Bonding and Nomenclature
HTHS Chemistry
1
Ionic Bonding
Keeping Track of Electrons
Valence electrons are the electrons responsible for the chemical properties of atoms and are those
electrons in the outermost energy level (highest number). Valence electrons are found in the s and p
sublevels in the outermost energy level (valence shell). The core electrons are those electrons in the
energy levels below the valence shell (inside energy levels).
Atoms in the same column have the same valence electron configuration. The valence electron
configuration can easily be found by looking up the group number on the periodic table.
Group number
#of valence e-
1A
1
2A
2
3A(13)
3
4A(14)
4
5A(15)
5
6A(16)
6
7A(17)
7
8A(18)
8
Electron Dot diagrams (Lewis structure)
The electron dot diagram is a way of keeping track of and presenting the valence electron structure. To
write the electron dot structure; you write the element symbol, which represents the core (inner electrons
and nucleus) and then put one dot for each valence electron. Don’t pair up until they have to
Write the electron dot diagram for
·
Na
Na
Mg
C
O
F
Ne
He
·
Mg ·
·
·C·
·
··
·O:
·
··
· F:
··
··
: Ne :
··
··
He Electrons are paired because the valence shell is full.
Electron Configurations for Cations
Metals will have few valence electrons that are shielded from the nuclear charge. These valence
electrons are weakly attracted to the nucleus and will come off easily during chemical reactions. Metals
lose valence electrons to attain noble gas (stable) configuration (Metals are born losers). Metals react to
make positive ions (cations). If we look at an electron configuration, it makes sense. Na 1s22s22p63s1
loses 1 valence electron to become Na+ with an electron configuration of 1s22s22p6 that is the same as
the noble gas Ne. The Na+ has a more stable electron configuration than the Na atom.
Electron Configurations for Anions
Nonmetals have a many valence electrons that are poorly shielded by the core electrons (stongly
attracted to the nucleus)and their valence shells are almost full. Nonmetals strongly attract valence
electrons and can become more stable by gaining electrons to attain noble gas configuration. Nonmetals
Bonding and Nomenclature
HTHS Chemistry
2
make negative ions (anions). If we look at the electron configuration of Sulfur this makes sense. S
1s22s22p63s23p4 has 6 valence electrons and if it gains two valence electrons it becomes S-2
1s22s22p63s23p6 that has noble gas electron configuration for Ar.
Stable Electron Configurations
All atoms react to achieve the stable (low energy) noble gas configuration. Noble gases have 2 s and 6 p
electrons a total of 8 valence electrons. (He the only exception.)
All atoms react by changing their number of valence electrons to acquire a complete valence shell
(called the octet rule). Not all stable ions result in the noble gas configuration; there are a few
exceptions mainly in the transition metals. Zn 1s22s22p63s23p63d104s2 loses the two valence electrons to
become Zn2+ 1s22s22p63s23p63d10 that is stable but does not have the configuration of a noble gas. It
does have a complete valence shell. This is known as the pseudo noble gas electron configuration. Other
ions like Cu+, Ag+, Au+ and Cd2+ have pseudo noble gas configurations.
Ionic Bonding
Ionic bonds occur when the reacting elements have a large difference in electronegativity values (a
difference of approximately 1.7). Metals are born losers and nonmetals are winners, so when they react
there is a transfer of valence electrons. Cations and anions with stable noble gas configurations result
from the transfer of the valence electrons. The resulting electrostatic (charged) ions emit positive and
negative force fields. These ions can become more stable by neutralizing the force fields. A positive
force field neutralizes a negative force field, a force of attraction results. Anions and cations are held
together by these oppositely charged force fields. This force of attraction is called an ionic bond. Ionic
compounds contain ionic bonds between a metal ion and a nonmetal ion. Ionic compounds are called
salts. The simplest ratio of cation to anion is called the formula unit. The bond is formed through the
transfer of electrons. All the electrons must be accounted for, the total number of electrons lost must
equal the total number of electrons gained.
