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CHEM 101 - General Chemistry 1
Chapter 2 – Atoms and Elements
Handout prepared by Dr.Saravanan Rajendrasozhan,
Department of Chemistry, University of Hail.
Assistant
Professor,
Reference book: ‘Principles of Chemistry: A Molecular Approach, First edition, by
Nivaldo Tro. Published by Pearson Education International.
MATTER
Matter is anything that has both mass and volume (occupies space). It can be a
solid, liquid or gas.
Early Philosophy of Matter
¾ Because of the lack of sophisticated instruments in the early days, there are
two different opinions given by philosophers about matter.
1) Matter had an ultimate, tiny, indivisible particle (Leucippus and
Democritus)
2) Matter was infinitely divisible (Plato and Aristotle)
¾ In the late 17th century, research on matter was improved that could not
explain the infinitely divisible matter concept.
Atoms
¾ Atoms are basic unit of matter or building blocks of matter.
Example: Carbon (C), Oxygen (O), Hydrogen (H), Gold (Au)
Elements
¾ A pure chemical substance consisting of one type of atom.
Example: Pure gold bar (made of only gold atoms)
¾ Atoms are the smallest particle of the element which retains the properties
of that element.
Pure gold bar
(element)
Gold bar (element)
consisting of only gold atoms
Gold atom
(smallest particle
of the element)
FUNDAMENTAL CHEMICAL LAWS
Scientific law is a description of an observed phenomenon.
Law of Conversion of Matter
¾ In a chemical reaction, mass (weight) is neither created nor destroyed.
Total mass of reactants = Total mass of products
Reactants
Products
Reaction
2Na (s) + Cl2 (g)
2NaCl (s)
Mass of chemicals
7.7 g Na + 11.9 g Cl2
19.6 g NaCl
Total mass
19.6 g
19.6 g
Law of Definite Proportions
A given compound always contains exactly the same proportion of elements by
mass (regardless of their source or how they were prepared).
Sample 1
100 g Sodium chloride salt (NaCl)
produced by reacting sodium metal
and chlorine gas
Sample 2
58.44 g of NaCl prepared by
seawater evaporation process
Elements: 39.3 g Sodium and 60.7 g Elements: 22.99 g Sodium and
chlorine
35.44 g chlorine
Proportion=Mass of Cl/Mass of Na
Proportion=Mass of Cl/Mass of Na
= 60.7 g / 39.3 g
= 35.44 g / 22.99 g
= 1.54
= 1.541
Law of Multiple Proportions
¾ When two elements (A and B) combine to form two different compounds,
they combine in a ratio of small whole numbers.
¾ The masses of B that combine with 1 g of A can be expressed as a ratio of
small whole numbers.
Example:
Carbon and Oxygen can combine at different ratio to form carbon monoxide
(CO) and carbon dioxide (CO2).
CHEM 101‐Chapter 2 Handout Page 2 CO: 1 g carbon combine with
1.33 g oxygen
CO: 0.752 g carbon combine with
1 g oxygen
CO2: 1 g carbon combine with
2.67 g oxygen
CO2: 0.375 g carbon combine
with 1 g oxygen
Ratio of oxygen combine with
1 g carbon = 2.67 g/1.33 g
Ratio of carbon combine with
1 g oxygen = 0.752 g/0.375 g
=2
=2
Dalton’s Atomic Theory
Scientific theory is an explanation of the observed phenomenon.
1) Each element is composed of tiny, indestructible particles called atoms.
2) All atoms of a given element have the same mass and other properties that
distinguish them from atoms of other elements.
3) Atoms combine in simple, whole-number ratios to form molecules of
compounds.
4) In a chemical reaction, atoms of one element cannot change into atoms of
another element.
During a reaction, the atoms are simply rearranging the way they are
attached.
EARLY EXPERIMENTS TO CHARACTERIZE ATOMS
¾ Three experiments are considered important in characterizing atoms.
1) Thomson’s Experiment – Postulated the existence of electrons.
2) Millikan’s Oil Drop Experiment – Measured the mass and charge of
electrons.
3) Rutherford’s Experiment – Proposed the nuclear model of atom.
