1
Periodic Trends in
Atomic Properties
2
Characteristic properties and trends of the
elements are the basis of the periodic
table’s design.
3
These trends allow us to use the periodic
table to accurately predict properties and
reactions of a wide variety of substances.
4
Metals and Nonmetals
5
Chemical Properties
of Metals
• metals tend to lose
electrons and form
positive ions (cations).
Chemical Properties
of Nonmetals
• nonmetals tend to gain
electrons and form
negative ions (anions).
When metals react with nonmetals electrons
are usually transferred from the metal to the
nonmetal.
6
Nonmetals
arefound
foundtotothe
theleft
right
metalloids.
Metals are
ofofthethemetalloids
7
11.1
Atomic Radius
8
Radii of atoms
increase down a
group.
For each step down a group, electrons enter
the next higher energy level.
9
Radii of atoms tend to decrease
from left to right across a period.
This increase
For
Each
time an in
representative
electron
positive
nuclear
is added a
elements
proton
charge
is
pulls
awithin
added
all to
the nucleus.
electrons
same period
closer to
the energy
nucleus.level
remains constant
as electrons are
added.
10
Ionization Energy
11
The ionization energy of an atom is the
energy required to remove the outermost
electron from an atom.
Na + ionization energy → Na+ + e-
12
Ionization energies gradually increase from left to
right across a period.
Noble
Gases
1
2
VIIA
VA
IA
IVA
IIA
VIA
3
4
IIIA
Periodic relationship of the first ionization energy for
representative elements in the first four periods.
13
Gases
nonmetals have higher ionization
potentials than metals
VIIA
VA
IA
VIA
IVA
IIA
IIIA
Distance of Outer Shell Electrons From Nucleus
Ionization energies of Group A elements decrease
from top to bottom in a group.
Noble
nonmetals
metals
Periodic relationship of the first ionization energy for
representative elements in the first four periods.
14
Lewis “Dot” Structures of Atoms
15
Metals form cations and nonmetals form
anions to attain a stable valence electron
structure.
16
Thesestable
This
rearrangements
structure often
occur
consists
by losing,
of twogaining,
s and
sixsharing
or
p electrons.
electrons.
17
The Lewis structure of an atom is a
representation that shows the valence
electrons for that atom.
• Na with the electron structure 1s22s22p63s1
has 1 valence electron.
• Fluorine with the electron structure 1s22s22p5
has 7 valence electrons
18
The Lewis structure of an atom uses dots to
show the valence electrons of atoms.
Paired
electrons
B
Unpaired
electron
Symbol of
the element
2
1
2s 2p
The number of dots equals the number of s
and p electrons in the atom’s outermost shell.
19
The Lewis structure of an atom uses dots to
show the valence electrons of atoms.
S
2
4
3s 3p
The number of dots equals the number of s
and p electrons in the atom’s outermost shell.
20
Lewis Structures of the first 20 elements.
21
The chemistry of many elements,
especially the representative ones, is to
attain the same outer electron structure
as one of the noble gases.
22
With the exception of helium, this
structure consists of eight electrons in
the outermost energy level.
23
The Covalent Bond:
Sharing Electrons
24
A covalent bond consists of a pair of
electrons shared between two atoms.
In the millions of chemical compounds
that exist, the covalent bond is the
predominant chemical bond.
25
Substances which covalently bond exist
as molecules.
Carbon dioxide bonds covalently.
It exists as individually bonded
covalent molecules containing
one carbon and two oxygen
atoms.
26
The term molecule is not used when
referring to ionic substances.
Sodium chloride bonds ionically.
It consists of a large aggregate of
positive and negative ions. No
molecules of NaCl exist.
27
Covalent bonding in the hydrogen molecule
Two 1s orbitals from each of
two hydrogen atoms overlap.
Each 1s orbital contains 1
The two nuclei are
electron.
shielded from each
other by the electron
pair. This allows the
two nuclei to draw
close together.
The most likely
The
orbital
of the
region
to find
the
electrons
includes
two electrons
is
both
hydrogen
between
the two
nuclei.
nuclei.
