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Chemical Bonds
Chapter 22
Chapter 22, Section 1
Compounds & Chemical Formulas
• Compounds: two or more elements that are chemically bonded
together
Compounds & Chemical Formulas
• Chemical formula: tells what elements a compound contains and
the exact number of atoms of each element in a compound
▫ A combination of chemical symbols and subscripts (numbers
written below the symbols)
▫ Each subscript indicates the exact number of atoms of the
element it comes directly after
▫ If there are no subscripts, the compound contains only one atom
of that element
Compounds & Chemical Formulas
• Ex: H2O has 2 atoms of H and 1 atom of O
• PRACTICE - How many atoms of each element are in the
following compounds:
1. SiO2
2. C12H22O11
3. Mg(OH)2
4. CH3COOH
Atomic Stability
• All elements want to reach stability
• An atom is considered chemically stable when its outermost
energy level has the maximum number of electrons
▫ Ex: hydrogen and helium are stable with two electrons
▫ Ex: the outer energy levels of all other elements are stable when they
contain eight electrons
Atomic Stability
• Valence electrons: electrons in the outermost energy level of
an atom
• Groups 1A, 2A, 3A, 4A, 5A, 6A, 7A, and 8A indicate the number
of valence electrons in the Group # before the “A”
▫ Ex: Elements in Group 3A (Group 13) have 3 valence electrons
• Recall that the Noble Gases all have full outer energy levels
▫ This indicates that elements in Group 8A (Group 18) are stable
Electron Dot Structures of Some Group A Elements
Group
Period
1
2
3
4
1A
2A
3A
4A
5A
6A
7A
8A
Bonding
• Atoms with partially stable outer energy
levels can lose, gain, or share electrons to
obtain stable outer energy levels
• They do this by combining with other
atoms that also have partially complete
outer energy levels
▫ Ex: sodium, Na bonding with chlorine, Cl
 Na+ + Cl-  NaCl
Octet Rule
• Octet rule: when forming compounds, atoms want to fill their
outer energy levels with a full set of valence electrons
• An octet is a set of 8
• Atoms want to look like the Noble Gases
Ions
• Ion: a charged particle because it has either more or fewer electrons
than protons
• Cation: positively charged ion
▫ Formed when an atom loses one or more valence electrons
▫ Typically formed by metals
▫ Ex: sodium ion  Na+
• Anion: negatively charged ion
▫ Formed when an atom gains one or more valence electrons
▫ Typically formed by nonmetals
▫ Ex: chloride ion  Cl-
Chapter 22, Section 2 – Part 1
Chemical Bonds & Ionic Compounds
• Chemical bond: the force that holds atoms together in a
compound
• Compounds have different physical and chemical properties
from those of the atoms that make up the compound
• Ionic compound: a compound made of positively charged ions
and negatively charged ions
▫ Ionic compounds are neutral because the total positive charge of the
cations equals the total negative charge of the anions
Chemical Bonds & Ionic Compounds
• Recall that all atoms want to have a full set of valence
electrons like the Noble Gases
• Ex: when sodium and chlorine react to form a compound, the
sodium atom transfers its one valence electron to the chlorine
atom that is lacking one valence electron to have a full set
▫ Sodium and chlorine can combine in a 1:1 ratio so that both achieve
stable octets!
Valence Electrons and Ions
• Metals in Group 1A tend to lose their 1 valence electron
• Metals in Group 2A tend to lose their 2 valence electrons
• Metals and Metalloids in Group 3A tend to lose their 3 valence
electrons
• Elements in Group 4A do not tend to form ions
• Nonmetals in Group 5A tend to gain 3 valence electrons
• Nonmetals in Group 6A tend to gain 2 valence electrons
• Nonmetals in Group 7A tend to gain 1 valence electron
Modeling Ionic Compounds
• Recall that a chemical formula shows the numbers of
atoms of each element in a compound using
subscripts
• In order to determine how many atoms of each
element are required to form a bond so that the
compound is electrically neutral, we can use electron
dot diagrams
• Ex: When calcium and
chlorine combine
▫ For every 1 calcium
atom, 2 chlorine atoms
are required to
“balance” the charges
and make the compound
neutral
▫ The chemical formula is
thus: CaCl2
Modeling Ionic Compounds
• Use electron dot diagrams to predict the formulas of
the ionic compounds formed from the following
elements:
1. Potassium and oxygen
2. Aluminum and bromine
3. Magnesium and nitrogen
Ionic Bonds
• Ionic bond: the force of attraction between the
opposite charges of ions in an ionic compound
• Ionic bonds are formed when metals bond with
nonmetals
• Ionic bonds involve the TRANSFER of electrons to form
compounds
Ionic Bonds
• Ex: in the example before, calcium “loses” or “transfers”
its 2 valence electrons to two chlorine atoms (1 to each)
▫ The calcium atom thus attains a charge of +2
▫ Each chlorine atom thus attains a charge of -1
▫ The result is a compound with a net charge of zero!