Na reacts with Cl. The Na loses 1 electron to form Na+ while the Cl gains 1 electron to form Cl-. The
ratio is 1:1, shown in NaCl.
Al reacts with S. The Al loses 3e- to form the Al3+ ion while S gains 2e- to form the S2- ion. The number
of electrons lost, 3, does not equal the number of electrons gained. To get the total lost equal to the total
gained there must be 2 Al (loss of 6) and 3 S (gain of 6) so the compound would be in a ratio of 2:3
Al2S3.
••
•S‫׃‬
•
•
Al•
•
•
Al•
•
••
•S‫׃‬
•
2Al3+ + 3S2 -
Al2S3
••
•S‫׃‬
•
Properties of Ionic Compounds
As mentioned above, an ionic compound can be more stable when the positive and negative force fields
are neutralized. Because the bonds are strong ionic compounds have high melting points and form brittle
solids.
Bonding and Nomenclature
HTHS Chemistry
3
Positive fields are neutralized when surrounded by as many negative fields as space will allow. Negative
force fields are neutralized when surrounded by as many positive fields as space will allow (closely
packed like charges repel each other). The oppositely charged ions arrange themselves in a regular
repeating pattern that neutralizes the force fields. This repeating pattern of ions creates a crystalline
structure called the crystal lattice. This crystalline structure is rigid because ions are strongly bonded.
In NaCl the Na+ is surrounded by 6 Cl- and 6 Na+ surrounds each Cl-. The coordination number gives the
number of oppositely charged ions that surround each ion in the crystal lattice (structure). The
coordination number for the Na+ ion in NaCl is 6. NaCl crystals are small cubes. This represents the face
of the cube of NaCl; the pattern would be repeated. There is a Cl- in front and behind the central Na+ and
there is a Na+ ion in each vertices of the cube face. This arrangement is called a face-centered cube.
Na +
Cl –
Cs is in the same family as Na but forms a larger ion that results in a different crystal lattice (structure).
Cl –
Cs +
Notice the Cs+ is surrounded by 8 Cl- ions so the coordination number is 8. The Cs+ is in the center of
the cube. This arrangement is called a body-centered cube.
Do ionic compounds conduct an electric current?
To conduct electricity there must be a movement of charges. In a solid the ions are locked in place, ionic
solids are non-conductors. When melted, the ions can move around so melted ionic compounds conduct.
Dissolving ionic compounds in water allows the charged ions to move, solutions of ionic compounds
conduct an electric current.
Bonding and Nomenclature
HTHS Chemistry
4
Ionic Bonding review questions
1. How many valence electrons does each of the following atoms have?
a. Potassium
b. carbon
c. magnesium
d. oxygen
2. State how many electrons will be gained or lost for the following elements in forming an ion
a. calcium
b. fluorine
c. aluminum
d. oxygen
3. Write the formula for the ion when the element becomes isoelectronic with a Noble gas.
a. sulphur
b. sodium
c. fluorine
d. barium
4. Use electron dot symbols (Lewis structure) to determine the formula unit and then name the
ionic compounds formed when the following elements combine.
a. potassium and iodine b. calcium and chlorine
c. aluminum and sulfur d. magnesium and phosphorus
5. Define the term valence electrons.
6. Write the electron dot structures for each of the following elements.
a. S
b. C
c. Al
d. Be
e.
Cl
7. The atoms of the noble gas elements are stable. Explain.
8. Write the complete electron configuration for the following atoms and ions. Comment on the
results:
Ar, K+, Ca2+
9. What are cations and anions? How and why are cations and anions produced?
10. Why are ionic compounds electrically neutral?
11. Write the correct chemical formula for the compound formed by the following pairs of ions:
a. Na+ and Fb. K+ and S2c. Ca2+ and N3d. Al3+ and O212. Which of the following pairs of elements are most likely to form ionic compounds?
a. chlorine and bromine b. potassium and helium
c. lithium and fluorine
d. iodine and sodium
13. Write the formula for the ions in the following compounds.
a. LiF
b. BaO
c. Na2S
d. Al2O3
e. Ca3N2.