Thomson’s Experiment
¾ Glass tube (almost all the air has been evacuated) was connected to metal
electrodes.
¾ When high voltage is applied, the cathode starts to emit a glowing ray,
which is called cathode ray.
¾ When electric field applied around the glass tube, the cathode ray deflects
towards the positively charged plates.
¾ The ray also deflects when magnetic field is applied.
-
From the observation, Thomson postulated that cathode ray contains
some particles with negative charge and mass.
CHEM 101‐Chapter 2 Handout Page 3 -
Also calculated the charge/mass ratio of the particle: −1.76 × 108 C/g
¾ These cathode ray particles became known as electrons.
Thomson’s Plum Pudding Model of Atom
¾ The structure of the atom contains many negatively
charged small particles, called electrons.
¾ The electrons are held within a positively charged
sphere (positively charged cloud).
¾ This model of atom is called as palm-pudding model
(Later experimetns proved that this model is not
correct).
Millikan’s Oil Drop Experiment
¾ The apparatus contains 2 horizontal metal plate electrodes (anode and
cathode).
¾ A fine mist of oil droplets was sprayed (by atomizer) into a chamber above
the plates.
¾ When the oil droplets enter the space between the electrodes (through the
hole in plate electrode), the oil droplets were ionized by applying ionizing
radiation such as x-ray.
¾ The ionized droplets could be
made to rise and fall by changing
the voltage across the electrodes
(anode and cathode).
¾ From this experiment, the mass
and charge of electron was
calculated.
Mass of electron = 9.1 X 10-28 g
Charge of electron = −1.60 × 1019 C.
Radioactivity
¾ The discovery radioactivity initiated further experiments to characterize the
atomic structure.
¾ Certain elements would constantly emit small, energetic particles and rays.
This is called radioactivity.
¾ These particles and rays can penetrate matter.
Rays
Charge
Form
α: Alpha particles
Positive charge
Particles with a mass of 4H atoms
β: Beta particles
Negative charge
Particles with a mass ~1/2000 H atom
γ: Gamma Rays
No charge
Energy rays (not particles)
Rutherford’s Experiment
¾ Alpha particles (positively charged) were directed at an ultra-thin sheet of
gold foil.
¾ The deflection in alpha rays was monitored.
#
Results
Conclusion
1
98% of α-particles penetrate gold
foil and went straight through.
An atom is mostly empty space.
2
0.005% of α-particles bounced
back from the gold foil.
An atom contains a dense particle
(small in volume compared to the
atom but large in mass).
2
2% of the α-particles penetrate
gold foil, but deflected by large
angles.
The dense particle in the atom is
positively charged which deflect
positively charged α-particles
Rutherford’s Nuclear Model of Atom
¾ The atom has a tiny dense center called nucleus. It is responsible for the
entire mass of an atom and it has positively charged particles (named
protons).
¾ Most of the volume of atom is empty space. The negatively charged electrons
are dispersed in the empty space surrounding the nucleus.
¾ The atoms are neutral as it has equal number of positively charged protons
and negatively charged electrons.
¾ Later, Rutherford and his student J. Chadwick discovered the presence of
neutrons (no charge) in the nucleus of the atom. Protons and neutrons are
collectively responsible for the mass of an atom.
STRUCTURE OF ATOM
¾ Thus, an atom consist of
1) Protons
2) Neutrons
3) Electrons
¾ Protons and neutrons are present in the tiny nucleus at high density. Both
of them accounts for almost all the atom’s mass (atomic mass).
¾ Electrons form the outer shell of the atom, so that it allows one atom to
interact with another. Thus, it is responsible for the chemical properties of
the atom.
¾ Atoms are neutral in charge. So they must have equal number of positively
charged protons and negatively charged electrons.
Subatomic
Particle
(& Charge)
Proton
Location Charge
+1
1.6726 x10-24 Determine the identity of an
atom.
Nucleus
0
1.6749 x10-24 Prevent the repulsion
between positively charged
protons and help to pack
the protons in nucleus.
Move
around
nucleus
-1
0.0009 x10-24 Determine the size and
chemical properties of the
atom.