• Because of the new charges attained by each atom, the
ions stick together by electric forces
▫ Positives are attracted to negatives
Cations
• Lose electrons
• Formed by metals
• Groups 1A, 2A, and 3A form charges equal to their group
number
Group 1A: Lose 1 electron to form 1+ ions
H+
Li+
Na+
K+
Group 2A: Lose 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+
Group 3A: Lose 3 electrons to
form 3+ ions
B3+
Al3+
Ga3+
Naming Cations
• The names of Groups 1A, 2A, and 3A are the same as the name
of the metal, followed by the word ion or cation
Ion
Na+
Ca2+
Al3+
Name
Sodium ion
Calcium ion
Aluminum ion
Anions
• Gain electrons
• Formed by nonmetals
• Groups 5A, 6A, and 7A form charges equal to:
18 – group number
Ex: elements in Group 7A
17 – 18 = -1
Elements in Group 7A form anions with a -1 charge!
Group 14 (4A): there aren’t any
common ions for this
group.
N3P3As3-
Nitride
Phosphide
Arsenide
Group 15 (5A): Gain 3
electrons to form
3- ions
O2-
Oxide
S2-
Sulfide
Se2-
Selenide
Group 16: Gain 2 electrons
to form 2- ions
F- Fluoride
Br- Bromide
Cl- Chloride
I- Iodide
Group 17: Gain 1
electron to form 1ions
Group 18/ 8A: Stable
Noble gases - do
not form ions!
Naming Anions
• Anion names start with the root of the element name
and end in –ide
Ion
N3P3O2-
Name
Nitride
Phosphide
Oxide
Practice Naming Ions:
• Write the symbol for the ion formed, classify it as a cation or
anion, and name the ion for each of the following:
1. Potassium

K+ , cation, potassium ion
2. Iodine

I-, anion, iodide
3. Sulfur

S2-, anion, sulfide
Naming Ions formed by Transition Metals
• Ions formed by transition metals follow a different rule:
we will not focus on these yet…
Naming Ionic Compounds
• Recall that an ionic compound is made up of a metal cation
and a nonmetal anion
• Always name the cation first, followed by the anion
1. Cation = name of the element
 Ex: Ca2+ = calcium (ion)
2. Anion = root of element name + -ide
 Ex: Cl- = chloride
 So the name of the ionic compound formed is calcium chloride
Practice Naming Ionic Compounds
1. CaBr2
▫ Calcium bromide
2. The ionic compound formed by combining potassium
and oxygen
▫ Potassium oxide
3. Be3N2
▫ Beryllium nitride
Writing Chemical Formulas for Ionic Compounds
Steps to writing chemical formulas for ionic compounds:
1. Write the chemical symbols with charges for both ions
▫
Be sure to write the cation’s symbol first!
2. Add subscripts when necessary to balance the charges by
using the CRISSCROSS method
Writing Chemical Formulas for Ionic Compounds
Example: Aluminum sulfide
1. Write the formulas for the cation and anion, including
charges
▫
Al3+
S2-
2. Use subscripts to balance the charges using the CRISSCROSS
method, if necessary
▫
▫
Al3+2 S2-3
Correct formula: Al2S3
Practice Writing Chemical Formulas for Ionic Compounds
1. Sodium nitride
▫ Na3N
2. Magnesium oxide
▫ MgO
3. Gallium bromide
▫ GaBr3
Binary Ionic Compounds
• Binary compound: a compound made up of two
elements
▫ Ex: NaCl, MgBr2
• In order to write the correct chemical formula, you
need to know the correct charges of each ion involved
in the bond
▫ CRISSCROSS method
Oxidation Numbers
• Recall that all elements in the same group have the same
number of valence electrons (electrons in their outermost
energy levels)
• Thus, they must gain or lose the same number of electrons to
achieve a full outer energy level in a bond
• Oxidation number: the charge of an ion
• For ionic compounds, the oxidation number is the same as the
charge of the ion
▫ Ex: the sodium ion (Na+) has a charge of 1+ and an oxidation number of
1+
Transition Metals & Other Metals
• Transition metals are the elements in the main body of the
periodic table
• These are known as Type II metals
• These metals tend to form cations with different oxidation
numbers
• The charges of the ions formed are determined by the number
of electrons lost
• Ex: iron (Fe) can lose 2 electrons (Fe2+) or 3 electrons (Fe3+)
Transition Metals & Other
Metals
• Common metals (Type II
metals) that have more than
one oxidation number:
Name
Oxidation
Number
Copper (I)
1+
Copper (II)
2+
Iron (II)
2+
Iron (III)
3+
Chromium (II)
2+
Chromium (III)
3+
Lead (II)
2+
Lead (IV)
4+
Naming Compounds of Metals with Multiple
Oxidation Numbers
• Place a Roman Numeral in parentheses
after the name of the element to
indicate the charge (oxidation number)
▫
Ex: Fe2+ is iron (II)
• Write the chemical formula in the same
way you do for other ionic compounds
by crisscrossing the charges
▫ Ex: Fe3+ bonded with O2 Fe2O3
Practice Naming!
