Download Introductory Chemistry, 2nd Edition Nivaldo Tro

Atoms & the
Periodic Table
Chapter Outline
•
•
•
•
•
•
What is Atom?
Chemical properties of Atoms: the Periodicity
Isotopes
Electrons in Atom: Quantum physics’ view
Valence electrons and the Periodic Table
Periodic trend: Atomic Radius, Metallic
Character, Ionization Energy
2
Experiencing Atoms
• Atoms: incredibly small, yet compose everything
• atoms are the pieces of Elements
• properties of the atoms determine the properties
of the elements
3
Within an Atom
• Atoms = (Protons +
Neutrons) + Electrons
• The nucleus (Protons +
Neutrons) is only about 10-13
cm in diameter yet with most
of the mass of the atom
• The electrons move outside
the nucleus with an average
distance of about 10-8 cm
• The atom is electrically
neutral : #proton (#p) =
#electron (#e)
Proton
Neutron
Electron
Nucleus
4
Comparison among Proton,
Electron, Neutron
Subatomic
Particle
Mass
g
Mass Location Charge Symb
amu in atom
ol
Proton
1.67 x 10-24
Electron
9 x 10-28
Neutron
1.67 x 10-24
1
nucleus
+1 p, p+, H+
~0 empty space -1
e, e-
1
n, n0
nucleus
0
5
Elements
• each Element has a unique number of protons
in its nucleus
• Atomic number: the number of Protons in
the nucleus of an atom
the elements are arranged on the Periodic Table
in order of their atomic numbers
• Name and Symbol of an Element
symbol either one or two letters
one capital letter or one capital letter + one lower case
6
The Periodic Table of Elements
7
The Size of Atoms
Atomic Mass Unit (amu):
1 amu = 1.66  10-24 g
Hydrogen the smallest atom
mass of H atom= 1.67 x 10-24g ~ 1 amu
volume of H atom = 2.1 x 10-25cm3
8
Isotopes
The same element could have atoms with different
masses
Examples:
• 2 isotopes of chlorine atoms in nature: one
weighs about 35 amu (Cl-35); another weighs
about 37 amu (Cl-37)
• Carbon-12 (C-12) is much more abundant than
C-13.
9
Isotopes
all isotopes of an element:
• chemically identical
undergo the exact same chemical reactions
• the same number of protons
• different masses due to different numbers of
neutrons. Example: C-14 atom has eight neutrons;
C-12 atom has six neutrons.
• identified by their mass numbers
protons + neutrons
10
Isotopes
• Atomic Number (Z)
 Number of protons
• Mass Number (A)
 Protons + Neutrons
Abundance = relative amount
found in a sample
Example: Cl-35 (75%) vs. Cl-37
(25%)
11
Isotopic Symbol
• Cl-35 has a mass number = 35, 17 protons and 18
neutrons (35 - 17). The symbol for this isotope
would be
35
17
Cl
Atomic Symbol
A = mass number
Z = atomic number
#neutrons = A - Z
AX
Z
12
Example:
How many protons, neutrons,
and electrons in an atom of
238
92
U
Isotopic symbol  element
 atomic number
 #p  #e
 #n
Mass number = Atomic number (# protons, or #p) + #neutrons
U = uranium
Atomic Number = 92
#p = atomic number = 92
#e = #p = 92
Mass Number = #p + #n
238 = 92 + #n
146 = #n
#proton = 92
#neutron = 146
#electron = 92
13
Mass Number is Not the Same
as Atomic Mass
• Atomic mass (or Atomic Weight) is an experimental
number determined from all naturally occurring
isotopes
• Mass number refers to the number of protons +
neutrons in one isotope
natural or man-made
When given the relative abundance of all isotopes, we