(No charge)
Electron
(Negative)
Role
Nucleus
(Positive)
Neutron
Mass (g)
CHEM 101‐Chapter 2 Handout Page 6 Changes in the Structure of Atom
¾ The number of protons and neutrons in an atom cannot be changed by the
chemical reaction.
¾ The number of electrons in an atom may be changed during the chemical
reaction to form a charged atoms (positively or negatively charged ions).
ELEMENTS and Number of Protons
¾ The elements were identified by the number of protons in the nucleus.
¾ To identify the different elements (atoms), we have a standard way of
representing each element:
Chemical symbol with atomic number and mass number.
Chemical Symbol: One or two letter abbreviation (One capital letter or one capital
letter with one lower case letter). It is based on current name or old Latin name.
Current Name
Latin Name
Symbol
Hydrogen
-
H
Carbon
-
C
Calcium
-
Ca
Oxygen
-
O
Iron
Ferrum
Fe
Copper
Cuprum
Cu
Silver
Argentum
Ag
Sodium
Natrium
Na
Atomic number (Z) = Number of protons
Mass number (A) = Number of protons + Number of neutrons
¾ Atomic number (Z) is written as a subscript to the left of chemical symbol.
¾ Mass number (A) is written as a superscript on the left of chemical symbol.
¾ Z is always smaller than A (except for Hydrogen).
M
A
Example:
N
N
Carbon
Chemical symbol = C, Atomic number = 6, Mass number = 12
The element (carbon) is written as:
CHEM 101‐Chapter 2 Handout Page 7 ¾ Surprisingly in the periodic table, atomic number (Z) is mentioned on top of
the chemical symbol.
Example: Carbon is written as
6
C
Exercise:
How do you symbolize an element with the atomic number of 12 and mass
number of 24.
Z = 12, A = 24, From the periodic table: Element with the atomic number 12 = Mg
Number of Proton, Neutron and Electrons
Number of protons = Atomic number (Z)
Number of electrons = Number of protons (only for neutral atoms)
Number of neutrons = Mass number - Number of protons
(because, Mass number = Number of protons + Number of neutrons)
Example: How many protons, electrons, and neutrons are in an atom of
?
Atomic number (Z) = 24
Mass number (A) = 52
Number of protons = 24
Number of electrons = 24
Number of neutrons = 52 – 24 = 28
Exercise:
Calculate the number of proton, electron and neutrons in the following atoms:
Atom
# of Protons
# of Electrons
# of Neutrons
Pb
82
82
208-82 = 126
F
9
9
19-9 = 10
Ca
20
20
41-20 = 21
U
92
92
235-92 = 143
ISOTOPES and Number of Neutrons
¾ Atoms with the same atomic number but different mass numbers are called
isotopes.
¾ Isotopes have same number of proton and electron, but different number of
neutron.
CHEM 101‐Chapter 2 Handout Page 8 ¾ As the chemical property of the atoms depends mainly on electrons, the
isotopes have same chemical property.
Example: Sodium ( Na) has 11 electrons and 11 protons.
If sodium has 12 neutrons, A = 11+12 = 23, it is written as
Na or Na-23.
If sodium has 13 neutrons, A = 11+13 = 24, it is written as
Na or Na-24.
9
Na (Na-23) and
Na (Na-24) are called as isotopes.
CHARGED ATOMS and Number of Electrons
¾ Atoms or group of atoms (molecules) with a net positive or negative charge
are called ions.
Example: Na+, NH4+, Cl–, OH–
¾ Ions are classified into 2 groups based on the number of atoms.
1) Mono-atomic ions (Example: Na+, Cl–)
2) Poly-atomic ions (Example: NH4+, OH–)
The net charge of each atom is equal to the charge of poly-atomic ion.
¾ Ions are classified into 2 groups based on the charge.
1) Positive ions, called cations (Example: Na+, NH4+)
2) Negative ions, called anions (Example: Cl–, OH–)
¾ Only the number of electrons is changed in ions as compared to the parent
atoms (NO change in the number of protons or neutrons).
Cations
¾ When atoms or molecules lose one or more electrons, they become positively
charged ions, called cations.
¾ The name of the cation is same as the element.