28
Covalent bonding with equal sharing of
electrons occurs in diatomic molecules
formed from one element.
hydrogen
chlorine
iodine
nitrogen
A dash may replace a pair of dots.
H-H
29
Electronegativity
Linus Pauling
30
electronegativity The relative attraction
that an atom has for a pair of shared
electrons in a covalent bond.
31
• If the two atoms that constitute a
covalent bond are identical then there
is equal sharing of electrons.
• This is called nonpolar covalent
bonding.
• Ionic bonding and nonpolar covalent
bonding represent two extremes.
32
• If the two atoms that constitute a
covalent bond are not identical then
there is unequal sharing of electrons.
• This is called polar covalent bonding.
• One atom assumes a partial positive
charge and the other atom assumes a
partial negative charge.
– This charge difference is a result of the
unequal attractions the atoms have for
their shared electron pair.
33
Polar and Non-Polar
34
Partial positivePartial
charge
negative charge
on hydrogen. on chlorine.
Polar Covalent Bonding in HCl
+
:
:
H Cl
-
Chlorine
hasthat
a greater
attraction
forelement
the
Shared
Theof
shared
electron
electron
pair. pair
The attractive
force
an atom
an
has
shared electron pair than is
hydrogen.
to chlorine than
for shared electrons in a molecule closer
or a polyatomic
ion
to hydrogen.
35
is known as its electronegativity.
A scale of relative electronegativities
was developed by Linus Pauling.
36
Electronegativity generally
decreases increases
down a left
group
to right
for
representative
across
a periodelements
.
.
37
A dipole is a molecule that is
electrically asymmetrical, causing it to
be oppositely charged at two points.
A dipole can be written as
+
38
An arrow can be used to indicate a dipole.
The arrow points to the negative end of the
dipole.
Molecules of HCl, HBr and H2O are polar .
O
H
Cl
H
Br
H
H
39
A molecule containing different kinds of
atoms may or may not be polar depending
on its shape.
The carbon dioxide molecule is nonpolar
because its carbon-oxygen dipoles cancel
each other by acting in opposite directions.
40
Lewis Structures of
Compounds
41
In writing Lewis structures, the most
important consideration for forming a
stable compound is that the atoms attain
a noble gas configuration.
42
• The most difficult part of writing
Lewis structures is determining the
arrangement of the atoms in a molecule
or an ion.
• In simple molecules with more than
two atoms, one atom will be the central
atom surrounded by the other atoms.
43
Cl2O has two possible arrangements.
The two chlorines can be bonded to each other.
Cl-Cl-O
The two chlorines can be bonded to oxygen.
Cl-O-Cl
Usually the single atom will be the central atom.
44
Practice Writing
Lewis Structures
45
Valence Electrons of Group A Elements
Atom
Group
Valence Electrons
Cl
VIIA
7
H
IA
1
C
IVA
4
N
VA
5
S
VIA
6
P
VA
5
I
VIIA
7
46
3-Dimensional Shapes
Linear
180
Bent
105
Trigonal
Planar
120
Tetrahedral
109.5
Trigonal
Pyramidal
107
47
Covalent Bonding Structures
Molecular
Formula
Structure
Name
Bond
Angle
Polar or
Non-polar
H2O
Bent
105
Polar
CO2
Linear
180
Non-Polar
PH3
Trigonal
Pyramidal
107
Polar
NO3–
Trigonal
Planar
120
Non-Polar
Tetrahedral
109.5
CH4
Lewis “dot”
Structure
3-D
Structure
Non-Polar
48
“Define ‘resonance’? Sure, that’s where you live.”
49
The Ionic Bond: Transfer of
Electrons From One Atom
to Another
50
After sodium loses its 3s electron it has attained the
same electronic structure as neon.
51
After chlorine gains a 3p electron it has attained the
same electronic structure as argon.
52
Formation of NaCl
53
The
3s electron
of sodium
transfersion
to (Cl
the-)half-filled
3p
A
sodium
ion (Na+)
and a chloride
are formed.
orbital
of chlorine.
The
force
holding Na+ and Cl- together is an ionic bond.
Lewis representation of sodium chloride formation.