can find the Atomic mass
14
The Modern Periodic Table
• Elements with similar chemical and
physical properties are in the same column
• columns are called Groups or Families
designated by a number and letter at top
• rows are called Periods
• each period shows the pattern of properties
repeated in the next period
15
The Modern Periodic Table
• Main Group = Representative Elements = ‘A’
groups
• Transition Metals = ‘B’ groups
Aka Transition elements
• Inner Transition Elements = Bottom rows =
Rare Earth Elements
metals
really belong in Period 6 & 7
16
Main group vs. Transition metals, Inner transition metals
= Metal
= Metalloid
= Nonmetal
IA
VIIIA
IIA
IIIA
IIIB
VIIB
VIIIB
IB IIB
Metals: Physical vs. Chemical Properties
• solids at room temperature, except Hg
• reflective surface
 shiny
• conduct heat, electricity
• Malleable (can be shaped)
• Tend to Lose electrons and form
Cations in reactions. Na  Na+ + e • about 75% of the elements are metals
• lower left on the table
18
Nonmetals: Physical vs. Chemical Properties
• Elements found in all 3 states
• poor conductors of heat or
electricity
• solids are brittle
• Tend to gain electrons in
reactions to become anions:
Cl + e -  Cl-
• upper right on the table
except H
19
Metalloids: between Metals and
Nonmetals
• show some
properties of metals
and some of
nonmetals
• also known as
semiconductors
Properties of Silicon
shiny
conducts electricity
does not conduct heat well
brittle
20
= Alkali Metals
= Halogens
= Alkaline Earth Metals
= Lanthanides
= Noble Gases
= Actinides
= Transition Metals
21
= Transition Metals
= Rare Earth Metals
= Transuranium element
U
22
Important Element - Hydrogen
• nonmetal
• colorless, diatomic gas H2
 very low melting point & density
• reacts with Nonmetals to form molecular
compounds
 HCl is acidic gas
 H2O is a liquid
• reacts with Metals to form hydrides
 metal hydrides react with water to form H2
 Nickel-metal hydride (NiMH) used in rechargeable
battery
• HX dissolves in water to form acids
23
Important Element - Carbon
Three forms of pure carbon:
• Diamond: hardest substance in nature
• Graphite: soft and slippery solid
• Buckminsterfullerene: a molecule made of 60
(Images from public domain wikipedia.com)
• carbon atoms in a sphere
24
Carbon as backbone for
Organic/Biochemical Molecules
Carbon atoms capable of forming robust
bonds with many other elements and
themselves.
Examples:
• Small molecules: Butane, Sugar, Fatty acid,
Vitamins
• Big molecules (Polymers): Starch, Kevlar,
Teflon, Protein, and DNA
25
Group IA: Alkali Metals
• Usually Hydrogen is included
• All metals: soft, low melting
points
• Flame tests  Li = red, Na =
yellow, K = violet
lithium
sodium
Chemical Property:
• Very reactive. React with water
to form basic (alkaline) solutions potassium
and H2.
releases a lot of heat
rubidium
• Tend to form water soluble
compounds, such as table salt
and baking soda.
cesium
colorless solutions
26
Group IIA: Alkali Earth Metals
Physical properties: harder, higher
melting, and denser than alkali
metals
• flame tests  Ca = red, Sr = red,
Ba = yellow-green
Chemical properties:
• reactive, but less than
corresponding alkali metal
• form stable, insoluble oxides.
