Chapter 2
Atoms
What is Chemistry?
• The study of matter and its properties and
transformations
What is Matter?
• Anything that has mass and volume
– Mass = the amount of a substance, measured in
grams, g
– Volume = the space occupied by a substance,
measured in cm3, mL (milliliters) or L (Liters)
Brief History of Chemistry
• Ideas about matter date back to ancient Greece 
2000 years ago
• 2 schools of thought
Democritus – all matter made of tiny, indivisible
particles called atoms (from Greek word atomos =
uncuttable)
Aristotle – matter is continuous, it is infinitely divisible
• Aristotle’s ideas dominated for almost 2000 yrs
So, who was right?
• Today we know that Democritus was right
• Atoms are the basic building blocks of matter
• we will discuss evidence for the existence of
atoms later
Why do we believe in atoms?
• First atomic theory based on scientific
evidence proposed in 1808 by English chemist
John Dalton (1766 -1844)
• Theory based on three scientific laws
discovered in late 1700s, early1800s
• Law of Conservation of Mass (Antoine
Lavoisier (1743-1794))
– Matter cannot be created or destroyed in an
ordinary chemical reaction
• Law of Constant Composition (Joseph Proust
(1754-1826))
– no matter where you find a specific compound, it
is always made up of the same proportion of
elements by mass  elements combine to form
compounds in fixed proportions
• Law of Multiple Proportions (John Dalton)
– elements always combine to make compounds in
whole number ratios or multiples of whole
number ratios, never in fractions
(Not mentioned in book)
Postulates of Dalton’s Atomic Theory
1. All matter is made of tiny, indivisible particles
called atoms (in honor of Democritus’ idea)
2. All atoms of a given element have the same
properties; atoms of different atoms have
different properties
3. Compounds are formed by the chemical
combinations of two or more different types of
atoms
4. During chemical reactions, atomic arrays are just
rearranged into new combinations
Dalton’s Model of the Atom
• Atoms are solid, indivisible spheres, like
billiard balls
• His model was referred to as the “Billiard Ball”
model
• Dalton’s model of the atom was to endure for
almost 100 years, until the discovery of
radioactivity and the first subatomic particle
(the electron) in the late 1890s.
What are atoms made of?
• First subatomic particle, the electron,
discovered by English Physicist J.J. Thomson in
1897.
• The proton was discovered in 1919 by Ernest
Rutherford
• The last subatomic particle to be identified
was the Neutron in 1932 by James Chadwick.
Properties of Subatomic Particles
Particle Location
Relative Actual
electrical Mass (g)
charge
Relative
Mass
(amu)
Electron
In space
surrounding nucleus
1-
9.11 x 10-28
Proton
In nucleus
1+
1.673 x 10-24
1
Neutron
In nucleus
0
1.675 x 10-24
1
__1__
1840
Terminology for Atomic Structure
Atomic number (Z) – the number of protons in the
nucleus of an atom, also the number of electrons as
atoms are electrically neutral
Mass number (A) – the number of protons and
neutrons in the nucleus of an atom
• Number of neutrons in the nucleus :
#neutrons = mass no. – atomic no.
Subatomic Particles
Determining the number of protons and
electrons in an atom from the periodic table
6
6
12.01
C
C
Atomic number = # protons
=6
Carbonn
= # electrons
Atomic
number
Symbol
=6
Subatomic Particles
Determining the number of neutrons in an
atom: Mass # - Atomic #
- Must be given the mass number!
- Mass number is not the same as the atomic
mass
- e.g. Sodium with a mass number of 23
Na atomic # = 11, 11 protons, 11 eneutrons = 23 – 11 = 12 neutrons
Discovery Of Isotopes
• After neutrons discovered, it was found that
not all atoms of the same element were the
same (as Dalton had said)
• Almost every element has examples of atoms
that have the same number of protons, but
different numbers of neutrons
Isotopes = Atoms that have the same number of
protons (atomic number), but different
numbers of neutrons (different mass numbers)
Nuclear Notation
 Contains
the symbol of the
element, the mass number and the
atomic number.
Mass
number
Atomic
number
X
Examples
Mass number
12
C
6
Atomic number
Number of protons = 6
Number of electrons = 6
Number of neutrons = 12 – 6 = 6
Mass number can also be used at the end of the
element’s name e.g. carbon-12
Isotope Examples
Isotope
Carbon13
Cobalt 58
Sodium23
Mass
number
(A)
Atomic #
#
number protons neutrons
(Z)
#
electrons
Isotope Examples
Isotope
Mass
number
(A)
Atomic #
#
number protons neutrons
(Z)
#
electrons
Carbon13
13
6
6
7
6
Cobalt 58
58
27
27
31
27
Sodium23
23
11
11
12
11
Introduction to the Periodic Table
Dmitri Mendeleev (1834-1907)
• When Mendeleev arranged the elements by
increasing atomic
weight, he noticed a
periodic repetition in
atomic properties (e.g.
density, melting point)
•Because the properties
of the atoms were
repeated periodically, the
table he created was
called periodic table
Mendeleev’s 1872 Table
Modern Periodic Table
• After Moseley discovered atomic number, elements
were rearranged from increasing atomic weight to
increasing atomic number
• Vertical Columns called groups or families.