14. In your own words describe a crystal.
15. Most ionic compounds are brittle. Why?
16. Why does molten MgCl2 conduct an electric current although crystalline MgCl2 does not? How
else can we get MgCl2 to conduct a current?
NOMENCLATURE - IONIC COMPOUNDS
Binary ionic compounds contain two elements and are composed of cations (metals) and anions
(nonmetals). The total number of electrons lost must equal the total number of electrons gained because
compounds are neutral. To name binary ionic compounds you first name the cation! BUT!!!!
Bonding and Nomenclature
There two types of cations:
HTHS Chemistry
5
(i)
cations with a constant (fixed) charge - Group A cations have only one possible charge. The
Alkali metals form 1+ ions, Alkaline earth metals form 2+ ions. To name these cations you use
the element name. K1+ is the potassium ion, Ca2+ is the calcium ion, Al3+ is the aluminium ion.
(ii)
cations with a variable charge ( more than one charge is possible) most transition and heavy
metals can have more than one charge, Iron can form a 2+ or 3+ ion, Copper can form two ions
1+ or 2+. To name these ions you must identify the charge that the ion has and then use a Roman
numeral in parenthesis after the element name. Fe2+ iron (II) ion, Fe 3+ iron(III) ion, Cu2+
copper(II) ion. To name Fe2O3 you must know which iron ion is there. Nonmetal anions have
constant charges. O is 2- but there 3 of them so the oxygen has gained a total of 6 e-. The iron
must have lost a total of 6 electrons; there are two iron ions so each must have lost 3 e- Fe3+
iron(III); so the name is iron(III) oxide.
To complete the name of the binary ionic compound you must name the anion (nonmetal). The anion
name is derived from the root of the element name and the ending changed to ide. O2- is oxide, Cl1chloride, N3- nitride.
Name the following binary ionic compounds:
CuO
KCl
Na3N
AlN
Ca3P2
PbO
PbO2
Na2S
CoCl2
Ternary ionic compounds have more than two elements. Ternary compounds usually contain a cation
and a polyatomic ion. Polyatomic ions are groups of atoms that stay together, act as a single unit and
have a charge. When a polyatomic ion is taken more than once in a compound the ion is placed within
parenthesis and a subscript is used. Mg is 2+ NO3 is 1- so magnesium nitrate is Mg(NO3)2. Most
poyatomic ions have an ending of ate or ite (a few exceptions have an ide ending).
You must memorize the polyatomic ions (more listed on the back of your periodic table handout)
Ethanoate (acetate) CH3COO- or C2H3O2-1
Nitrate NO3-1
Nitrite NO2-1
Hydroxide OH-1
Permanganate MnO4-1
Cyanide CN-1
Sulfate SO4-2
Sulfite SO3-2
Carbonate CO3-2
Chromate CrO4-2
Dichromate Cr2O7-2
Phosphate PO4-3
Phosphite PO3-3
The only polyatomic cation is Ammonium NH4+1
Naming Ternary Ionic Compounds
1. Name the cation as shown above
2. Name the polyatomic anion
Examples:
CaSO4 cation is Calcium; the anion is sulphate. The compound is calcium sulphate.
Bonding and Nomenclature
HTHS Chemistry
6
Fe2(SO4)3 SO4 is 2- x 3 = 6-; Fe total = 6+ so each Fe must be 3+. The name is iron(III) sulfate.
Name the following ternary ionic compounds:
NaNO3
CaSO4
CuSO3
(NH4)2O
LiCN
Fe(OH)3
(NH4)2CO3
NiPO4
Write Formulas for Ionic compounds
For all ionic compounds the charges have to add up to zero; that is, the total number of electrons lost
must equal the total number of electrons gained. To write the formula of a compound:
1. Write the cation symbol with the charge.
2. Write the anion symbol with the charge.
3. Write the subscript for each so that the total e- lost = total e- gained
Put polyatomics in parenthesis when the subscript is greater than one.
Write the formula for calcium chloride.