oxides are basic = alkaline earth
• reactivity with water to form H2,
 Be = none; Mg = steam; Ca,
Sr, Ba = cold water
beryllium
magnesium
calcium
strontium
barium
27
Group VIIA: Halogens
•
•
•
•
•
nonmetals
F2 & Cl2 gases; Br2 liquid; I2 solid
all diatomic
very reactive
Cl2, Br2 react slowly with water
Cl2 + H2O  HCl + HOCl
(chlorine)
• react with metals to form ionic
compounds
• HX all acids
 HF weak < HCl < HBr < HI
fluorine
chlorine
bromine
iodine
28
Group VIIIA: Noble Gases
• all gases at room
temperature,
very low melting and
boiling points
• very unreactive,
practically inert
• very hard to remove
electron from or give an
electron to
29
Atomic Orbitals
Quantum Physicists including
Schrödinger:
• Electrons move very fast
around the nucleus
• Electrons show up with a
particular probability at certain
location of the atom
• Orbital: A region where the
electrons show up a very high
probability when it has a
particular amount of energy
 generally set at 90 or 95%
30
Electron Shells
• Electron Shell: the main energy level for the
orbital. Principal quantum number n = 1, 2, …
For a chlorine atom, three shells of electrons:
• The innermost shell (n = 1, 2 electrons) has the
lowest energy
• The outmost shell (n = 3, 7 electrons) has the
highest energy
7e-
Cl (17p+ & 17e-)
8e2e17 p+
31
Each Shell have Subshells
• Each Electron Shell has one or more Subshells (s, p,
d, f)
 the number of subshells = the Principal quantum number n
n = 1, one subshell (1s);
n = 2, two subshells (2s, 2p)
n = 3, three subshells (3s, 3p, 3d)
n = 4, three subshells (4s, 4p, 4d, 4f)
• each Subshell has orbitals with a particular shape
 the shape represents the probability map
 90% probability of finding electron in that region
32
Shapes of Subshells
s Orbital
p Orbitals: px , py , pz
d Orbitals
33
f orbitals
34
Tro: Chemistry: A Molecular Approach, 2/e
34
Shapes of f orbitals: 4f orbitals
(downloaded from public domain)
The coloration corresponds to the sign of function.
35
Shells & Subshells
36
How does the 1s Subshell Differ
from the 2s Subshell?
37
Subshells and Orbitals
• Among the subshells of a principal shell, slightly different
energies
 for multielectron atoms, the subshells have different
energies: s < p < d < f
• each subshell contains one or more Orbitals
s : 1 orbital
p : 3 orbitals
d : 5 orbitals
f : 7 orbitals
within one subshell, different orbitals have the same
energy. Example: 2px, 2py and 2pz
38
Electron Configurations
Definition: The distribution of electrons into the
various energy shells (n = 1,2,3,…) and subshells
(s, p, d, f) in an atom in its ground state
• Each energy shell and subshell has a maximum
number of electrons it can hold
Subshell s = 2, p = 6, d = 10, f = 14
Shell n: 1 = 2e, 2 = 8e, 3 = 18e, 4 = 32e
• Electrons fill in the energy shells and subshells in
order of energy, from low energy up
 Aufbau Principal (“Construction” in German)
39
7s
6s
Energy
5s
4s
6p
5p
6
d
5d
5f
4f
4d
4p
3d
3p
3s
2p
2s
1s
40
Order of Subshell Filling
in Ground State Electron Configurations
1. Diagram
putting each energy shell on
a row and listing the subshells,
(s, p, d, f), for that shell in
order of energy, (left-to-right)
2. draw arrows through
the diagonals, looping back
to the next diagonal
each time
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
41
Spinning Electron(s) in Orbital
• Experiments showed Electrons spin on an axis
 generating their own magnetic field
Pauli Exclusion Principle
• each Orbital may have a maximum of 2 electrons,
with opposite spin
• Two electrons sharing the same orbital must have
Opposite spins
 so there magnetic fields will cancel
 analogous to two bar magnets in parallel: only opposite
alignment could stabilize each other.
42
Orbital Diagrams
• often an orbital as a square
• the electrons in that orbital as arrows
 the direction of the arrow represents the spin of the
electron
unoccupied
orbital
orbital with
1 electron
orbital with
2 electrons
43
Filling the Orbitals in a Subshell
with Electrons
• Energy shells fill from lowest energy to high
1 → 2 → 3 → 4
• Subshells fill from lowest energy to high
s → p → d → f
• Orbitals of the same subshell have the same
energy. Three 2p orbitals; Five 3d orbitals
 Electrons prefer “spreading out” in orbitals of
same subshell before they pair up in orbitals.