Example 1:
Sodium ion ( Na+) formation from sodium ( Na) atom.
Na
Lose one electron
⇓
11 Protons, 11 Electrons
Net charge = +11-11 = 0
Neutral
Na+
⇓
11 Protons, 10 Electrons
Net charge = +11-10 = +1
Positive charge
Example 2:
Calcium ion ( Ca2+) formation from calcium ( Ca) atom.
Ca
⇓
20 Protons, 20 Electrons
Lose two electrons
Ca2+
⇓
20 Protons, 18 Electrons
Anions
¾ When atoms or molecules gain one or more electrons, they become
negatively charged ions, called anions.
¾ Anions are named by changing the ending of the element name to -ide.
Example 1:
Chloride ion ( Cl–) formation from chlorine atom ( Cl).
Cl
Gain of one electron
⇓
17 Protons, 17 Electrons
Net charge = +17-17 = 0
Neutral
Cl–
⇓
17 Protons, 18 Electrons
Net charge = +11-18 = -1
Negative charge
Example 2:
Oxide ion ( O2–) formation from oxygen ( O) atom.
O
O2–
Gain two electrons
⇓
8 Protons, 8 Electrons
⇓
8 Protons, 10 Electrons
Name of common monoatomic anions:
Name of Atom
Name of Anion
Name of Atom
Name of Anion
H: Hydrogen
H– : Hydride
O: Oxygen
O2– : oxide
F: Fluorine
F– : Fluoride
S: Sulfur
S2– : Sulfide
Br: Bromine
Br– : Bromide
N: Nitrogen
N3– : Nitride
I: Iodine
I– : Iodide
P: Phosphorus
P3– : Phosphide
Cl: Chlorine
Cl– : Chloride
Number of Proton, Neutron and Electrons in Ions
¾ The number of protons and neutrons in an ion is same as that of the parent
atom.
Example: The number of protons and neutrons in sodium ion (Na+) is same
as that of sodium atom (Na).
¾ The number of electrons in an ion may be increased or decreased as
compared to number of proton (atomic number).
Number of electrons in an ion = Atomic number – Charge
Exercise:
Calculate the number of proton and electron in the following ions:
Ions
Z (from Periodic
table)
# of Protons
# of Electrons
Zn 2+
30
30
30-(+2) = 28
N3–
7
7
7-(-3) = 10
Br–
35
35
35-(-1) = 36
Ions and Compounds
¾ Ions behave much differently than the neutral atom.
Example: Sodium metal (Na) is highly reactive and quite unstable.
Sodium ion (Na+) found in table salt (NaCl) is non-reactive and
stable.
¾ Cations (positive charged ions) and
anions (negative charged ions) attract
each other because of the opposite
electrical charges.
¾ Cation and anion with equal amount of
positive and negative charges join
together to form a compound, because
compounds are neutral (Net charge=0).
9 Cation: Na+ (charge: +1)
9 Anion: Cl– (charge: -1)
9 Compound: 1Na+ + 1Cl– → NaCl
¾ If the cation and anion has different amount of positive and negative
charges, multiple ions join together to make equal positive and negative
charges to form a compound.
9 Cation: Mg2+ (charge: +2), Anion: Cl– (charge: -1), Compound: MgCl2
9 Cation: Fe3+ (charge: +3), Anion: O2– (charge: -2), Compound: Fe2O3
CHEM 101‐Chapter 2 Handout Page 11 PERIODIC TABLE
¾ All the known elements are arranged in periodic table depending on the
atomic number (Z) and chemical properties.
¾ Each row (left to right) of the periodic table is called ‘Period’.
¾ Each column (top to bottom) of the periodic table is called ‘Group’.
9 Atomic number is mentioned on top of the chemical symbol of element. It
is increased from left to right of periodic table (period).
9 The chemical properties of elements are similar in each group (top to
bottom of periodic table).
¾ From the periodic table you can find out number of protons and number of
electrons present in an atom.
Example:
2
Atomic number (Z) = 2
He
Number of protons = 2
Number of electrons = 2
Periodic Pattern
¾ Periodic pattern are specific trends that are present in the periodic table,
which help to predict the properties of elements.
