Transcript Properties of Atoms and the Periodic Table

Chapter 22 – Chemical Bonds
Chapter preview
sections
1 Stability in Bonding
2 Types of Bonds
3 Writing Formulas and Naming Compounds
Elements Form Chemical Bonds
Just as skydivers link together to make a stable
formation, the atoms in elements can link together
with chemical bonds to form a compound. You will
read about how chemical bond form and learn how
to write chemical formulas and equations
Chemical Bonds
Section 1—Stability in
Bonding
Combined Elements



Some matter around you is in the
form of uncombined elements, such
as copper, sulfur and oxygen.
They can unite chemically to form a
compound, when conditions are right.
The green coating, on the statue of
Liberty is such a compound, called
copper sulfate.
New Properties


An observation you will make, is that the
compound formed when elements
combine, often has chemical and
physical properties that aren’t anything
like those of the individual elements.
Na
+
Cl

Sodium
+
Chlorine

Na+ ClSodium Chloride (Table Salt)
Formulas



The chemical symbols Na and Cl represent
the elements sodium and chlorine.
When written as NaCl, the symbols make up
a formula, or chemical shorthand, for the
compound sodium chloride.
A chemical formula tells what elements a
compound contains and the exact number of
atoms of each element in a unit of that
compound.
Learn Table 1 on page 689
Some Familiar Compounds
Familiar Name
Chemical Name
Formula
Sand
Silicon dioxide
Milk of magnesia
Magnesium hydroxide
Mg(OH) 2
Cane sugar
Sucrose
C12H22O11
Lime
Calcium oxide
Vinegar
Acetic Acid
Laughing gas
Dinitrogen oxide
Grain alcohol
Ethanol
Battery acid
Sulfuric acid
Stomach acid
Hydrochloric acid
SiO2
CaO
CH3COOH
N 2O
C2H5OH
H2SO4
HCl
Atomic Stability

The electric forces
between
electrons
and protons, which
are opposite, are the
forces that cause
compounds to form..

The Unique Noble
Gases can be
understood if you
look at an electron
dot diagram of them.
Chemical Stability


An atom is
chemically stable
when its
outermost energy
level has the
maximum number
of electrons.
Helium and
Hydrogen are
stable with two
electrons

All other elements
are stable when
they have eight
electrons.
He
Helium He

Stable
Unstable
Atomic components
•
•
When you look at the
elements in Groups
13 through 17, you
see that none of the
elements has a
stable energy level.
Each group contains
too few electrons for
a stable level of eight
electrons.
1
2
13

14 15 16 17
Unstable

18
Outer Levels – Getting Their Fill




Hydrogen does not have a
full outer energy level.
How does it or any other
such element, become
stable?
Atoms with partially stable
energy levels can lose,
gain, or share electrons.
They combine with other
atoms that also have
partially complete outer
energy levels.

Ea: Sodium will lose one
electron, while chlorine
gains
one
electron,
making the combination
stable.
Stability is Reached




Sodium had only one
electron in its outer energy
level which it lost when it
combined with chlorine to
for sodium chloride.
An atom that has lost or
gained an electron is called
an ion.
An ion is a charged particle
because it now has either
more or fewer electrons
than protons.
It is this electric force that
holds these compounds
together.


Water (H2O)
In water, Hydrogen and oxygen
each contribute one electron to
each hydrogen—oxygen bond.
The atoms share those electrons
instead of giving them up.
Types of Bonds
Section 2
Gain or Loss of Electrons


Atoms can loan
electrons to
another atom, so
they both become
stable.
An attractive force
occurs with this
and we call it a
chemical bond.
A Bond forms

This type of sharing
occurs between atoms
in Group 1 and Group
17. The atoms become
ions, with one being
positive and the other
negative. Ea: sodium
(Na) joins with chlorine
(Cl) to create sodium
chloride [Na]+ [Cl]- 
NaCl
The Ionic Bond


When ions attract in this way, a bond
is formed.
An ionic bond is the force of
attraction between opposite charges
of ions in an ionic compound.
Zero Net Charge


The result of this
bond is a neutral
compound.
The compound as a
whole is neutral
because the sum of
the charges is zero.
Magnesium
+
2 chlorine

Magnesium Chloride
Sharing Electrons
 Some
atoms are unlikely to lose
or gain electrons.
• Group 14 have 4 electrons in outer
shell.
• They become more stable by
sharing atoms , since removing one
makes the atom hold others more
tightly.
Single Covalent Bond


