Bonding
A chemical bond is a force that
holds two or more atoms together.
Compound – two or more
elements chemically combined
by gaining, losing, or sharing
electrons.
Molecule – a particle made of 2
or more atoms covalently
bonded together.
Electrons
closest to the nucleus -least
amount of energy and lowest
energy level.
farthest from the nucleus have
the greatest amount of energy
and are in the highest energy
level.
Outermost electrons are easily
attracted to the positively
charged nucleus of other
atoms. This forms a chemical
bond. These outermost
electrons are called
valence electrons.
Using the Periodic table to determine
the number of valence electrons.
Groups 1,2, 13 – 18 have the same
number of valence electrons as the
ones digit of the group number,except
helium which only has 2.
Groups 3- 12 vary in numbers of valence
electrons.
Lewis Electron Dot Diagram p. 271
• Lewis Electron Dot diagram is a
model that represents valence electrons
in an atom as dots around the element’s
chemical symbol.
(See handout for steps to do this.)
• Elements with unpaired dot are reactive
or chemically unstable.
• Atoms will bond to other atoms to
become chemically stable.
Isotopes
Isotopes : are atoms of the same
element with different numbers of
neutrons.
Example: Isotopes of the element, carbon.
Carbon - 12
6 protons 6 neutrons
Carbon - 14
6 protons 8 neutrons
The same element, same number of protons,
different number of neutrons.
(If proton number changes, element changes)
Covalent Bond
(nonmetal to nonmetal)
Covalent Bonding – A type of
bonding in which two or more
atoms share one or more pairs of
valance electrons.
When covalent bonding occurs,
molecules are formed.
Covalent Bond
• Single covalent bond – atoms share one
pair of valence electrons.
• Double covalent bond – atoms share two
pairs of valence electrons.
• Triple covalent bond – atoms share three
pairs of valence electrons.
Covalent bond examples
• The more valence electrons that two
atoms share, the stronger the covalent
bond is between the atoms.
• Examples :
H20 CO2 N2 C6H12O6
H2
Valence number – The number
of electrons an atom needs to
gain, lose, or share to become
stable.
(Gain electrons – negative
valence)
(Lose electrons – positive
valence.)
Chemical formulas – a group of
chemical symbols and numbers
that represent the number of atoms
of each element that make up a
compound . ( H2O)
In a chemical formula, the
element with the positive
charge is always written first.
Subscript – the number that
shows how many atoms of an
element are in the compound.
Examples
H20, CO2
The number 2 is the subscript in each.
How to Write Formulas
• 1. Write the symbol correctly for both
elements.
• 2. Write the valence number above each
symbol. Make sure you have a positive
first, and then a negative.
• 3. Criss-cross valences to make a
subscript.
• 4. NEVER write 1 as a subscript. (It is
understood to be 1 if there is no number.)
• 5. Reduce to lowest terms.
Examples
• Sodium Chloride
• Copper (I) Sulfide
Aluminum Oxide
Iron (III) Iodide
1. Calcium Sulfide
2. Copper (I)Iodide
3. Potassium Chloride 4. Iron(III) Bromide
5. Sodium Nitride
6. Antimony (V) Oxide
Problem check
1. Calcium Sulfide
2. Copper (I)Iodide
CaS
CuI
3. Potassium Chloride 4. Iron(III) Bromide
KCl
FeBr 3
5. Sodium Nitride
Na3N
6. Antimony (V) Oxide
Sb2O5
Practice Problems
• Aluminum Bromide
• Calcium Iodide
Magnesium Nitride
Potassium Sulfide
• Copper (I) Chloride Barium Fluoride
• Hydrogen Oxide
Lead Bromide
Practice Problems
Aluminum Bromide Magnesium Nitride
AlBr3
Mg3N2
Calcium Iodide
Potassium Sulfide
CaI2
K2S
Copper (I) Chloride Barium Fluoride
CuCl
Ba F2
Hydrogen Oxide
Lead Bromide
H 2O
PbBr2
Practice Problems
1. Lead Iodide
2. Barium Fluoride
3. Tin (II) Nitride
4. Lithium Chloride
5. Sodium Bromide
6. Iron (II) Oxide
7. Potassium Fluoride 8. Silver Bromide
9. Calcium Chloride
10. Copper (I) Iodide
11. Aluminum Nitride
12. Cobalt (II) Sulfide
Practice Problems
•
•
•
•
•
•
1. PbI2
3. Sn3N2
5. NaBr
7. KF
9. CaCl2
11. AlN
2. BaF2
4. LiCl
6. FeO
8. AgBr
10. CuI
12. CoS
Binary Compound – a
compound composed of ONLY
2 elements.
Polyatomic Ion- a group of
positively or negatively charged
covalently bonded atoms.
