
1mol C6H10O2S
1molS
32gS
×
×
146gC6H10O2S 1mol C6H10O2S 1molS
5.5mol S
10. Calculate the volume of NH3 solution which has 0.14 M, containing 10g
of NH3.
Solution:
# mols of solute
Molarity =
L solution
1mol NH3
10g NH3 ×
14g + 3 (1g)
0.14M =
L solution
L solution = 4.2 L
11. Arterial blood contains about 0.25mg of oxygen per millilitre. Determine
the pressure exerted by the oxygen in one liter of arterial blood at normal
body temperature of 38°C?
Solution:
ρ = 0.25
mg
g
= 0.25
mL
L
nRT
V
m
m
g
= 0.25
n=
;ρ =
MM
V
L
ρRT
mRT
=
P=
(MM)( V ) MM
P=
g ⎞⎛
L − atm ⎞
⎛
⎜ 0.25 L ⎟ ⎜ 0.0821 mol − K ⎟ ( 37 + 273 )
⎝
⎠
⎝
⎠
= 0.2 atm
P=
( 2 × 16g)
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
11. How many nitrogen atoms are there in 200g of NH4N3?
Solution:
molar mass NH4N3 ⇒ 4 (14g) + 4 (1g )
molar mass NH4N3 ⇒ 60g = 1mol
# Nitrogen atoms :
200g NH4N3 ×
1mol NH4N3
4mol N
6.022 × 1023 atoms N
×
×
60gNH4N3 1mol NH4N3
1molN
8.03 × 1024 atoms N
12. What is the mass in grams of C2H6O2 in 0.769 mol of C2H6O2?
Solution:
molar mass C2H6O2 ⇒ 2 (12g) + 6 (1g ) + 2 (16g)
molar mass C2H6O2 ⇒ 62g = 1mol
mass (g) C2H6O2 :
0.769 mol C2H6O2 ×
62g C2H6O2
1mol C2H6O2
47.7 g C2H6O2
13. Determine the molarity of 2.6 L of a 3M solution after it has been diluted
to 5.9 L.
Solution:
# mols solute
L solution
# mols solute = ( 3M)( 2.6L )
Molarity =
# mols solute=7.8 mols solute
After dilution:
7.8 mols solute
5.9 L
Molarity = 1.3 L
Molarity =
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
14. What is the molar fraction of solvent in an aqueous solution of alcohol,
with 5 moles of alcohol and 10 moles of water?
Solution:
# mols solvent
# mols solute + # mols solvent
# mols H2O
=
# mols alcohol + # mols H2O
Xsolvent =
XH2O
10
5 + 10
= 0.67
XH2O =
XH2O
15. Determine the molality of NH3 in aqueous solution if the mole fraction
NH3 is 0.343.
Solution:
XNH3 =
# mols NH3
# mols NH3 + # mols H2O
# mols NH3
0.343
=
1
# mols NH3 + # mols H2O
Let :
# mols NH3 = 0.343mols NH3
then,
# mols NH3 + # mols H2O = 1
# mols H2O=1 − # mols NH3 = 1 − 0.343
# mols H2O=0.657 mols
molality =
# mols NH3
kg H2O
0.343 mols NH3
molality =
0.657 mols H2O ×
( 2 (1) + 16 ) gH2O × 1kgH2O
1 mol H2O
1000 g
molality = 29 m
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
16. Calculate the volume in liter occupied by 0.92 mol of H2 at STP.
Solution:
@ STP :
V = 0.92 mol H2 ×
22.4L
1mol N2
V = 20.6 L
17. What is the volume occupied by 0.252 mol of nitrogen gas at 100°C and
85.6 kPa.
Solution:
@ STP :
V1 = 0.252 mol N2 ×
T1 = 273 K ;
22.4L
= 5.6448 L
1mol N2 1
P1 = 101.325 kPa
Solving for V2 @ T2 = 100°C and P2 = 85.6 kPa :
P2 V2 P1V1
=
T2
T1
V2 =
V2 =
P1V1T2
P2T1
(101.325 kPa )( 5.6448 L )(100 + 273 ) K
( 85.6 kPa )( 273K )
= 9.1 L
18. Determine the molar mass of 5.21g of gas which occupies 3.92 L at 105
kPa and 25°C.
