Atoms, Molecules and
Periodic Table
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Dalton’s Atomic Theory (1808)
1. Elements are composed of extremely small particles
called atoms.
2. All atoms of a given element are identical, having the
same size, mass and chemical properties. The atoms of
one element are different from the atoms of all other
elements.
3. Compounds are composed of atoms of more than one
element.
4. A chemical reaction involves only the separation,
combination, or rearrangement of atoms; it does not
result in their creation or destruction.
John Dalton
• 1800 -Dalton proposed a modern atomic model
based on experimentation not on pure reason.
•
•
•
•
All matter is made of atoms.
Atoms of an element are identical.
Each element has different atoms.
Atoms of different elements combine
in constant ratios to form compounds.
• Atoms are rearranged in reactions.
• His ideas account for the law of conservation of
mass (atoms are neither created nor destroyed)
and the law of constant composition (elements
combine in fixed ratios).
Dalton’s Atomic Theory
Law of Multiple Proportions
16 X
+
8Y
Law of Conservation of Mass
8 X2Y
Ernest Rutherford (movie: 10 min.)
• Rutherford shot alpha () particles at gold foil.
Zinc sulfide
Thin gold foil
screen
Lead block
Radioactive
substance path of invisible
-particles
Most particles passed through.
So, atoms are mostly empty.
Some positive -particles
deflected or bounced back!
Thus, a “nucleus” is positive &
holds most of an atom’s mass.
Rutherford’s Model of
the Atom
atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
“If the atom is the Houston
Astrodome, then the nucleus is a
marble on the 50-yard line.”
mass p ≈ mass n ≈ 1840 x mass e-
Atomic number, Mass number and Isotopes
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different
numbers of neutrons in their nuclei
Mass Number
A
ZX
Atomic Number
1
1H
235
92
2
1H
U
Element Symbol
(D)
238
92
3
1H
U
(T)
The Isotopes of Hydrogen
How many protons, neutrons, and electrons are
14
in 6 C ?
6 protons, 8 (14 – 6) neutrons, 6 electrons
How many protons, neutrons, and electrons are
11
in 6 C ?
6 protons, 5 (11 – 6) neutrons, 6 electrons
Properties of Waves
Wavelength (l) is the distance between identical
points on successive waves.
Amplitude is the vertical distance from the
midline of a wave to the peak or trough.
Frequency (n) is the number of waves that pass
through a particular point in 1 second (Hz = 1
cycle/s). The speed (u) of the wave = l x n
Bohr’s Model of
the Atom (1913)
1. e- can only have
specific (quantized)
energy values
2. light is emitted as emoves from one
energy level to a
lower energy level
1
En = -RH ( 2
n
)
n (principal quantum number) = 1,2,3,…
RH (Rydberg constant) = 2.18 x 10-18J
E = hn
E = hn
The Planck constant was first discovered as the
proportionality constant between the energy (E) of a
photon and the frequency of its associated
electromagnetic wave (ν) length. This relation between the
energy and frequency is called the Planck relation or the
Planck–Einstein equation:
E=hv
h=6.626×10−34
Since the frequency ν, wavelength λ, and speed of
light c are related by λν = c, the Planck relation can
also be expressed as
c=3x108 speed of the
light
Bohr’s model
• Electrons orbit the nucleus in “shells”
• Electrons can be bumped up to a higher
shell if hit by an electron or a photon of light.
There are 2 types of spectra: continuous spectra &
line spectra. It’s when electrons fall back down that
they release a photon. These jumps down from
“shell” to “shell” account for the line spectra seen in
gas discharge tubes (through spectroscopes).
Schrodinger Wave Equation
y is a function of four numbers called
quantum numbers (n, l, ml, ms)
Quantum Numbers of Electrons in Atoms
Name
Symbol
Permitted Values
Property
principal
n
positive integers(1,2,3,…) orbital energy (size)
angular
momentum
l
integers from 0 to n-1
orbital shape (The l values
0, 1, 2, and 3 correspond to
s, p, d, and f orbitals,
respectively.)
magnetic
ml
integers from -l to 0 to +l
orbital orientation
spin
ms
+1/2 or -1/2
direction of e- spin
The Pauli Exclusion Principle - No two electrons in the same atom can have the
same four q.n. Since the first three q.n. define the orbital, this means only two
electrons can be in the same orbital and they must have opposite spins.