 Hund’s Rule
 Example: 2p3 _ _ _ instead of  ____
44
Electron Configuration of Atoms
in their Ground State
• Electron configuration: a listing of the subshells in order
of filling with the number of electrons in that subshell
written as a superscript
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
• a shorthand way : use the symbol of the previous noble
gas in [] for the inner electrons, then just write the last set
Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1
45
Example: Ground State Orbital Diagram and
Electron Configuration of Magnesium
1s22s22p63s2 = [Ne]3s2


1s
2s
  
2p

3s
3p
46
Practice: Write Electron Configuration for the
following atoms at the Ground state
•
•
•
•
Calcium
Sulfur
Sodium
Chlorine
Important: you are required to be able to write electron
configuration for first three rows of elements
47
Valence Electron vs. Core Electron
Valence Electron: the electrons in all the subshells with
the highest principal energy shell
• Example: electrons in bold
Mg = [Ne]3s2 O = [He]2s22p4
Br = [Ar]4s23d104p5
• Core electrons: electrons in lower energy shells
• Chemists have observed that one of the most important
factors in the way an atom behaves, both chemically
and physically, is the Number of Valence electrons
48
Valence Electrons
Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1
• the highest principal energy shell that contains
electrons is the 5th : 1 valence electron + 36 core
electrons
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
• the highest principal energy shell that contains
electrons is the 4th : 8 valence electrons + 28 core
electrons
49
Electrons Configurations and
the Periodic Table
50
Electron Configurations from
the Periodic Table
Example: Be 2s2 B 2s22p1 C 2s22p2 N 2s22p3 O 2s22p4
• Elements in the same period (row) have Valence
Electrons in the same principal energy shell.
• #Valence electrons increases by one from left to right
• Elements in the same group have the same #valence
electron and they are same kind of subshell
Example: IIA: Be 2s2 Ca 3s2 Sr 4s2 Ba 5s2
•
VIIA: F 2s22p5 Cl 3s23p5 Br 4s24p5 I 5s25p5
51
Electron Configuration & the
Periodic Table
• Elements in the same Group have similar
chemical and physical properties  their
valence shell electron configuration is the same
• No. Valence electrons for the main group
elements is the same as the Group Number
Example:
• Group IA: ns1 ;
• Group IIIA: ns2np1
• Group VIIA: ns2np5
52
s1
1
2
3
4
5
6
7
Electron Configuration & the
Periodic Table p1 p2 p3 p4
s2
p 5 s2
p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
53
Electron Configuration from
the Periodic Table
• Inner electron configuration = Noble gas of the
preceding period
• Outer electron configuration: from the preceding
Noble gas the next period (Subshells) 
Element
the valence energy shell = the period number
the d block is always one energy shell below the
period number and the f is two energy shells below
54
Electron configuration &
Chemical Reactivity
• Chemical properties of the elements are
largely determined by No. Valence
electrons
• Why elements in groups? Since elements in
the same column have the same #valence
electrons, they show similar properties
55
Electron Configuration:
Noble Gas
• Noble gases have 8 valence electrons
except for He, which has only 2
electrons
• Noble gases are especially nonreactive
He and Ne are practically inert
 The reason: the electron configuration of
the noble gases is especially stable
56
Everyone Wants to Be Like a Noble Gas!
Alkali Metals (Group 1A)
• have one more electron than the previous
noble gas, [NG]ns1
• tend to lose their extra ONE electron,
resulting in the same electron
configuration as a noble gas
 forming a cation with a 1+ charge
 Na  Na+
 Li  Li+
57
Everyone Wants to Be Like a Noble Gas!