Solution:
PV
n=
RT
(105 kPa )( 3.92L )
n=
L − kPa ⎞
⎛
⎜ 8.314 mol − K ⎟ ( 25 + 273 ) K
⎝
⎠
n = 0.166 mol
5.21g
MM=
0.166 mol
MM=31.36 g/mol
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
19. Calculate the molar mass of a gas with the density of 2.73 g/L and
pressure of 1.5 atm at 22°C.
Solution:
Assume 1L-volume: mass = 2.73g
PV
n=
RT
(1.5atm )(1L )
n=
L − atm ⎞
⎛
⎜ 0.0821 mol − K ⎟ ( 22 + 273 ) K
⎝
⎠
n = 0.062 mol
mass
MM=
n
2.73g
MM=
0.062 mol
MM=44 g/mol
20. What is the molinity of 550-g solution having 25 moles of solute?
Solution:
# moles of solute
molinity =
kg of solution
25 mols
molinity =
0.550kg
molinity = 45.45 M
21. How many neutrons in the nucleus of Fluorine element, with atomic
mass of 19amu and atomic number of 9?
Solution:
Atomic mass = # neutrons + # protons
Atomic number = # protons = 9
# neutrons = Atomic mass − # protons
# neutrons = 19 − 9
# neutrons = 10
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
22. Determine the pH value of a solution with 3.2 x 10-8 hydrogen ion
concentration.
Solution:
( )
pH = − log ( 3.2 × 10 )
pH = − log H+
−8
pH = 7.5
23. What is the volume of 5.0 kg of Hg with the density of 13.6 g/mL?
Solution:
m
ρ=
V
m
V=
ρ
5000g
g
13.6
mL
V = 368mL
V=
24. What is the mass of oxygen that reacts with 7.89 g Aluminum to produce
14.78 g of Aluminum Oxide?
Solution:
mass Oxygen = 14.78g − 7.89g
mass Oxygen = 6.89g
25. If 7.35 g of sulphur reacts with 4.92g of aluminum to form the only
compound of sulphur and aluminum, how much sulphur will react with
9.12 g of aluminum?
Solution:
mass Sulfur 7.35gS
=
9.12gAl
4.92gAl
mass Sulfur = 13.62g
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GENERAL ENGINEERING AND APPLIED SCIENCES
CHAPTER 1
Chemistry
26. Calculate the molar mass of the compound N2H6(NO3)2.
Solution:
molar mass N 2H6 ( NO3 )2 :
= 4 (14amu ) + 6 (1amu ) + 6 (16amu )
= 158amu
27. Determine the concentration in molarity of a 320-mL solution which
contains 2.5 moles of solute.
Solution:
2.5mols solute
0.320 L
M = 7.8 M
M=
28. How many moles of solute required to prepare 4.5m aqueous solution
containing 267g of H2O?
Solution:
molality =
# moles solute
kg solvent
# moles solute = ( 4.5m )( 0.267kg)
# moles solute = 1.2 mols
29. What is the mole fraction of NaCl in a solution containing of 0.032 moles
of NaCl in 75g of water?
Solution:
XNaCl =
# mols NaCl
# mols NaCl + # mols H2O
0.032 mols NaCl
⎛
1mol ⎞
0.032 mols NaCl + ⎜ 75g ×
⎟ HO
⎜
2 (1) + 16 ⎟⎠ 2
⎝
= 0.00762
XNaCl =
XNaCl
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CHAPTER 1 - Chemistry
GEAS
GENERAL ENGINEERING & APPLIED SCIENCES
30. What is the hydrogen ion concentration of a solution with pH 5.5.
( )
5.5 = − log ( H )
pH = − log H+
+
H+ = 3.16 × 10 −6 M
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