Bonds formed by sharing
are called covalent bonds.
A neutral particle that
results is called a molecule.
• A single covalent bond is
•
usually made up of two shared
electrons.
Water contains two single
covalent bonds.
Multiple Bonds

A covalent bond also
can contain more
than one pair of
shared electrons.
• An example is the bond
•
•
in nitrogen (N2)
Each atom contributes
three electrons.
Each pair of electrons
represents a bond.
Unequal Sharing


Electrons are not always shared equally
between atoms in a covalent bond.
The strength of attraction of each atom
to its electrons is related to size of atom.
• One example is found in a molecule of
•
hydrogen chloride (HCl).
Chlorine atoms have a stronger attraction
than Hydrogen.
Tug-of-War:
Partial positive
charge δ +
Partial negative
charge δ -
You might think of a covalent bond as the rope in a tug-of-war, and the
shared electrons as the knot in the center. Each atom in the molecule
attracts the electrons they share. Some atoms aren’t the same size.
Therefore these stronger atoms pull harder in their direction.
Chloroform
Hydrogen Fluoride
Polar or Nonpolar?

For molecules involved in this tug-ofwar, there is another consequence.
The atom holding the electrons more
closely will have a slightly negative
charge, and the other atom will be
slightly positive.
•
•
This type of molecule is called polar. A
polar molecule is one that has a slightly
positive end, and a slightly negative end.
Water is an example, and its polarity is
responsible for many of its unique
properties
Polar or Nonpolar?


Two atoms that are exactly alike can
share their electrons equally, forming a
nonpolar molecule.
This is true of molecules with identical
atoms, or symmetric atoms (CCl4)
Hydrogen molecule 
Nonpolar  Oxygen molecule
carbon tetrachloride
Nonpolar (symmetric)
Properties of Compounds

Recall: Atoms can form two types of
bonds, covalent and ionic.
• Sugar is a covalent compound
• Table salt is an ionic compound
• Their comparison can be seen in the table on
page 701 of you books.
Comparison of Covalent & Ionic Compounds
(page 701)
Covalent Compound
Ionic Compound
Electron Sharing
Electron Transfer
Lower
Higher
Poor
Good
State at Room
Temperature
Solid, liquid, or gas
Solid
Forces Between
Particles
Strong bonds
between atoms;
weak attraction
between molecules
Strong attraction
between positive
and negative ions
Bond Type
Melting & Boiling
Points
Electrical
Conductivity
Covalent & Ionic Properties


The difference in properties is due to
differences in attractive forces of the
bonds.
Covalent Compounds:
•
•
•
•
Covalent bonds between atoms are strong
Attraction between molecules is weak
Melting & boiling points are relatively low (sugar
melts at 185ᵒC.
They form soft solids, which have poor electrical
and thermal conductivity. (Ea: candles & propane
gas)
Covalent & Ionic Properties
(continued)

Ionic Compounds:
• Bonds between ions are relatively strong.
• They have high melting & boiling points – Salt
•
•
•
melts at 801ᵒC
Solids are hard and brittle.
When in liquid or aqueous state they readily
conduct electric current.
Ionic compounds are stable due to the
strength of attraction.
Writing Formulas and
Naming Compounds
Section 3
The numbers in columns 1,
2, 13, 14, 15, 16, 17 & 18; is
the most common oxidation
number if elements in that
group (family).
Binary Ionic Compounds
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
A binary compound
is
one
that
is
composed of two
elements.
Before you can write
a formula, you must
have all needed
information.

What you will need
to know:
1. Identify the symbols of the
cation (first part of the
name) and anion
2. Identify the charge for each
and place above the
symbol in parenthesis
3. Balance the positive and
negative charges
4. Write the formula placing
the subscripts right after
the symbol they go with.
Oxidation Numbers

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You will need to know which elements are involved
and what number of electrons they lose, gain, or
share to become stable.
Because all elements in a group have the same
number of electrons in their outer energy levels, they
must gain or lose the same number of electrons.
Metals always lose electrons and nonmetals always
gain electrons when they form ions.
The charge on the ion is known as the oxidation
number of the atom.
Oxidation Numbers & Periodic Table

The numbers with positive or negative signs are the
oxidation numbers for the elements in the column
below them.
Writing Formulas

When writing an ionic formula between a metal and a nonmetal follow
these 5 steps:
1. Write the symbols for the metal and the nonmetal.
2. Write the valences as superscripts above each symbol.
3. Drop the + and - sign.
4. Crisscross the valences so they become the subscript for the other element.
5. Reduce subscripts whenever possible. Only when both are divisible by a
number
greater than one.

