Chapter 1-Atomic Concept and Mole
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Transcript of Chapter 1-Atomic Concept and Mole
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Atomic Concept and Mole
1. Matter
2. Atoms and Ions
3. Atomic and Molecular Mass
4. Mole Concept
5. Names of Chemical Compounds
6. Chemical Formula
7. Chemical Equation and Stoichiometry8. Volumetric Analysis
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Matter
Matter can be described as
anything that has mass and it
must take up space.
Examples
States of Matter
Classification of Matter
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States of Matter
Depending on its temperature,matter can be solid, liquid or
gas.
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Classification of Matter
MixturesPure Substances
Homogeneous Heterogeneous
Elements Compounds
IonsAtoms Molecules
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Atom
An atom is the smallest particledifferentiable as a certain
chemical element.
When an atom of an element is
divided, it ceases to be thatelement.
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Atoms
An atom consists of three sub-atomic
particles which are electrons,
protons and neutrons.
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The atomic number ( Z ) is the number of protons in thenucleus of each atom of an element ( Z = p ).
In a neutral atom the number of protons is equal to thenumber of electrons ( p = e ).
The mass number ( A ) is the total number of neutrons andprotons present in the nucleus of an atom of an element.In general the mass number is given by:
Mass number = number of protons + number of neutrons
= atomic number + number of neutrons
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The accepted way to denote the atomic
number and mass of an atom of an
element ( X ) is as
Mass number or
AXNucleon number
ZAtomic number orProton number
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Atom Example
(a) Sodium (Na),(b) Iron (Fe),
(c) Gold (Au),
(d) Oxygen (O),
(e) Hydrogen (H),
(f) Chlorine (Cl),
(g) Carbon (C).10
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Molecules
A molecule is the smallest indivisible
portion of a pure compound that
retains a set of unique chemical
properties.
A molecule consists of two or more
atoms bonded together.
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Molecules - examples
(a) Oxygen (O 2 ),
(b) Hydrogen (H 2 ),(c) Chlorine (Cl 2 ),
(d) Ozone (O 3 ),
(e) Water (H 2 O),
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IONIC COMPOUNDS
Sodium chloride (NaCl),
Magnesium oxide (MgO)
Sodium oxide (Na 2 O)
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Pure Substances
Any sample of matter can be
classified as a pure substance or
a mixture. A pure substance can be
either an element or a compound.
The composition of a pure substance
is definite and fixed.
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Pure Substances
Example:
2.Pure gold - an element.
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A chemical element or simply
element , is a substance that cannot
substances by ordinary chemical
methods.
The smallest particle of such an
element is an atom , which consistsof electrons centered around a
nucleus of protons and neutrons.
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Examples of elements
(a)Metals - Iron (Fe), Gold (Au),Silver (Ag), Mercury (Hg).
(a)Gases - Oxygen (O 2 ), Nitrogen (N 2 ),Chlorine (Cl 2 ), Helium (He),
Neon (Ne).
(a)Non-metals - Carbon (C).
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C d
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Compounds
A chemical compound is a chemical
substance formed from two or more
elements, with a fixed ratio
determining the composition.
For example, dihydrogen monoxide
(water) is a compound composed oftwo hydrogen atoms for every
oxygen atom, H 2 O.
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Mixtures
A mixture is a combination of two
or more substances, often chemicals,
in which the substances remain
chemically distinct, retaining their
particular composition and
properties.
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Mixtures
There are two types of mixtures:heterogeneous mixtures and
homogeneous mixtures.
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Heterogeneous Mixtures
A Heterogeneous Mixture does not have
uniform properties throughout; the
composition of one part (or phase) .
Example:1. A mixture of oil and water.
2. A mixture of nuts and cake.
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Homogeneous Mixtures
An Homogeneous Mixture is the samethroughout. It has uniform composition
and appearance throughout.
Example:1. A cup of hot coffee.