Calcium is Ca+2 loss of two
Chloride is Cl-1 gain of one
Need 2 x Cl-1 for a gain of two. Ca+2 Cl2-1 the formula unit is CaCl2
Write the formulas for these ionic compounds:
Lithium sulfide
tin(II) oxide
Magnesium fluoride
Copper(II) sulfate
Iron(III) phosphide
Sodium nitrate
Iron(III) sulfide
Ammonium sulfide
Barium nitrate
Tin(IV) oxide
Ammonium phosphate
Covalent bonding
How does H2 form?
Bonding and Nomenclature
HTHS Chemistry
7
Two hydrogen atoms approach each other. The electron clouds overlap as the atoms collide. The
electron cloud density is greatest between the nuclei, which keeps the nuclei from repelling each other.
The electrons caught between the nuclei are attracted equally to both nuclei. As the atoms (nuclei) move
closer the nuclei repel each other and the atoms move away from each other. The shared electrons are
attracted to both nuclei and pull the nuclei back together. These electrons hold the atoms together and
are called bonding electrons, the atoms are now joined together to form a molecule. The overlapping
atomic orbitals form a new molecular orbital, which is densest in between the two nuclei to form a
bonding orbital. The bonding electrons are shared between the two valence shells of the atoms so the
bond is called a covalent bond. There is one pair of bonding electrons being shared so this is a single
covalent bond. Each hydrogen atom effectively has two valence electrons and therefore has the electron
configuration of Helium (noble gas configuration). The atoms in molecular compounds neither gain nor
lose electrons but instead share them. Molecular compounds also follow the octet rule. They share
electrons in a way that allows both atoms in a bond to have an octet of valence electrons. A covalent
bond is the electrostatic attraction between the nuclei of two adjacent atoms and a pair of shared
bonding electrons.
Why do covalent bonds form?
Nonmetals have high electronegativity values so they tend to hold onto their valence electrons and can’t
give away electrons to bond. Nonmetals still want noble gas configuration but get the stable electron
configuration by sharing valence electrons with each other. By sharing both atoms get to count the
electrons toward noble gas configuration.
Another example of a covalent bonding:
..
..
:F.
.F:
··
··
Fluorine has seven valence electrons and a second atom also has seven. Note there are three pairs of
electrons (paired electrons are stable) and one electron not paired. To become stable each Fluorine atom
must gain one electron and each has one electron that can be shared. The fluorine atoms each share a
valence electron to form a molecule that is held together by a single covalent bond. The bonding
electrons are equally attracted to both nuclei (atoms have the same electronegativity). Names given to
these pairs of electrons, as shown below, are bonding electron pairs for the shared pairs and lone pairs
for unshared electrons that fill a valence level for only one atom.
.. ..
:F : F:
·· ··
The single covalent bond is a sharing of two valence electrons (one pair). Covalent bonding occurs with
nonmetals. Covalent bonding is different from an ionic bond because the atoms actually form molecules
(at least two specific atoms are joined by bonding electrons).
Multiple Bonds
Sometimes atoms share more than one pair of valence electrons.
Oxygen atoms have six valence electrons; two pairs and two unpaired valence electrons. To become
stable, each oxygen atom must gain two electrons but each atom has two electrons available for sharing.
If each atom offers two electrons for sharing there will be two pairs of bonding electrons shared between
the two oxygen nuclei.
Bonding and Nomenclature
HTHS Chemistry
8
..
..
:O.
.O:
·
·
Since there are two pairs of bonding electrons being shared; there is a double covalent bond holding the
oxygen atoms together to form an oxygen molecule. A double bond is when atoms share two pair of
electrons (4e-). The bonding electrons are equally attracted to both nuclei (atoms have the same
electronegativity).
..
..
:O :: O:
Nitrogen atoms have five valence electrons; one pair and three unpaired valence electrons. To become
stable, each nitrogen atom must gain three electrons but each atom has three electrons available for
sharing. If each atom offers three electrons for sharing there will be three pairs of bonding electrons
shared between the two nitrogen nuclei.
.
.
:N.