• Horizontal rows called periods.
Introduction to the Periodic Table
Period number
Group number
IA
1
2
3
4
5
6
7
VIIA
IIA
IIIA IVA VA VIA VII
A
3 Main Categories of Elements
1. Metals – Shiny, good conductors of electricity and
heat, tend to have 3or less valence e-, malleable,
ductile, located to the left of the stair step on the
periodic table
2. Nonmetals – dull, brittle, poor (some non)
conductors of electricity and heat, have 4 or more
valence e-, located to the right of the stair step on
the periodic table
Metalloids
Stair Step
Metals
Nonmetals
3 Main Categories of Elements
3. Metalloids – located along the stair step on
the periodic table, have properties of both
metals and nonmetals
e.g. Silicon – is shiny, brittle, semiconductor
of electricity
Introduction to the Periodic Table
Alkali
metals
Introduction to the Periodic Table
Alkaline
earth
metals
Introduction to the Periodic Table
Halogens
Introduction to the Periodic Table
Noble gases
Introduction to the Periodic Table
Transition
metals
Introduction to the Periodic Table
Lanthanide
series
Actinide
series
Introduction to the Periodic Table
Metalloids
Introduction to the Periodic Table
Main Group
Elements
Periodicity
•
Trend within a group of elements or across a
period of elements in the periodic table
How are the electrons in an atom
arranged?
• Atom is mostly empty space with central
dense core called nucleus
• Electrons are located at a distance away and
have to be constantly moving to avoid being
pulled into the positively charged nucleus
• Because e- are moving, they possess kinetic
energy
• In 1913, Niels Bohr discovered
Niels Bohr 1913
• Discovered that only certain values are
possible for the energy of the hydrogen
electron
• The energy of the electron is quantized 
only certain values are allowed
Quantized Energy Levels
• The energy levels of all atoms are quantized.
• Electrons are confined to specific regions of
space, called principal energy levels or shells
• These energy levels or shells radiate away
from the nucleus and given whole integer
numbers of 1, 2, 3, 4, etc
• Each energy level can accommodate only a
certain number of electrons, given by the
formula 2n2
Energy level
(Shell) n
1
2
3
4
maximum number of
electrons (2n2)
2
8
18
32
Energy levels are further divided into
sublevels or subshells
• Sublevels are designated by the letters s, p, d
and f
n = 1  1 sublevel = 1s
n = 2  2 sublevels = 2s and 2p
n= 3  3 sublevels = 3s, 3p and 3d
n = 4  4 sublevels = 4s, 4p, 4d, and 4f
Within these sublevels, electrons are
grouped in orbitals
Orbital = most probable region in space of finding an
electron
• According to quantum theory, there is a limit to what
we can know about the electron
• Therefore, we can only discuss its location in terms
of probability.
• Orbitals are probability maps that have definite
shapes and orientations in space
• Each orbital can hold a maximum of 2 electrons
Sublevel designation s, p, d and f also
designates the shape of the electron
orbital
S orbitals = spherical an shape
1s
2s
3s
p orbitals
• p orbitals are dumbbell shaped
• There are three p orbital shapes
• The s and p types of sublevel
d Orbitals
7-
f-orbtals
http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/ao4.html
Electron Configuration
Electron configuration: The arrangement of
electrons in the extranuclear space (i.e. the empty
space surrounding the nucleus).
• The energy of the electrons in an atom is quantized,
which means that an electron in an atom can have
only certain allowed energies.
• These allowed energies correspond to specific
regions in space surrounding the nucleus called
energy levels or shells.
Ground-state electron configuration: The electron
configuration of the lowest energy state of an atom.
Electron configurations
• Tell us the orbital location of an atom’s
electrons
• Like an address
Electron Configuration
Table 2.5 Distribution of Electrons in Shells
Relative
N u mb er of
energies
electrons sh ell of electrons
in each sh ell
S hell can hold
4
3
2
1
32
18
8
2
higher
low er
Assigning electrons to orbitals
• Orbital filling diagrams  use boxes or circles to
represent orbitals
• See handout
• Rules for filling orbitals
– Bottom up rule – atoms place their electrons in the lowest
possible energy orbitals first
– Each orbital can hold a maximum of 2 electrons, which
must be spinning in opposite directions
– For p, d and f orbitals, one electron in each orbital before
pairing up
Electron Configuration
Table 2.6 Distribution of Orbitals within
Shells
Maximum Number
of Electrons Shell
Shell Orbitals Contained in Each Shell
Can hold
4 One 4s , three 4 p, five 4 d, and seven 4 f orbitals 2 + 6 + 10 + 14 = 32
2 + 6 + 10 = 18
3 One 3s, three 3 p, and five 3 d orbitals
2+6=8
2 One 2 s and three 2 p orbitals
1 One 1s orbital
2
Electron Configuration
Figure 2.13 Energy levels for orbitals through
the third shell.