Name the ionic compounds formed when combining the
following ions:
1. Cu2+ and Br▫ Copper (II) bromide
2. Pb4+ and S2▫ Lead (IV) sulfide
3. Cr3+ and N3▫ Chromium (III) nitride
Practice Writing Chemical Formulas!
Write the correct chemical formula for the ionic
compound formed by combining the following ions:
1. Copper (I) and selenide
▫ Cu2Se
2. Iron (III) and oxide
▫ Fe2O3
3. Lead (II) and phosphide
▫ Pb3P2
Polyatomic Ions
• Not all ionic compounds are binary (include only two
elements)
• Some ionic compounds contain more than two elements
because they include polyatomic ions
• Polyatomic ion: a group of atoms that are chemically bonded
together and have a positive or negative charge together
▫ Ex: C2H3O2- = acetate
▫ Ex: NH4+ = ammonium
• Look at your table of common polyatomic ions to help
Practice Identifying Polyatomic Ions
Write the correct chemical formula and charge for each of the
following polyatomic ions:
1. Hydroxide
▫
OH-
2. Phosphate
▫
PO43-
3. Nitrate
▫
NO3-
Naming Ionic Compounds with Polyatomic Ions
1. Write the name of the cation (positive ion) first
2. Then write the name of the anion (negative ion)
• Ex: K2SO4
▫ Potassium sulfate
Writing Chemical Formulas for Ionic Compounds
with Polyatomic Ions
• Follow the rules for naming binary compounds by
crisscrossing charges
▫ List cation first, followed by anion
• Use parentheses around each polyatomic ion before
adding the appropriate subscript
• Ex: barium chlorate
▫ Ba(ClO3)2
Practice Naming!
Write the correct names for the following ionic
compounds:
1. Sr(OH)2
▫ Strontium hydroxide
2. CaSO4
▫ Calcium sulfate
Practice Writing Chemical Formulas!
Write the correct chemical formulas for the following
ionic compounds:
1. Beryllium nitrate
▫ Be(NO3)2
2. Sodium carbonate
▫ Na2CO3
Ionic Bonding Quiz Important Information
Name
Chemical Formula
Type II Metals
Chlorate
ClO3-
Titanium
Chlorite
ClO2-
Nickel
Cyanide
CN-
Lead
Acetate
CH3COO-
Iron
Phosphate
PO43-
Sulfate
SO42-
Carbonate
CO3
2-
Cobalt
Chromium
Gold
Covalent vs. Ionic Bonds
• Recall that an ionic bond involves the transfer of
valence electrons between a metal and a nonmetal
• In a covalent bond, valence electrons are shared
between two nonmetals
Covalent Bonds
• Atoms involved in a covalent bond still want to
achieve a full outer energy level
• By sharing valence electrons, all atoms involved
can count the electrons toward achieving a
complete octet (set of 8 valence electrons)
Molecules
• Molecule: a neutral group of atoms joined together by
covalent bonds
• Molecular compound: a compound composed of
molecules
• Diatomic molecule: a molecule that contains two
identical atoms
▫ Br2, I2, N2, Cl2, H2, O2, F2
Single Covalent Bonds
• Single covalent bond: two atoms held together by sharing one
pair of (2) electrons
• Ex: All halogens have seven valence electrons
▫ By sharing valence electrons, two of the same halogens end up with a
complete octet
Single Covalent Bonds
• Ex: Similar to the halogens, hydrogen exists as a diatomic
molecule, sharing its 1 pair of electrons
Double Covalent Bonds
• Double covalent bond: a bond that involves two shared pairs
of electrons
▫ Ex: Carbon dioxide (CO2)
Triple Covalent Bonds
• Triple covalent bond: a bond formed by sharing three pairs of
electrons
▫ Ex: Nitrogen (N2)
Steps to Drawing Lewis Structures
1. Count the total number of valence electrons for the given
molecule by adding each individual element’s valence
electrons
2. Draw the “skeleton” of atoms in the molecule
▫ Hydrogens are always around a central atom
3. Join all the atoms using only single bonds
Steps to Drawing Lewis Structures
4. Add the electrons as dots to fulfill the octet rule around each
element in the compound
 Remember that hydrogen requires only 2 electrons
5. If the total comes out to what you calculated in step 1,
you’re done! If not, you may need to add double or triple
bonds as necessary
 Remember that each atom still needs to satisfy the octet
rule!