.N:
·
·
Since there are three pairs of bonding electrons being shared; there is a triple covalent bond holding the
nitrogen atoms together to form an nitrogen molecule. A triple bond is when atoms share three pair of
electrons (6). The bonding electrons are equally attracted to both nuclei (atoms have the same
electronegativity).
..
..
N ::: N
All the examples above formed homonuclear molecules. What happens if the atoms are different
elements? How are the molecules formed? You put the atoms together to end up with each of the atoms
having a stable electron configuration. To have the correct formula you must know the elements in the
molecular compound. For example- show how water is formed with covalent bonds. Water has
hydrogen and oxygen present. Each hydrogen atom has 1 valence electron and each hydrogen atom
wants 1 more electron. The oxygen has 6 valence electrons and the oxygen wants 2 more electrons. A
hydrogen atom has one electron to share and needs to gain one electron; each oxygen atom has two
electrons to share and needs to gain two electrons. There must be two hydrogen atoms to satisfy the
oxygen atom.
..
:O.
.H
·
·H
The oxygen atom forms two single covalent bonds; one bond with each of the hydrogen atoms. Every
atom has full energy level (valence shell). Are these electrons shared equally? Oxygen has an
electronegativity value of 3.5. Hydrogen has an electronegativity value of 2.1. Oxygen atoms have a
greater ability to attract valence-bonding electrons. The bonding electrons will spend more of their time
around oxygen; oxygen will have an electron cloud density higher than usual. The oxygen atom has a
negative charge but it is not a complete negative because the electron has not been transferred. There is a
partial negative charge because the bonding electrons spend a greater part of the time with oxygen and
the lesser time with hydrogen. The hydrogen has lost the bonding electron a great deal of time and so the
hydrogen atom has a partial positive charge. Because there is a partial separation of charge the bond is
called a polar covalent bond. Polar refers to the charge separation.
Bonding and Nomenclature
HTHS Chemistry
9
δ- ..
:O:H δ+
··
H
δ+
Water has two single polar covalent bonds.
Carbon dioxide
Carbon is central atom (I have to tell you). Carbon has 4 valence electrons and wants 4 more electrons.
Carbon has 4 electrons available for sharing. Oxygen has 6 valence electrons and wants 2 more. Oxygen
has two electrons available for sharing. To complete carbon the must be two oxygen atoms.
..
·
··
:O.
·C·
·O:
·
·
·
3.5
2.5
3.5 Electronegativity values.
Two of the C valence electrons pair with the two valence electrons from the oxygen atom to form double
covalent bonds. The CO2 molecule has two double covalent bonds. Because there is a fairly large
difference in electronegativity values the bonding electrons are not shared equally; the bonding electrons
spend more time around the oxygen and less around the carbon. The unequal sharing of bonding
electrons creates two double polar covalent bonds.
..
..
δ- :O::C::O: δδ+
Covalent bonds can vary in the equality of valence electron sharing. Chemists use the following scheme
to identify the class of bond formed:
Electronegativity Difference
0 - .4
.4 – 1.7
≥ 1.7
Covalent bond
Polar covalent
Ionic bond
A continuous spectrum where there is no definite boundary between bond types. As
sharing becomes less equal the bond becomes polar and as the unequal sharing becomes
greater the bond becomes ionic in character.
Coordinate Covalent Bond
A coordinate covalent bond occurs when one atom donates both bonding electrons in a covalent bond.
Once formed the coordinate covalent bond is the same as any other covalent bond. Example:
To form a coordinate covalent bond the atom must have lone pair (nonbonding electrons). Nitrogen has
a lone pair of electrons and 3 bonding pairs of electrons in ammonia. The ammonia molecule is a dipole;
the nitrogen is δ- while the hydrogen is δ+. The δ- attracts positive charges, like a H+. The H+ needs two
electrons to fill the valence shell, the nitrogen atom has a pair of nonbonding electrons and a partial
negative charge the H+ can take a share of these electrons and bond with the ammonia to form the
polyatomic ammonium ion.
Most polyatomic ions contain a combination of covalent and coordinate bonds.