Electron Configuration
Electron configurations are governed by
three rules:
Rule 1: Orbitals fill in the order of increasing
energy from lowest to highest.
– Elements in the first, second, and third periods
fill in the order 1s, 2s, 2p, 3s, and 3p.
Electron Configuration
Rule 2: Each orbital can hold up to two
electrons with spins paired in opposite
directions.
– With four electrons, the 1s and 2s orbitals are
filled and are written 1s2 2s2.
– With an additional six electrons, the three 2p
orbitals are filled and are written either 2px2 2py2
2pz2, or they may be written 2p6.
Electron Configuration
Orbitals have definite shapes and
orientations in space
Electron Configuration
Figure 2.14 The pairing of electron spins.
Electron Configuration
Rule 3: When there is a set of orbitals of
equal energy, each orbital becomes half filled
before any of them becomes completely
filled.
– Example: After the 1s and 2s orbitals are filled, a
5th electron is put into the 2px, a 6th into the 2py,
and a 7th into the 2pz. Only after each 2p orbital
has one electron is a second added to any 2p
orbital.
Electron Configuration
Orbital box diagrams
– A box represents an orbital.
– An arrow represents an electron.
– A pair of arrows with heads in opposite
directions represents a pair of electrons with
paired spins.
Example: carbon (atomic number 6)
Electron configuration
Exp anded : 1s2 2s2 2p x1 2py 1
1s
2s
2px 2py 2pz
Con dens ed: 1s2 2s2 2p 2
Electron Configuration
Noble gas notation
– The symbol of the noble gas immediately
preceding the particular atom indicates the
electron configuration of all filled shells
Example: carbon (atomic number 6)
Orbital box diagram
Electron
Configuration
(conden sed)
2
2
1s 2s 2p
2
N oble Gas
N otation
2
[He]2s 2p
2
Electron Configuration
Valence shell: The outermost incomplete shell.
Valence electron: An electron in the valence shell.
Lewis dot structure:
– The symbol of the element represents the nucleus and
filled shells.
– Dots represent valence electrons.
1A
H
2A
3A
4A
5A
6A
7A
8A
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Electron Configuration
Main group elemen ts;
s block (2 elemen ts )
1A
Main group elemen ts;
p block (6 elemen ts )
8A
3A 4A 5A 6A 7A 1s 1
2p
2
Tran sition elemen ts;
d block (10 elemen ts )
1 1s 2A
2s
2
Heliu m is
als o an s block
element
3
4
5
3s
5s
4d
4p
3
4
5
6
6s
5d
6p
6
7
7s
6d
7p
7
4s
3B 4B 5B 6B 7B 8B 8B 8B 1B 2B
3d
Inn er transition
elements; f block
(14 elemen ts)
6
7
4f
5f
3p
5p
Electron Configuration
Table 2.9 Noble Gas Notation and Lewis dot
structures for the Alkali Metals (Group 1A
Elements)
N ob le
Lew is
Gas
dot
Element N otation Structure
Li
[He]2s1
Li•
Na
[N e]3s1
N a•
K
[A r]4s1
K•
Rb
[Kr]5s1
Rb•
Cs
[Xe]6s1
Cs •
3
Li
6.941
11
Na
22.990
19
K
39.098
37
Rb
85.468
55
Cs
132.91
1A
Periodic Property
• As we have seen, the Periodic Table was
constructed on the basis of trends (periodicity) in
chemical properties.
• With an understanding of electron configuration,
chemists realized that the periodicity of chemical
properties could be understood in terms of
periodicity in electron configuration.
• The Periodic Table worked because elements in the
same column (group) have the same configuration
in their outer shells.
• We look at two periodic properties: Atomic size and
Atomic Size
The size of an atom is determined by the size of its
outermost occupied orbital.
• Example: The size of a chlorine atom is determined by the
size of its three 3p orbitals, the size of a carbon atom is
determined by the size of if its three 2p orbitals.
Cl
Cl
C
C
154 pm
198 pm
The radiu s of a chlorine
atom is 99 pm
The radiu s of a carbon
atom is 77 pm
Figure 2.16
Atomic radii of
the maingroup
elements (in
picometers).
Atomic Size
Ionization Energy
Ionization energy: The energy required to
remove the most loosely held electron from
an atom in the gaseous state.
– Example: When lithium loses one electron, it
becomes a lithium ion; it still has three protons
in its nucleus, but now only two electrons
outside the nucleus, and therefore has a positive
charge.
+
Li (g) + en ergy
Lith iu m Ionization
energy
Li (g)
Lithiu m
ion
+ e
Electron
Ionization Energy
Ionization energy is a periodic property:
– In general, it increases across a row; valence electrons are
in the same shell and subject to increasing attraction as
the number of protons in the nucleus increases.
– It increases going up a column; the valence electrons are
in lower principle energy levels, which are closer to the
nucleus and feel the nuclear charge more strongly.