Exceptions to the Octet Rule
• The octet rule cannot be satisfied in molecules
whose total number of valence electrons is an
odd number
▫ Ex: NO2 contains 17 total valence electrons – each
oxygen contributes 6 and the nitrogen contributes 5
• There are also molecules in which an atom has
less, or more, than a complete octet of valence
electrons
▫ Ex: sulfur hexafluoride (SF6)
Practice Drawing Lewis Structures
1.
2.
3.
4.
SCl2
CH4
NH3
H2O
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:
1. The size of the atom
2. The charge of the nucleus
3. The total number of electrons in the atom
Unequal Sharing
• Electronegativity: the ability of an atom to attract
electrons in a chemical bond
▫ High electronegativity means that it pulls the electrons
toward it
▫ Electronegativity decreases as you move down a group, but
increases as you move from left to right across a period
▫ Fluorine is the most electronegative atom on the periodic
table
Bond Polarity
• The unequal sharing of electrons can give one atom a slight
negative charge (δ-) and another atom a slight positive charge
(δ+)
• The atom holding the electrons more closely is the more
electronegative atom and has a slight negative charge (δ-)
Bond Polarity
• Polar molecule: a molecule with a slightly positive end and a
slightly negative end, although the overall molecule is neutral
▫ Ex: the O-H bonds in a water molecule
Bond Polarity
• Nonpolar molecule: a molecule in which electrons are shared
equally
▫ Ex: diatomic molecules (identical atoms), or symmetric molecules
Difference 0 or less
NonPolar
Difference greater than 0 but less than 2 Polar
Difference greater than 2
Ionic Bond
Ionic vs. Covalent Bonds
• Covalent bonds are strong, but the attractions between
individual molecules are weak
• Ionic bonds are strong because of the strong attraction
between positive and negative ions
▫ Therefore, ionic compounds have high melting and boiling
points compared to covalent compounds
Hydrates
• Hydrate: a compound that has water chemically
attached to its ions and written into its chemical
formula
• To write the chemical formula for a hydrate, write the
formula for the compound, then place a dot followed
by the number of water molecules
▫ Ex: calcium sulfate dihydrate is CaSO4•2H2O
Binary Molecular Compounds
• Binary molecular
compounds are made up of
two nonmetals
• Prefixes are used to
describe the number of
atoms of an element that
are present in each
molecular compound
Naming Molecular Compounds
• To name a binary molecular compound:
1. Write the names of the elements in the order listed in
the chemical formula
2. Use prefixes appropriately to indicate the number of
each kind of atom
 Omit the prefix mono- for the first element in the
compound
3. End the name of the second element with the suffix –ide
Practice Naming Molecular Compounds
1. CO
▫ Carbon monoxide
2. Cl2O8
▫ Dichlorine octoxide
3. PCl3
▫ Phosphorus trichloride
Writing Chemical Formulas for Binary Molecular
Compounds
• To write the correct chemical formula for a binary molecular
compound:
1. Use the prefixes to tell you the subscript of each element
in the formula
2. Write the correct symbols for the elements with the
appropriate subscripts
Practice Writing Chemical Formulas for Molecular
Compounds
1. Tetraphosphorous trisulfide
▫ P4S3
2. Iodine heptafluoride
▫ IF7
Naming Hydrates
• Use the same prefixes to name hydrates
• To name a hydrate:
1. Name the ionic compound as you normally would
2. Write the prefix to indicate the number of water molecules
in the compound
3. Follow the prefix with the word hydrate
Writing Chemical Formulas for Hydrates
• To write the formula for a hydrate:
1. Write the chemical formula for the ionic compound
2. Draw a dot
3. Write the number of waters that are in the compound
Practice Naming Hydrates
Write the name of each of the following hydrates:
1. CoCl2•6H2O
▫ Cobalt (II) chloride hexahydrate
2. CuSO4•5H2O
▫ Copper (II) sulfate pentahydrate
Practice Writing Formulas for Hydrates
Write the chemical formula for the following hydrates:
1. Iron (III) sulfate nonahydrate
▫ Fe2(SO4)3•9H2O
2. Sodium carbonate decahydrate
▫ Na2CO3•10H2O