¾ The main periodic trends include: metallic character, electronegativity,
ionization energy, electron affinity and atomic radius.
CHEM 101‐Chapter 2 Handout Page 12 Pattern in Metallic Character
¾ Metallic character (ability to lose electrons) decreases from left to right of
the periodic table.
¾ Metallic character increases from top to bottom of the periodic table.
¾ Based on the metallic properties, the elements are classified into:
1) Metals,
2) Non-metals, and
3)Metalloids
Metals
¾ Metals conduct heat and electricity.
¾ Metals are solids at room temperature (except Hg). They are malleable (can
be shaped) and ductile (pulled into wires).
¾ Most of the elements (about 75%) are metals.
¾ Metals frequently lose electrons to become positive ions (cations).
¾ It can form ionic compound with a non-metal ion (negatively charged
anion).
¾ Metals are in the left side and middle of the periodic table.
Non-metals
¾ Nonmetals are poor conductors of heat and electricity.
¾ Nonmetals found in all three states (solid, liquid, gas) at room temperature.
Solids are brittle (easily breakable), so they are NOT malleable and NOT
ductile.
CHEM 101‐Chapter 2 Handout Page 13 ¾ Non-metals frequently gain electrons to become negative ions (anions).
¾ They can form ionic compound with a metal ion (positively charged cation).
¾ They can also form covalent compound by bonding with another nonmetal.
¾
Non-metals are in the right side of the periodic table (diagonal from B to At)
and H.
Metalloids
¾ Metalloids show some properties of metals and some properties of
nonmetals.
¾ They are semiconductors of heat and electricity.
¾ They are present in between the metals and nonmetals in periodic table.
MODERN PERIODIC TABLE
¾ Elements with similar chemical and physical properties are in the same
column.
Remember: Columns are called Groups. Rows are called Periods
¾ They are classified into 3 groups:
1) Main group elements
2) Transition elements
3) Inner-transition elements.
1) Main group elements
CHEM 101‐Chapter 2 Handout Page 14 9 Also called as Representative elements or Group A elements
(present in Group 1A to 8A).
9 They are both metals and non-metals.
2) Transition elements
9 Also called as Group B elements (present in Group 1B to 8B)
9 They are all metals.
3) Inner-transition elements
9 Also called as Rare Earth Elements (present in bottom rows,
really belongs to periods 6 and 7)
9 They are all metals.
Important Groups
¾ Important groups in the periodic table are:
Hydrogen, Alkalai metals, Alkalai earth metals, Halogens and Noble gases.
Hydrogen
¾ Nonmetal present on top of the Group 1A.
¾ Colorless, diatomic (H2) gas with very low melting point and low density.
9 Reacts with nonmetals to form molecular compounds (Example: HCl
and H2O)
¾ Reacts with metals to form hydrides. Metal hydrides react with water to
form H2 gas.
¾ HX dissolves in water to form acids.
Alkalai Metals
¾ Metals present in group 1A (Li, Na, K, Rb, Cs, Fr) react with water to form
alkaline solutions (pH>7, basic). So they are called alkaline metals.
Example: 2 Na + 2 H2O
2 NaOH + H2 + Heat
¾ Hydrogen is also present in Group 1A, but it is a nonmetal (NOT belongs to
alkalai metals).
¾ They are soft solids with low melting point and low density.
¾ They are very reactive.
9 React with nonmetals to form water-soluble compounds.
Example: 2Na + Cl2
2 NaCl
CHEM 101‐Chapter 2 Handout Page 15 Alkalai Earth Metals
¾ Metals present in group 2A (Be, Mg, Ca, Sr, Ba, Ra) react with oxygen to
form metal oxides, which can dissolve in water to form alkaline solution.
Example: 2Mg + O2
2MgO
All these metals are found in Earth and remain solids (earths) in fire. So
they are called alkaline earth metals.
¾ They are hard solids with very high melting point and high density.
¾ They are very reactive, but the reactivity is lesser than alkaline metals
(Group 1A).
¾ Alkaline earth metals (except Be-beryllium) react with water to form
hydrogen.