54
Formation of MgCl2
55
2+
2+) and two
Two
A
The
magnesium
3s
forces
electrons
holding
ionof(Mg
Mg
magnesium
two
transfer
chloride
Cl- together
toions
the are
half-filled
(Cl-ionic
) are
3p
formed.
bonds.
orbitals of two chlorine atoms.
56
In NaCl
the crystal
is made
each
upsodium
of cubicion
crystals.
is surrounded by six
chloride ions.
57
In the crystal each chloride ion is surrounded by six
sodium ions.
58
The ratio of Na+ to Cl- is 1:1
There is no molecule of NaCl
59
• Metals usually have one, two or three
electrons in their outer shells.
• When a metal reacts it:
– usually loses one two or three electrons
– attains the electron structure of a noble
gas
– becomes a positive ion.
• The positive ion formed by the loss of
electrons is much smaller than the
metal atom.
60
• Nonmetals are usually only a few
electrons short of having a noble gas
structure.
• When a nonmetal reacts it:
– usually gains one two or three electrons
– attains the electron structure of a noble
gas
– becomes a negative ion.
• The negative ion formed by the gain of
electrons is much larger than the
nonmetal atom.
61
Predicting Formulas of
Ionic Compounds
62
In almost all stable chemical compounds of
representative elements, each atom attains a
noble gas electron configuration.
63
Ions are always formed by adding or
removing electrons from an atom.
64
Most often ions are formed when metals
combine with nonmetals.
•Metals will lose electrons to attain
a noble gas configuration.
•Nonmetals will gain electrons to attain
a noble gas configuration.
65
The charge on an ion can be predicted
from its position in the periodic table.
66
elements
elements of
elements
elements
of of of
elements of Group
VIA have
a
Group IIA have Group
a Group
VAGroup
have
VIIA
a have
a
IA have a +1 charge
-2 charge
+2 charge
-3 charge
-1 charge
67
Writing Formulas From
Names of Compounds
68
A chemical compound must have a
net charge of zero.
69
If the compound contains ions,
then the charges on all of
the ions must add to zero.
70
Write the formula of calcium chloride.
Step 1. Write down the formulas of the ions.
Ca2+ ClStep 2. Combine the smallest numbers of Ca2+
- so that the sum of the charges
and
Cl
The cation
Theisanion is
equals
written
written
first. zero.
second.
(Ca2+) + 2(Cl-) = 0
(2+) + 2(1-) = 0 The lowest
common multiple
The correct formula is CaCl2
of +2 and –1 71is 2
Write the formula of barium phosphide.
Step 1. Write down the formulas of the ions.
Ba2+ P3Step 2. Combine the smallest numbers of Ba2+
3- so that the sum of the charges
and
P
The cation
The anion
is
is
equals
zero.
written
written
first. second.
3(Ba2+) + 2(P3-) = 0
3(2+) + 2(3-) = 0 The lowest
common multiple
The correct formula is Ba3P2
of +2 and –3 72is 6
Write the formula of magnesium oxide.
Step 1. Write down the formulas of the ions.
Mg2+ O2Step 2. Combine the smallest numbers of Mg2+
and O2- so that the sum of the charges
equals zero.
(Mg2+) + (O2-) = 0
(2+) + (2-) = 0The lowest
common multiple
The correct formula is MgO
of +2 and –2 73is 1
Write the Formula of Sodium Peroxide
1
Na O
2
2
gives
Na2O2
or
NaO
74
Write the Formula of Sodium Peroxide
Don’t mess with the subscripts of
polyatomic ions!!
1
Na O
2
2
gives
Na2O2
not
NaO
NaO does not contain the peroxide anion
75
Combine to Give Compounds (Do Not Name!)
ions
Br–
O–2
NO3–
PO4–3
CO3–2
NH4+
NH4Br
(NH4)2O
NH4NO3 (NH4)3PO4 (NH4)2CO3
Sn+2
SnBr2
SnO
Sn(NO3)2 Sn3(PO4)2 SnCO3
Al+3
AlBr3
Al2O3
Al(NO3)3
AlPO4
Al2(CO3)3
H+
HBr
H 2O
HNO3
H3PO4
H2CO3
76
77