Bonding and Nomenclature
HTHS Chemistry
10
Attractive Forces that hold particles of matter together
Matter can be found in three states solid, liquid or gas. The state of matter can be determined by the
strength of the attractive forces between the particles that make up matter. Strong forces of attraction
yield solids, weaker forces of attraction will yield liquids and extremely weak forces of attraction will
yield gases.
The strongest force of attraction is the attraction between electrostatic force fields of charged particles
(cations attracting anions) is called an ionic bond. Ionic bonds are found in ionic compounds, which are
all crystalline solids with high melting and boiling points. If substances are not ionic then they are
molecular and have different types of forces holding the molecules together.
Intermolecular Forces
The forces of attraction between molecules are called intermolecular forces. Intermolecular forces are
weak in comparison to ionic forces. Intermolecular forces vary in strength; the strength of these forces
determines if the molecular substance will be a solid a liquid or a gas. A strong intermolecular force
leads to a solid; a weaker intermolecular force leads to a liquid and extremely small or no intermolecular
force leads to a gas. There are two kinds of intermolecular forces:
1. Dipole Interactions
2. Dispersion forces
1. Dipole-Dipole forces.
These are the forces experienced between two polar molecules. The partial negatively charged region of
one polar molecule attracts the partial positively charged region of another neighboring molecule. These
opposite charges attract but are not as strong as the forces in ionic solids.
δ+H – F δ- - - - - - - - - δ+H – F δdipole-dipole
attraction
A special kind of dipole interaction involves hydrogen and is thus called Hydrogen Bonding. Hydrogen
Bonding occurs when a hydrogen atom is bonded to a very electronegative atom of a neighboring
molecule. A hydrogen bond is formed when the partial positively charged hydrogen is attracted to a lone
pair of electrons on the highly electronegative atom of a neighboring molecule. The hydrogen partially
share the lone pair in the molecule next to it. Hydrogen bonding forces are stronger than dipole forces.
Example:
2. Dispersion forces
These are the forces experienced between two nonpolar molecules. In any atom or molecule, the
electrons are constantly in motion. The electrons are dispersed around the particle in the electron cloud.
At any instant, the electron distribution may be slightly uneven one part of the cloud becomes denser
than another. This dense area momentarily takes on a slight (partial) negative charge and another area of
the molecule has a momentary partial positive charge. The resulting separation of charge is very short
lived because the electrons move on. The attraction of these momentary charge separations causes weak
Bonding and Nomenclature
HTHS Chemistry
11
temporary forces. These weak are called van der Waal’s forces. The strength of these forces depend on
two conditions:
(i) The strength of the dispersion force depends on the size of the atom/molecule - the larger the particle
the longer the duration of the charge separation thus the greater the size of the force of attraction. The
greater the size the longer the charge separation exists, the more interaction of dispersion charges.
The e- represents an area of high electron cloud density and + represents an area of low electron cloud
density.
(ii) The magnitude of the dispersion force also depends on the number of electrons; the more electrons
an atom/molecule has the greater the chance of unequal distribution within the electron cloud; thus
stronger dispersion forces.
Example:
Fluorine, F2 and Chlorine, Cl2, have few electrons, weak forces and are gases. Bromine has more
electrons, stronger forces, and is a liquid. Iodine has even more electrons, stronger forces, and is a solid.
When you compare the sizes of these molecules the smallest are gasses (weak forces), the middle size is
a liquid (stronger forces) and the biggest molecules is a solid (strongest forces).
Dipole interaction forces are stronger than dispersion forces but the strongest of the intermolecular
forces is Hydrogen Bonding.
Characteristics of Ionic and Covalent Compounds
Characteristic
Representative Unit
Bond formation
Type of elements
Physical state
Melting point
Solubility in water
Electrical conductivity (aquous)
Ionic compound
Formula unit
Electron transfer
Metal with nonmetal
Solid
High(>300ºC)
High
Good conductor
Covalent compound
Molecule
Electron sharing
Nonmetal with nonmetal
Solid, liquid or gas
Low(<300ºC)
High to low
Poor to nonconducting
Bonding and Nomenclature
HTHS Chemistry
12
Covalent Bonding Review questions
1.