2. A mixture of alcohol and water.
3. A solution.
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A d I
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Atoms and Ions
An atom consists of three sub-atomic
particles which are electrons,
protons and neutrons.
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A d I
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Atoms and Ions
An ion is an atom or group of atoms
with a net electric charge.
H +
Cl - Na + H +O 2 -
H +
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At d I
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Atoms and Ions
A negatively charged is known as
an anion, and a positively charged
is known as a cation.
Cation Anion
H +
H +
O 2 -O 2-
Na +Cl -H +
H +
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Cations & anions
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At i d M l l M
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Atomic and Molecular Mass
A tomic mass is the mass of one atomexpressed in units ( atomic mass unit ,amu) where 1 amu is equal to 1/12 th
of the actual mass of carbon-12.
Atomic mass is also called as
relative atomic mass or atomicweight.
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Relative Atomic Mass
Relative Atomic Mass is unitless. It is the
atomic mass (amu) divided by the mass
of one C-12 atom (amu).
Relative atomic mass = atomic mass (amu) x 12mass of 12 C (12)
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Average Atomic Mass
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Average Atomic Mass
Most elements have several naturally
occurring isotopes with differentabundance.
The atomic mass shown in the
periodic table for an element is
actually a weighted average of the
masses of all isotopes of the element.
Example 129
Average Atomic Mass
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Average Atomic Mass
- Example 1 -
Bromine :
50.69% bromine - 79, mass = 78.9183
49.31% bromine - 81, mass = 80.9163
? Average atomic mass of Br
= (50.69% x 78.9183)
+ (49.31% x 80.9163)
=79.9035
8030
Average Atomic Mass
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Average Atomic Mass
- Example 2 -
Magnesium :
78.99% magnesium-24, mass = 23.9850
10.00% magnesium-25, mass = 24.9858
11.01% magnesium-26, mass = 25.9826
? Average atomic mass of Mg
= (78.99% x 23.9850) + (10.00% x 24.9858)
+ (11.01% x 25.9826)= 24.3050
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Average Atomic Mass: Exercise
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Average Atomic Mass: Exercise
Isotopes Mass Abundance (%)6 Li 6.015 7.59
7 Li 7.016 92.41
16 O 15.995 99.75717 O 16.999 0.038
18 O 17.999 0.205
Calculate the average atomicmass of lithium and oxygen.
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Carbon 12
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Carbon - 12
Carbon-12 is the most abundant (98.89%)of the two stable isotopes of the elementcarbon. It contains 6 protons, 6 neutronsand 6 electrons.
Carbon-12 is of particular importance asit is used as the standard from which allother isotopes' atomic weight is measured
and thus the measurement of Avogadro'snumber.
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Molecular Mass
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Molecular Mass
The molecular mass of a substance(less accurately called molecular
weight and abbreviated as MW ) is
substance, relative to the unified
atomic mass unit u (equal to 1/12
the mass of one atom of carbon-12).
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Molecular Mass
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Molecular Mass
calculated as the sum of the
atomic masses of all the atoms of
any one molecule.The molar mass of a substance is
numerically equal to the molecular
mass, but expressed in mass unitsper mole (e.g. grams per mole)
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Molecular Mass
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Molecular Mass
For example: the atomic mass of hydrogen
is 1.00784 u and that of oxygen is 15.9994
u; therefore, the molecular mass of water
with formula H 2 O is (2 x 1.00784 u) +
15.9994 u = 18.01508 u.
Therefore, one molecule of water weighs
18.01508 u, and one mole of water weighs
18.01508 grams.
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Molecular Mass
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Molecular Mass
Molecular mass or molar mass areused in stoichiometry calculations.
Since molecules are created bychemical reactions , not nuclearreactions , a molecule's molecularmass exactly equals the sum of the
atomic masses of its constituentatoms.
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Mole Concept
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Mole Concept
The mole (symbol: mol) is one of theseven SI base units and is commonly
used in chemistry.