Example: Ba + 2H2O
Ba(OH)2 + H2
Halogens
¾ Nonmetals present in group 7A (F, Cl, Br, I, At).
In Greek, ‘hal’ means ‘salt’ and ‘gen’ means ‘to produce’. Non-metals in
Group 7A react with sodium to produce salt of similar properties. So they
are called Halogens.
Example: Cl2 + 2Na
2NaCl
¾ Exist as diatomic molecules (F2 and Cl2 are gases, Br2 is liquid, and I2 is
solid).
¾ Very reactive. Reacts with metals to form ionic compounds.
¾ All HX are acids (Acidity: HI < HCl < HBr < HI)
¾ Cl2, Br2 react slowly with water.
Example: Br2 + H2O
HBr + HOBr
Noble Gases
¾ Nonmetals present in group 8A (He, Ne, Ar, Kr, Xe, Rn) are called noble
gases.
¾ They are all gases at room temperature. They have very low melting and
boiling points.
¾ They are most stable or very unreactive (practically inert). Also very hard to
produce ions (very hard to remove electron from or give an electron to noble
gas).
¾ They are odorless and colorless, and are used in many conditions when a
stable element is needed to maintain a safe and constant environment.
CHEM 101‐Chapter 2 Handout Page 16 IONIC CHARGE AND PERIODIC TABLE
¾ Different ions may have different magnitude of charge.
Example: Na+, Ca2+, Fe3+, Cl–, O2–
¾ Charge on an ion can be determined from an element’s position on the
periodic table.
¾ Metals always form positively charged cations.
9 For the main group metals (Groups 1A, 2A and Al from Group 3A)
Charge of Cation = Group number
¾ Nonmetals form negatively charged anions.
9 For many main group nonmetals (Groups 7A, 6A and N from Group 5A)
Charge of Anion = Group number – 8
¾ Charge on transition metal ion cannot be determined from its position in
the periodic table.
¾ Some transition metals form only one cations, but some transition metals
(and some metals form Groups 3A and 4A) form more than one cations.
Example: Ag can form only Ag+, but Fe can form Fe2+ and Fe3+
Transition
metal
Charge of Ion
Element
Charge of Ion
Ag
+1
Cu
+1 or +2
Zn, Cd
+2
Hg
Hg2+, Hg22+
Cr, Mn, Fe, Co
+2 or +3
Sn, Pb
+2 or +4
CHEM 101‐Chapter 2 Handout Page 17 ATOMIC MASS
¾ Atomic mass (atomic weight) is the mass of an atom in atomic mass units
(amu).
¾ Atomic mass can be determined using an instrument called ‘Mass
Spectrometer’.
¾ Atomic mass can also be derived from the mass number of the atom.
9 Mass number (A) = Number of Protons + Number of Neutrons
(Remember: It is indicated in the superscript in left side of the chemical
symbol or in the bottom of chemical symbol in periodic table).
Atomic mass = Mass number in amu
Example:
Atomic mass of carbon ( C) = 12 amu
Average Atomic Mass
¾ Many elements present in nature as mixtures of isotopes (same atomic
number but different mass number).
Example:
Carbon
C (C-12)
98.89% present as
1.11% present as
C (C-13)
< 0.01 % present as
C (C-14)
¾ Average atomic mass of an element is calculated from the atomic mass of
various isotopes of the elements and its natural abundance (percentage of
isotope present in the nature).
Average Atomic mass of C = (12 x 98.89) + (13 x 1.11) + (14 x 0.01)
100
= (1186.68 + 14.43 + 0.14)/100 = 12.0125
= 12.01 amu
¾ Average atomic mass of element is listed directly beneath the element’s
symbol in periodic table.
Example:
Atomic Number (Z) → 6
C
Atomic Mass → 12.01
← Chemical Symbol or
Element’s Symbol
COUNTING ATOMS BY MOLES
¾ If we can find the mass of a particular number of atoms, we can convert the
mass of atoms into number of atoms in the sample.
¾ The unit ‘mole’ is used to find out the mass of particular number of atoms.
CHEM 101‐Chapter 2 Handout Page 18 Mole
¾ Mole (also written as mol) is the unit of measurement used to measure the
number of things, usually atoms or molecules.