Let's use these rules to figure out the chemical formula for our
compound between aluminum and oxygen.
Tables for Writing Names
Table 3 Special ions
Name
Oxidation
Number
Copper (I)
1+
Copper (II)
2+
Iron (II)
2+
Iron (III)
3+
Chromium (II)
2+
Chromium (III)
3+
Lead (II)
2+
Lean (IV)
4+
Table 4 Elements in
Binary Compounds
Element
Oxygen
Phosphorus
-ide
Name
oxide
phosphide
Nitrogen
nitride
Sulfur
sulfide
Writing Names

(See rules on p.706)
You can name a binary ionic compound
from its formula by using these rules:
1.
2.
3.
4.
Write the name of the positive ion.
Use Table 3 – Check to see if there could
possibly be more than one oxidation number.
Write the root name of the negative ion. (the
first part of the electrons name. Ea: Chlorine
becomes chlor.
Add the ending –ide to the root. Table 4 list
several elements and their –ide counterparts.
Compounds with Polyatomic Ions
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Not all ionic compounds are binary.
Baking soda has the formula: NaHCO3.
This is an example of a non-binary ionic
compound.
A polyatomic ion is a positive or negatively
charged, covalently bonded group of elements.
The polyatomic ion in Baking soda is the
bicarbonate or hydrogen carbonate ion HCO3.
Polyatomic ions

Writing Names

Several Polyatomic ions are
named in Table 5.
First write the name of the
positive ion.
Then use table 5 to find the
name of the polyatomic ion.
Ea: K2SO4 is Potassium sulfate,
Sr(OH)2 is Strontium hydroxide




(continued)
Table 5
Polyatomic ions
Charge
Name
Formula
1+
ammonium
NH4+
1-
acetate
chloride
hydroxide
nitrate
C2H3O2ClO3OHNO3-
2-
carbonate
sulfate
CO32SO42-
3-
phosphate
PO43-
Polyatomic ions

Writing formulas

Follow the rules for binary
compounds,
with
one
addition. When more than
one
polyatomic
ion
is
needed, write parentheses
around the ion before adding
the subscript.
Barium chlorate: Barium is
Ba2+ and chlorate is ClO31The formula is: Ba(ClO3)2


(continued)
Figure 21 Naming Complex
Components
How would a scientist write the
chemical formula for ammonium
phosphate?
To write the formula
answer the following questions:
1. What is the positive ion and its charge?
The positive ion is NH41+ and its charge is 1+
2. What is the negative ion and its charge?
The negative ion is PO43- and its charge is 3-
3. Balance the charges to make the compound
neutral
a) three NH41+ ions (+3) balance one PO43- (3-)
b) The charge of one ion (without the sign)
becomes the subscript of the other. Add
parentheses for subscripts greater than 1.
NH41+ PO43- gives (NH4)3 PO4
Compounds with Added Water

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
Some ionic compounds have water molecules as part of their
structure. These compounds are called hydrates
A hydrate is a compound that has water attached chemically to
its ions and written into its chemical formula.
Common Hydrates: Copper(II) sulfate pentahydrate,
CuSO4 · 5H2O. The name and formula indicate that there are
five water molecules per copper sulfate formula unit.
Hydrated water molecules are generally indicated in formulas
as shown using a dot to separate the water molecules from the
formula of the salt.
If copper(II) sulfate pentahydrate is heated, the bright blue
crystals of the hydrate are converted to a white, powdery,
anhydrous salt. Anhydrous – “without” water
Naming Binary Covalent Compounds




Covalent compounds are those formed between
elements that are nonmetals.
For example: Nitrogen and oxygen can form N2O,
NO, NO2, N2O5.
These would all be called nitrogen oxide.
Using Prefixes: Using Greek prefixes, these would
be called dinitrogen oxide, nitrogen oxide, nitrogen
dioxide, dinitrogen pentoxide.
Naming Binary Covalent Compounds
(continued)



Notice that the last vowel of
the prefix is dropped when
the second element begins
with a vowel.
Carbon monoxide is an
example.
These same prefixes are
used when naming the
hydrates
previously
discussed.
Table 6 Prefixes for
Covalent Compounds
Number of
Atoms
Prefix
1
mono
2
di-
3
tri-
4
tetra-
5
penta-
6
hexa-
7
hepta
8
octa