Schrodinger Wave Equation
y is a function of four numbers called
quantum numbers (n, l, ml, ms)
principal quantum number n
n = 1, 2, 3, 4, ….
distance of e- from the nucleus
n=1
n=2
n=3
Schrodinger Wave Equation
quantum numbers: (n, l, ml, ms)
angular momentum quantum number l
for a given value of n, l = 0, 1, 2, 3, … n-1
from 0 to n-1
n = 1, l = 0
n = 2, l = 0 or 1
n = 3, l = 0, 1, or 2
l=0
l=1
l=2
l=3
s orbital
p orbital
d orbital
f orbital
Shape of the “volume” of space that the e- occupies
l = 0 (s orbitals)
l = 1 (p orbitals)
l = 2 (d orbitals)
Clover leaf
Schrodinger Wave Equation
quantum numbers: (n, l, ml, ms)
magnetic quantum number ml
for a given value of l
ml = -l, …., 0, …. +l
From –l to 0 then to + l
if l = 1 (p orbital), ml = -1, 0, or 1
if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2
orientation of the orbital in space
ml = -1, 0, or 1
3 orientations is space
ml = -2, -1, 0, 1, or 2
5 orientations is space
Schrodinger Wave Equation
(n, l, ml, ms)
spin quantum number ms
ms = +½ or -½
ms = +½
ms = -½
Schrodinger Wave Equation
quantum numbers: (n, l, ml, ms)
Shell – electrons with the same value of n
Subshell – electrons with the same values of n and l
Orbital – electrons with the same values of n, l, and ml
How many electrons can an orbital hold?
If n, l, and ml are fixed, then ms = ½ or - ½
y = (n, l, ml, ½) or y = (n, l, ml, -½)
An orbital can hold 2 electrons
How many 2p orbitals are there in an atom?
n=2
2p
If l = 1, then ml = -1, 0, or +1
3 orbitals
l=1
How many electrons can be placed in the 3d
subshell?
n=3 If l = 2, then m = -2, -1, 0, +1, or +2
l
3d
l=2
5 orbitals which can hold a total of 10 e-
Energy of orbitals in a single electron atom
Energy only depends on principal quantum number n
n=3
n=2
En = -RH (
n=1
1
n2
)
The most stable arrangement of electrons in
subshells is the one with the greatest number of
parallel spins (Hund’s rule).
s has one
orbital,
p has 3
orbitals,
d has 5
orbitals.
Each
orbitals can
hold upto
two electrons
Order of orbitals (filling) in multi-electron atom
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
Electron configuration is how the electrons are
distributed among the various atomic orbitals in an
atom.
number of electrons
in the orbital or subshell
1s1
principal quantum
number n
angular momentum
quantum number l
Orbital diagram
H
1s1
Isoelectronics, the number of electrons on
each ions the same In an isoelectronic
series of ions, the nuclear charge increases
with increasing atomic number and draws
the electrons inward with greater force. The
ion with fewest proton produces the
weakest attractive force on the electrons
and thus has the largest size
34
GROUND STATE VS EXCITED STATE of Electron
energy state.
The energy associated to an electron is that of its
orbital. The energy of a configuration is often
approximated as the sum of the energy of each
electron, neglecting the electron-electron
interactions.
The configuration that corresponds to the lowest
electronic energy is called the ground state. Any
other configuration is an excited state.
What is the electron configuration of Mg?
Mg 12 electrons
1s < 2s < 2p < 3s < 3p < 4s
1s22s22p63s2
2 + 2 + 6 + 2 = 12 electrons
Abbreviated as [Ne]3s2
[Ne] 1s22s22p6
What are the possible quantum numbers for the last
(outermost) electron in Cl?
Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s
1s22s22p63s23p5
2 + 2 + 6 + 2 + 5 = 17 electrons
Last electron added to 3p orbital
36
n = 3 l = 1 ml = -1, 0, or +1 ms = ½ or -½
Outermost subshell being filled with electrons
Paramagnetic
unpaired electrons
2p
Diamagnetic
all electrons paired
2p
THE MODERN PERIODIC TABLE
The Periodic Table
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elemen
4f
5f
42
Classification of the Elements
Nobel Gases
Metals
43
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1 Al3+ [Ne]
Atoms gain electrons
so that anion has a
noble-gas outer
electron configuration.
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
44
-3
-2
-1
+3
+1
+2
Cations and Anions Of Representative Elements
45
Effective nuclear charge (Zeff) is the “positive
charge” felt by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Z
Core Zeff
Radius (pm)
Na
11
10
1
186
Mg
12
10
2
160
Al
13
10
3
143
Si
14
10
4
132
46
Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
47
Atomic Radii
metallic radius
covalent radius
48
49
Cation is always smaller than atom from
which it is formed.
Anion is always larger than atom from
which it is formed.
50
Ionization energy is the minimum energy (kJ/mol)
required to remove an electron from a gaseous
atom in its ground state.