Halogens (Group 7A)
• one fewer electron than the next noble gas:
[NG]ns2np5
• Reactions with Metals: tend to gain an
electron and attain the electron configuration
of the next noble gas: [NG]ns2np5 + 1e 
[NG]ns2np6
 forming an anion with charge 1-: Cl  Cl-
• Reactions with Nonmetals: tend to share
electrons so that each attains the electron
configuration of a noble gas
58
Everyone Wants to Be Like a Noble Gas!
Summary
• Alkali Metals as a group are the most reactive
metals
 they react with many things and do so rapidly
• Halogens are the most reactive group of nonmetals
• one reason for their high reactivity: they are only
ONE electron away from having a very stable
electron configuration
 the same as a noble gas
59
Stable Electron Configuration
And Ion Charge
• Metals:  Cations
Atom
by losing enough
electrons to get the
same electron
Na
configuration as the
Mg
previous noble gas
• Nonmetals:  Anions Al
by gaining enough
O
electrons to get the
F
same electron
configuration as the
next noble gas
Atom’s
Electron
Config
[Ne]3s1
Na+
Ion’s
Electron
Config
[Ne]
[Ne]3s2
Mg2+
[Ne]
[Ne]3s23p1
Al3+
[Ne]
[He]2s2p4
O2-
[Ne]
[He]2s22p5
F-
[Ne]
Ion
60
Trends in Atomic Size
Increases down a group
valence shell farther from nucleus
effective nuclear charge fairly close
Decreases across a period (left to right)
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer
61
Trends in Atomic Size
62
Metallic Character
• Metals
 malleable & ductile
 shiny, lusterous, reflect light
 conduct heat and electricity
 most oxides basic and ionic
 form cations in solution
 lose electrons in reactions – oxidized
• Nonmetals
 brittle in solid state
 dull
 electrical and thermal insulators
 most oxides are acidic and molecular
 form anions and polyatomic anions
 gain electrons in reactions - reduced
63
Trends in Metallic Character
64
Electron Configuration Affects the
Size of Atoms and Metallic Character:
Within a Group
• Within the same Group, from top to
bottom:
As quantum number n increases for the valence
electron(s)
 valence electron(s) further away from the
nucleus
 Larger Atomic Radius
 weaker Coulombic force (electrostatic force)
withholding valence electrons
 electrons easier to be lost
 Stronger metallic character
65
Example: Group IIA
q q 
F  k
2
r
2e2e-
Be (4p+ & 4e-)
4 p+
2e-
Mg (12p+ & 12e-)
8e2e12 p+
2e8e-
Ca (20p+ & 20e-)
8e2e20 p+
66
Electron Configuration Affects the
Size of Atoms and Metallic Character:
Over the Period
• Within the same Period (row), from left to right:
Same quantum number n for the valence
electron(s)
As Nucleus has increasing number of protons
(p+)
 Stronger Coulombic force (electrostatic force)
withholding valence electrons
 Valence Electrons closer the nuclues
 Smaller Atomic Radius
Valence electrons harder to be lost
 Weaker metallic character
67
Example: Period 2
1e2e3+
2e2e4+
Li (3p+ & 3e-)
4e2e6+
C
(6p+
&
3e2e5+
Be (4p+ & 4e-) B (5p+ & 5e-)
6e- 2e
8+
6e-)
q q 
F  k
r2
8e-2e
10+
O (8p+ & 8e-) Ne (10p+ & 10e-)
68
Ionization Energy (IE)
In an atom, electrons (“-” charge) are
attracted to the nucleus (“+” charge).
Energy is required to remove the
electron from an atom.
Na + energy  Na+ + eNeutral atom
IE
Cation
Higher IE corresponds to lower Metallic
property.
69
Trends in Ionization Energy
Decreases down a group
valence shell farther from nucleus
effective nuclear charge fairly close
Increases across a period (left to right)
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer
70
Practice: Rank elements K, Mg, S, F:
A. increasing metallic character
B. increasing atomic radii
C. increasing ionization energy
F, S, Mg, K
F, S, Mg, K
K, Mg, S, F
71