1 mole = 6.022 x 1023 thing (atom or molecule)
6.022 x 1023 is Avogadro's number.
Example: 1 mole of carbon = 6.022 x 1023 carbon atoms
Exercise 1: Calculate the number of atoms in 1.9 moles of gold.
1 mol gold = 6.022 x 1023 gold atoms
23
Number of gold atoms = 1.9 mol gold x 6.022 x 10 atoms
1 mol gold
= 11.44 x 1023 atoms = 1.144 x 1024 atoms
Exercise 2: A copper coin contains 1.7 x 1026 atoms. How many moles of copper
are there in the coin?
1 mol copper = 6.022 x 1023 copper atoms
1 mol
Moles of copper = 1.7 x 1026 atoms x
6.022 x 1023 atoms
= 0.28 x 103 mol = 2.8 x 102 mol
Molar mass
¾ The mass of one mole of atoms is called the molar mass.
¾ Molar mass of an atom = Mass number or Atomic mass in grams
Example 1: What is the molar mass of
Molar mass of
C?
C = 12 grams
Thus, Mass of 1 mole (6.022x1023) of
C atoms = 12 grams
Example 2: What is the mass of carbon (C) atoms?
Atomic mass of C (from Periodic Table) = 12.01 amu
Molar mass of C = 12.01 gram
Substance Chemical Mass of 1 atom Atoms in 1 mole
Symbol (atomic mass)*
Mass of 1 mole
(molar mass)
Hydrogen
H
1.008 amu
6.022 x 1023 atoms
1.008 g
Carbon
C
12.01 amu
6.022 x 1023 atoms
12.01 g
Sulfur
S
32.06 amu
6.022 x 1023 atoms
32.06 g
*Atomic mass is taken from the Periodic Table
CHEM 101‐Chapter 2 Handout Page 19 Exercise 1: Calculate the number of moles in 7.2 grams of sodium metal.
Atomic mass of sodium (Na) = 22.99 amu
(Remember: Atomic mass is mentioned in the periodic table at the bottom of
chemical symbol)
Molar mass = 22.9 g
Therefore, 1 mol sodium = 22.99 g
Moles of sodium = 7.2 g sodium x
1 mol
22.9 g sodium
= 0.314 moles
Exercise 2: How many carbon atoms present in 0.265 g of pencil lead.
Atomic mass of carbon = 12.01 amu
1 mole carbon = 12.01 g (Molar mass)
1 mole = 6.022 x 1023 atoms
Number of carbon (C) atoms = 0.265 g C x
1 mol
12.01 g C
x
6.022 x 1023 atoms
1 mol
= 0.133 x 1023 atoms
= 1.33 x 1022 atoms
CHEM 101‐Chapter 2 Handout Page 20 REMEMBER
Atomic number = Number of protons
Mass number = Number of protons + Number of neutrons
Written symbol
Mass Number →
Atomic Number →
Periodic table
Atomic Number →
← Chemical
…..Symbol
Atomic Mass →
← Chemical
…..Symbol
.003 Number of protons = Atomic number
Number of electron in neutral atom = Number of protons = Atomic number
Number of electrons in ion = Atomic number – Charge of the ion
Number of neutron = Mass number – Atomic number
Cations = Metal ions = Ions with positive charge
Anions = Non-metal ions = Ions with negative charge
Charge of cations (+ve) = Group number
Charge of anions (-ve) = Group number – 8
Atomic mass = Mass of 1 atom = Mass number in amu (atomic mass unit)
Molar mass = Mass of 1 mole of atoms = Mass number or Atomic mass in grams
Avogadro number = 6.022 X 1023
1 mole = 6.022 X 1023 units (atoms or molecules)
Mole = Mass (g) / Molar mass (g/mol)
Mass (g) = Mole X Molar mass (g/mol)
Number of atoms = Mole X 6.022 X 1023
Mole = Number of atoms/6.022 X 1023
HOMEWORK
From the book ‘Principles of Chemistry: A Molecular Approach, First edition,
Nivaldo Tro.
Page # 33 to 37
Exercises: 24, 26, 30, 34, 36, 40, 45, 48, 50, 56