Mole of substance A = Mass of substance A(g)
Formula Mass of A(g.mol -1 )
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Mole Concept
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Mole Concept
It measures the amount of substance
of a system and is defined as the
amount ofsubstance that contains
as many elementary entities as there
are atoms in exactly 12 grams
carbon-12.
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Mole Concept
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Mole Concept
This quantity is known asAvogadro's number and is
approximately 6.0221415 x 103.
N A = 6.022 x 10 23
or6.02 x 10 23
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Names of Chemical Compo nds
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Names of Chemical Compounds
Different kinds of compounds are named
by different rules. Ordinary table salt is
named sodium chloride because of its
formula, NaCl.
But common table sugar is named
-D-fructofuranosyl- -D-lucopyranoside.[Please don't worry about it!]
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Names of Chemical Compounds
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p
Binary Compounds Acid
(nonmetals)
Ionic CompoundsSalts
PolyatomicHydratesIons
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Names of Binary Compound
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Names of Binary Compound
Chemical Name Symbol
Dinitrogen pentoxide N 2 O 5
Carbon tetrachloride CCl 4
Nitrous oxide N 2 O
Ammonia NH 3
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Names of Ionic Compound
Chemical Name Symbol
Potassium sulfide K 2 S
Iron(II) sulfide FeS
Calcium chloride CaCl 2
Zinc nitrate Zn(NO 3 ) 2
Sodium sulfate Na 2 SO 4
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Names of Hydrates
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Names of Hydrates
Chemical Name
Symbol
Sodium carbonate Na 2 CO 3 .10H2 Odecahydrate (washing
soda)
Magnesium sulfate MgSO 4 .7H 2Oheptahydrate(epsom salt)
Calcium sulfate dihydrate CaSO 4 .2H 2
O(gypsum)Sodium tetraborate Na 2 B 4 O 7 .10H
2 Odecahydrate (borax)
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N f P l t i I
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Names of Polyatomic Ions
Chemical Name Symbol
Cr 2 O 72-Dichromate
ClO 4-Perchlorate
ClO -Hypochlorite
HCO 3-BicarbonateSO 42-Sulfate
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Names of Acid
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Names of Acid
Chemical Name Symbol
Hydroiodic acid HI
Hydrosulfuric acid H 2 S
Phosphoric acid H 3 PO 4
Phosphorus acid H 3 PO 3
Chlorous acid HClO 2
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Names of Salts
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a es o Sa ts
Chemical Name Symbol
Sodium Fluoride NaF
Sodium Sulfite Na 2 SO 3
Sodium bicarbonate NaHCO 3
Sodium Na 2 HPO 3
monohydrogenphosphite
Potassium chlorate KClO 3
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Names of Binary Compounds
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1. Write an appropriateformula for
(a) dinitrogen pentoxide
(b) carbon tetrachloride(c) nitrous oxide
(d) ammonia
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Names of Binary Compounds
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2. Write the chemical namefor(a) NO 2
(b) SO 3
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Names of Ionic Compounds
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What are the names for thecompounds
(a) K 2 S
(b) FeS(c) CaCl 2
(d) Zn(NO 3 ) 2
(e) Na 2 SO 4
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Names of Hydrates
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1. Write formulas for the following
hydrates.
(a) Sodium carbonate decahydrate
(washing soda)
(b) Magnesium sulfate heptahydrate(Epsom salt)
(c) Calcium sulfate dihydrate
(gypsum)
(d) Sodium tetraborate decahydrate(borax)
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Names of Hydrates
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2. Write the chemical names for the
following hydrates.
(a) CoCl 2 .6H 2 O
(b) CuSO 4 .5H 2 O
(c) Na 2 S 2 O 3 .5H 2O
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Names of Polyatomic Ions
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y
Write formulas for thefollowing ions.