Calcium Phosphate
Tin (IV) Chromate
Ammonium Oxide
Magnesium Sulfate
Homework Check
1. K2O
6. Ca3(PO4)2
2. NaC2H3O2
7. FeBr2
3. Ag2SO4
8. HClO3
4. ZnCO3
9. MgSO3
5. CrS
10. (NH4)3PO4
Homework Check
Cd CrO4
Al2O3
FePO4
AgNO3
NaClO2
ZnI2
Cu(ClO3)2
PbC2O4
KClO3
SnS2
SbCl5
Li3PO4
Naming compounds – change
the ending of the second
element to – “ide”.
(Do not change the names of
polyatomic ions.)
Check to see if Roman numeral
is needed.
Ex: NaCl - Sodium Chloride
Fe2O3 – Iron (III) Oxide
Examples
• Zn3N2
______________________
• K2SO3
______________________
• CuO
______________________
• Ba(C2H3O2)2 _____________________
Ion – an atom or group of atoms that
has become electrically charged. It is a
charged particle.
When an atom loses an electron, it
loses a negative charge and becomes
a positive ion.
When an atom gains an electron, it
gains a negative charge and becomes
a negative ion.
Ionic Bond
(Metal to Nonmetal)
1. Ionic Bond – the attraction
between two oppositely charged
ions.
An ionic bond is formed when
atoms gain or lose electrons.
Ionic bonds form compounds.
Ionic bond
Homework Check
1.Ammonium Phosphate
2. Iron (ll) Bromide
3. Hydrogen Chlorate
4. Sodium Acetate
5. Silver Sulfate
6. Chromium (lll) Sulfide
7. Potassium Oxide
8. Tin (ll) Oxide
9. Hydrogen Oxide
10. Antimony (III) Phosphate
Predicting Reactions between
Elements
To form an ionic bond, a metal bonds to
a non-metal.
Metals to the left of stair-step line except H.
Non-metals to the right of stair-step line,
except H
Metalloids along stair-step except Al which
is a metal.
Predicting ionic bonds
• Metals react with non-metals to
form ionic compounds.
• Will Na react with Br? _____
• Will Ca react with Ag _____
• Will S react with Cl? _______
Metallic Bonds
(metal to metal)
• Metallic Bond – Bond formed when
many metal atoms share their pooled
valence electrons.
• Valence electrons are not bonded to one
atom but positive ions are surrounded by
a “sea of electrons” Ex: Aluminum
Properties of compounds.
• See handout and p. 288 in textbook.
Homework check
1.Iron (ll) Hydroxide
2. Aluminum Perchlorate
3. Chromium (ll) Acetate
4. Tin (lV) Chloride
5. Lead Chromate
6. Copper (ll) Sulfide
7. Potassium Hypochlorite
8. Mercury (ll) Dichromate
9. Calcium Hydroxide
10. Tin (ll) Sulfite
Homework check
11. Zinc Phosphate
12. Barium Fluoride
13. Sodium Nitrate
14. Cadmium Carbonate
15. Mercury (ll) Peroxide
16. Iron (lll) Sulfite
17. Antimony (V) Oxide
18. Copper(l) Permanganate
19. Chromium (lll) Oxalate
20. Ammonium Sulfate
Homework check
Chlorine - nonmetal Aluminum - metal
Boron - metalloid
Calcium - metal
Zinc - metal
Hydrogen - nonmetal
Neon - nonmetal
Nickel - metal
Sodium - metal
Magnesium - metal
Silver - metal
Fluorine - nonmetal
Silicon - metalloid
Potassium - metal
Helium - nonmetal
Uranium - metal
Bromine - nonmetal Carbon - nonmetal
Francium - metal
Argon - nonmetal
Homework Check
Sodium & Calcium - no -two metals metallic
Hydrogen & Oxygen -no two nonmetal covalent
Potassium & Bromine yes -metal +nonmetal
Francium & Fluorine - yes metal + nonmetal
Helium & Carbon - no two nonmetals covalent
Magnesium & Chlorine - yes metal+ nonmetal
Neon & Argon - no - two nonmetals covalent
Nickel & Copper - no two metals metallic
Phosphorus & Oxygen - no two nonmetals cov
Calcium& Uranium - no two metals metallic
Homework Check
1.Ammonium Phosphate 11. Manganese(ll) Oxide
2. Iron (ll) Bromide
12. Tin (lV) Iodide
3. Hydrogen Chlorate
13. Calcium Oxide
4. Sodium Acetate
14. Potassium Chloride
5. Silver Sulfate
15. Copper (l) Hydroxide
6. Chromium (lll) Sulfide
16. Ammonium Sulfide
7. Potassium Oxide
17. Aluminum Carbonate
8. Tin (ll) Oxide
18. Hydrogen Sulfite
9. Copper (ll) Phosphate
19. Iron (lll) Sulfate
10. Zinc Sulfate
20. Aluminum Nitride