I1 + X (g)
+ eI2 + X+(g)
eI3 + X2+(g)
+ e-
X+(g)
I1 first ionization energy
X2+(g) +
I2 second ionization energy
X3+(g)
I3 third ionization energy
I1 < I 2 < I3
51
General Trends in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
52
Electron affinity is the negative of the energy
change that occurs when an electron is accepted by
an atom in the gaseous state to form an anion.
X (g) + e-
X-(g)
F (g) + e-
X-(g) DH = -328 kJ/mol EA = +328 kJ/mol
O (g) + e-
O-(g) DH = -141 kJ/mol EA = +141 kJ/mol
53
54
Types of Elements
Metals,
Nonmetals,
metalloids,
transition elements,
55
Metals, found on both the left side and the in the
middle of the table, include active metals,
transition metals and the lanthanide and actinide
series of elements.
They are shiny solids except for mercury. They
have ability to deform without breaking which is
called malleability. They have high melting
points, and density.
They have low electronegativity, large atomic
radius, low ionization energy. They easily give
up electrones.
Nonmetals are found on the upper right side of the
table. Nonmetals are generally brittle in the solid
state and show little or no metallic luster. They have
high ionization energy, electron affinity and electro
negativity. They have small atomic radii.
Metalloids are also called semi metals. They
posses characteristics of metals and nonmetals.
For example Silicon has a metallic luster but is
brittle and a poor electric conductor
57
Group 1A Elements (ns1, n  2)
M
M+1 + 1e-
2M(s) + 2H2O(l)
2M2O(s)
Increasing reactivity
4M(s) + O2(g)
2MOH(aq) + H2(g)
58
Group 2A Elements (ns2, n  2)
59
Group 2A Elements (ns2, n  2)
M
M+2 + 2e-
Be(s) + 2H2O(l)
Mg(s) + 2H2O(g)
Mg(OH)2(aq) + H2(g)
M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba
Increasing reactivity
M(s) + 2H2O(l)
No Reaction
60
Group 3A Elements (ns2np1, n  2)
4Al(s) + 3O2(g)
2Al(s) + 6H+(aq)
2Al2O3(s)
2Al3+(aq) + 3H2(g)
61
Group 3A Elements (ns2np1, n  2)
62
Group 4A Elements (ns2np2, n  2)
Sn(s) + 2H+(aq)
Sn2+(aq) + H2 (g)
Pb(s) + 2H+(aq)
Pb2+(aq) + H2 (g)
63
Group 4A Elements (ns2np2, n  2)
64
Group 5A Elements (ns2np3, n  2)
N2O5(s) + H2O(l)
P4O10(s) + 6H2O(l)
2HNO3(aq)
4H3PO4(aq)
65
Group 5A Elements (ns2np3, n  2)
66
Group 6A Elements (ns2np4, n  2) (oxygen
family=chalcogen)
SO3(g) + H2O(l)
H2SO4(aq)
67
Group 6A Elements (ns2np4, n  2)
68
Group 7A Elements (ns2np5, n  2)
X2(g) + H2(g)
X-1
2HX(g)
Increasing reactivity
X + 1e-
69
Group 7A Elements (ns2np5, n  2)
70
Group 8A Elements (ns2np6, n  2)
Completely filled ns and np subshells.
Highest ionization energy of all elements.
No tendency to accept extra electrons.
71
Compounds of the Noble Gases
A number of xenon compounds XeF4,
XeO3, XeO4, XeOF4 exist.
A few krypton compounds (KrF2, for
example) have been prepared.
72
Properties of Oxides Across a Period
basic
acidic
73
The Modern Periodic Table
Noble Gas
Halogen
Group
Alkali Metal
Alkali Earth Metal
Period
A molecule is an aggregate of two or more atoms in a
definite arrangement held together by chemical forces
H2
H2O
NH3
CH4
A diatomic molecule contains only two atoms
H2, N2, O2, Br2, HCl, CO
diatomic elements
A polyatomic molecule contains more than two atoms
O3, H2O, NH3, CH4
An ion is an atom, or group of atoms, that has a net
positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
Cl–
17 protons
18 electrons
A monatomic ion contains only one atom
Na+, Cl-, Ca2+, O2-, Al3+, N3-
A polyatomic ion contains more than one atom
OH-, CN-, NH4+, NO3-
Common Ions Shown on the Periodic Table
How many protons and electrons are in
27 3+
13 Al
?
13 protons, 10 (13 – 3) electrons
How many protons and electrons are in
78
2Se
?