(a) chromate
(b) perchlorate
(c) hypochlorite
(d) bicarbonate(e) sulfate
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Names of Polyatomic Ions
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y
Write names for thefollowing ions.(a) PO 43-
(b) Cr 2 O 72-
thiocyanatedichromatethiosulfatephosphatechlorite(c) S 2 O 3 2-
(d) SCN -
(e) ClO 2-
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Names of Acids
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Write names for thefollowing acids.
(a) HF(b) H 2 SO 3
(c) H 2 CO 3
(d) CH 3 COOH(e) HClO
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Names of Acids
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Write formulas for thefollowing acids.
(a) Hydroiodic acid
(b) Hydrosulfuric acid
(c) Phosphoric acid
(d) Phosphorous acid(e) Chlorous acid
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Names of Salts
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Write names for thefollowing typical salts.
(a) NaF(b) Na 2 SO 3
(c) NaHCO 3
(d) Na 2 HPO 3
(e) KClO 3
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Chemical Formula
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There are two types of chemical
formulas:
(a) Empirical Formula,
(b) Molecular Formula .
Empirical formula and molecular
formula are used for different purposes.
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Empirical Formula
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The Empirical Formula gives the simplestwhole-number ratio of atoms of each
element present in a compound.
Example:
Acetylene (C 2 H 2 ) and benzene (C 6 H 6 )
- the simplest ratio C : H = 1 : 1
Empirical Formula = CH
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Molecular Formula
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The Molecular Formula gives the actual
number of atoms of each element present
in a molecule.
Example:a) Acetylene - C 2 H 2
b) Benzene - C 6 H 6
c) Glucose - C 6 H 6 O 6
d) Water - H 2 O
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Exercise 1
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When analyzed, an unknown compound
gave these experimental results: C, 54.0%;H, 6.00%; O, 40.0%. Four different
students used these values to calculate the
empirical formulas shown here. Which
answer is correct? Why did some students
not get the correct answer?(a) C 4 H 5 O 2 (c) C 7 H 10 O 4
(b) C 5 H 7 O 3 (d) C 9 H 12 O 5
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A process in which one or more substances is changed into oneor more new substances is a chemical reaction
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or more new substances is a chemical reaction
A chemical equation uses chemical symbols to show what
happens during a chemical reaction
3 ways of representing the reaction of H 2 with O 2 to form H 2O
productsreactants 63
How to "Read" Chemical Equations
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2 Mg + O 2 2 MgO
2 atoms Mg + 1 molecule O 2 makes 2 formula units Mg
2 moles Mg + 1 mole O 2 makes 2 moles MgO
48.6 grams Mg + 32.0 grams O 2 makes 80.6 g MgO
2 grams Mg + 1 gram O 2 makes 2 g MgO Is thiscorrect?
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Balancing Chemical Equations
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1. Write the correct formula(s) for the reactants on the
left side and the correct formula(s) for the product(s)
on the right side of the equation.
Ethane reacts with oxygen to form carbon dioxide and water
C 2 H 6 + O 2 CO 2 + H 2 O
2. Change the numbers in front of the formulas
( coefficients ) to make the number of atoms of each
element the same on both sides of the equation. Donot change the subscripts.
NOT2C 2 H 6 C 4 H 12
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Balancing Chemical Equations...contd
3 S b b l i h l h i l
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3. Start by balancing those elements that appear in only onereactant and one product.
C 2 H 6 + O 2 start with C or H but not OCO 2 + H 2 O
1 carbon2 carbonmultiply CO 2 by 2on righton left
C 2 H 6 + O 2 2CO 2 + H 2 O
6 hydrogen 2 hydrogenmultiply H 2 O by 3
on left on rightC 2 H 6 + O 2 2CO 2 + 3H 2 O
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Balancing Chemical Equations..contd
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4. Balance those elements that appear in two or more
reactants or products.
multiply O 2 by 7C 2 H 6 + O 2 2CO 2 + 3H 2 O2
4 oxygen + 3 oxygen = 7 oxygen2 oxygenon right(3x1)on left (2x2)
remove fractionC 2 H 6 + 7 O 22CO 2 + 3H 2 O
multiply both sides by 22
2C 2 H 6 + 7O 2 4CO 2 + 6H 2 O
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Balancing Chemical Equations.contd
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5. Check to make sure that you have the same number of
each type of atom on both sides of the equation.
2C 2 H 6 + 7O 2 4CO 2 + 6H 2 O
Reactants Products
4C 4C
12 H 12 H
14 O 14 O
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Mass Changes in Chemical Reactions
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1. Write balanced chemical equation
2. Convert quantities of known substances into moles3. Use coefficients in balanced equation to calculate the number of
moles of the sought quantity
4. Convert moles of sought quantity into desired units69
Methanol burns in air according to the equation
2CH 3 OH + 3O 2 2CO 2 4H 2 O
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2CH 3 OH + 3O 2 2CO 2 + 4H 2 O
If 209 g of methanol are used up in the combustion,what mass of water is produced?
grams CH 3 OH moles CH 3 OH moles H 2 O grams H 2
molar mass molar masscoefficientsCH 3 OH H 2 Ochemical equation
4 mol H 2 O 18.0 g H 2 O1 mol CH 3 OH=209 g CH 3 OH x x x
32.0 g CH 3 OH 2 mol CH 3 OH 1 mol H 2 O
235 g H 2 O
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Chemical Equation and Stoichiometry
(Exercise)
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1. Write the balanced equation for thecombustion of ethyl alcohol:
C 2 H 5 OH (l) + O 2 (g) CO 2 (g) + H 2 O (l)
Answer
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Chemical Equation and Stoichiometry
(Exercise)
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Determine the number of grams oflead(II) sulfide, PbS, that can be oxidised
by 5.22 mol of oxygen gas according to
the following equation.
2PbS(s) + 3O 2 (g) 2PbO(s) + 2SO 2 (g)
Answer
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Volumetric Analysis
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Concentrations of Solutions
Dilution
Titration
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Concentrations of Solutions
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Molarity ( M )
The most widely used to quantify
the concentration of solutions.
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Concentrations of Solutions
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Molarity ( M )
The molarity of a solution is
defined as
moles solutesMolarity= volume of solution ( L )
nor M =V
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Molarity - Exercise 1
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Calculate the molarity of a
solution made by dissolving
23.4 g of sodium sulphate in
enough water to form 125 mL ofsolution.
Answer
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Molarity - Exercise 2
H f di
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How many grams of sodium
sulphate, Na 2 SO 4 are required to
make 0.350 L of 0.500 M Na 2 SO 4 ?
Answer
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Dilution
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Solutions that are used routinelyin the laboratory are often
purchased or prepared in
concentrated form. They are calledas stock solutions .
Example: 12M HCl
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Dilution
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Solutions of lower concentrationscan be obtained by adding water.
This process is called as dilution .
Formula( Initial molarity )( initial volume ) = ( final molarity )( final volume )
or M 1 V 1 = M 2 V 2
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Exercise
How many milliliters of 3 0 M
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How many milliliters of 3.0 M
H 2 SO 4 are required to make 450mL of 0.10 M H 2 SO 4 ?
Answer
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Titration
d i h
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How can we determine the
concentration of a solution?
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Titration
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One common way is to use a secondsolution of known concentration ,
called a standard solution , that
undergoes a specific chemicalreaction of known stoichiometry
with the solution of unknown
concentration .
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This procedure is known as titration .
Example:
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aA + bB cC + dD
Formula( molarity of solution A )( volume of solution A ) ( molarity of solution B )( volume of solution B )=
a b
M a V a M b V b=or
a b 83
Exercise
What is the molarity of an
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What is the molarity of an
automobile battery acid, H 2 SO 4
solution if 22.53 mL of the acid
neutralizes 42.11 mL of 1.923 M
sodium hydroxide, NaOH?
Answer
84