34
34 protons, 36 (34 + 2) electrons
A molecular formula shows the exact number of
atoms of each element in the smallest unit of a
substance
An empirical formula shows the simplest
whole-number ratio of the atoms in a substance
molecular
empirical
H2O
H2O
C6H12O6
CH2O
O3
O
N2H4
NH2
Ionic compounds consist of a combination of cations
and an anions
• The formula is usually the same as the empirical formula
• The sum of the charges on the cation(s) and anion(s) in
each formula unit must equal zero
The ionic compound NaCl
Formula of Ionic Compounds
2 x +3 = +6
3 x -2 = -6
Al2O3
Al3+
1 x +2 = +2
Ca2+
1 x +2 = +2
Na+
O22 x -1 = -2
CaBr2
Br1 x -2 = -2
Na2CO3
CO32-
The most reactive metals (green) and the most reactive
nonmetals (blue) combine to form ionic compounds.
Chemical Nomenclature
• Ionic Compounds
– Often a metal + nonmetal
– Anion (nonmetal), add “ide” to element name
BaCl2
barium chloride
K2O
potassium oxide
Mg(OH)2
magnesium hydroxide
KNO3
potassium nitrate
• Transition metal ionic compounds
– indicate charge on metal with Roman numerals
FeCl2
2 Cl- -2 so Fe is +2
iron(II) chloride
FeCl3
3 Cl- -3 so Fe is +3
iron(III) chloride
Cr2S3
3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide
• Molecular compounds
− Element furthest to the left in a
period and closest to the bottom of a
group on periodic table is placed first
in formula
− If more than one compound can be
formed from the same elements, use
prefixes to indicate number of each kind
of atom
− Last element name ends in ide
Molecular Compounds
HI
hydrogen iodide
NF3
nitrogen trifluoride
SO2
sulfur dioxide
N2Cl4
dinitrogen tetrachloride
NO2
nitrogen dioxide
N2O
dinitrogen monoxide
1.
Which of the following increases with
increasing atomic number within a family (group)
on the periodic table?
a. electro negativity
b. electron affinity
c. atomic radius
d. ionization energy
2.
Silicon has a silvery luster at
room temperature. Silicon is brittle,
and does not conduct heat or
electricity well. Based on its position
in the periodic table, silicon is most
likely a:
a.
b.
c.
d.
Nonmetal
Metalloid
Metal
Chalcogen (Oxygen group)
3.
A natural sample of carbon contains
99% of 12C. How many moles of 12C are
likely to be fund in a 48.5 gram samle of
carbon obtained from nature?
a. 1
b. 2
c. 4
d. 12
e. 49.5
4. Which of the following is the correct
electron configuration for Zn2+
a.
b.
c.
d.
1s22s22p63s23p64s03d10
1s22s22p63s23p64s23d8
1s22s22p63s23p64s23d10
1s22s22p63s23p64s03d8
91
5.
Which of the following quantum
number sets describes a possible
element?
a.
b.
c.
d.
n=2; l=2; m1=1; ms=+1/2
n=2; l=1; m1=-1; ms=+1/2
n=2; l=0; m1=-1; ms=-1/2
n=2; l=0; m1=-1; ms=-1/2
92
6.
What is the maximum number of
electrons allowed in a single atomic
energy level in terms of the principal
quantum number of n?
a.
b.
c.
d.
2n
2n+2
2n2
2n2+2
93
8. An electron returns from an excited state to its
ground state, emitting a photon at 500nm.
What would be the magnitude of the energy
change if this process were repeated such that
a mole of these photons were emitted?
a.
b.
c.
d.
3.98x10-19 J
3.98x10-21 J
2.39x105 J
2.39x10-3 J
94
Answer 8
E=(6.26x10-34)x((3X108)/500 x10-9m)
E=(6.26x10-34)x(6x1014)
E=3.98x10-19J this is per photon
E=3.98x10-19x 6.022x1023=2.39x105J
95
Q.9 Which of the following statement is
NOT true of an electrons grounds state?
a. The electro is at its lowest
possible energy level.
b.The electron is in a quantized
energy level,
c. The electon is traveling along its
smallest possible orbital rdius
d. The electron is static
96
Q. 10
Which of the following experimental
conditions would NOT excite an electron out of
the ground state?
a. Radiation
b. High temperature
c. High pressure
d. None of the above
97
Q11. What determines the length of
an elements atomic radius?
I. the number of valence
electrons
II. The number of electron
shells
III. The number of neutrons in
the nucleus
a. I only
b. II only
c. I and II only
d. I, II and III
Q. 12. How many valence electrons are
present in elements third period?
a. 2
b. 3
c. The number decreases as the
atomic number increases
d. The number increases as the
atomic number increases
99