Draw the electron dot structure for the following diatomic halogen molecules.
a) chlorine
b) iodine
2.
Draw the electron dot structure for the following covalent molecules:
a) H2S
b) PH3
c) BrF
3.
Draw the electron dot structure for the polyatomic ions (a) NH4+
4.
Explain the difference between a nonpolar covalent bond and a polar covalent bond.
5.
Use electronegative differences to identify the types of bonds between atoms in the following
pairs of elements (ionic, polar covalent or nonpolar covalent).
a) H and Cl
b) K and Cl
c) N and O
d) I and F
e) Br and Br
6.
Which covalent bond is the most polar and which is the least polar?
a) H – Cl
b) H – Br
c) H – S
d) H – C
e) F – F
7.
Explain why the noble gasses are monatomic but the halogens are diatomic molecules.
8.
Draw the electron dot structure for each of the following:
a) H2O
b) H2O2
c) PCl3
d) NH3
9.
Classify the following compounds as ionic or covalent.
a) H2S
b) Na2S
c) HCl
d) MgCl2
10.
Explain why atoms form chemical bonds and describe the difference between an ionic bond
and a covalent bond.
(b) H3O +
Bonding and Nomenclature
HTHS Chemistry
11.
State the number of electrons shared by two atoms in a:
a) single covalent bond
b) double covalent bond
12.
13.
13
c) triple covalent bond
Draw electron dot structures for the following molecules:
a) I2
b) OF2
c) H2S
d) NI3
e) HCN
The bonds between the following pairs of elements are covalent. Arrange them according to
polarity, the most polar to the least polar.
a) H – Cl
b) C – H
c) H – F
d)H – O
e) H – H
f) S – Cl
Molecular Compounds
Writing names and formulas for Molecular compounds
Molecular compounds are made of just nonmetals. The smallest piece is a molecule that contains atoms
held together by the sharing of electrons. There is no transfer of electrons no charged particles.
Molecules are electrically neutral particles. The molecular formula shows the elements present and the
number of each atom present in the molecule.
SO3 means each molecule contains one sulfur atom and three oxygen atoms. How do you name
molecular compounds? The molecular compound name tells you the number of atoms. To state the
number you use prefixes:
Prefixes
1 mono6 hexa2 di7 hepta3 tri8 octa4 tetra9 nona5 penta10 decaTo write the name
1. write the name of the first element with a prefix to state the number. DO NOT USE MONO for the
first element. If there is one atom use the element name no prefix
2. write the name of the second element with the correct prefix use mono if there is only one. No double
vowels when writing names (oa oo).
Examples: SO3 sulfur trioxide; CO carbon monoxide; CO2 carbon dioxide; NO nitrogen monoxide; N2O
dinitrogen monoxide.
Name these:
NO2
Cl2O7
CBr4
Write formulas for these:
diphosphorus pentoxide
tetraiodide nonoxide
sulfur hexaflouride
nitrogen trioxide
CO2
N2O
P2O5
Carbon tetrahydride
phosphorus trifluoride
aluminum chloride
Acids
Writing names and Formulas
Acids are molecular compounds (nonmetal with nonmetal) that give off hydrogen ions when dissolved
in water. The water tears the molecules apart into cations (H+) and anions. The anion determines the
name of the acid.
Bonding and Nomenclature
HTHS Chemistry
14
Naming acids
Acids are either binary or ternary.
Binary acids (two elements) – the hydrogen is attached to a monatomic anion.
The hydrogen is named – using the prefix hydro
The anion name changes the ide ending to ic acid
HCl - hydrogen ion and chloride ion
HCl - hydrochloric acid
H2S hydrogen ion and sulfide ion
hydrosulfuric acid
Ternary acids (more than two elements) – the hydrogen is attached to a polyatomic anion. If the anion
has oxygen in it, the anion ending is -ate of -ite
To name ternary acids the ending (suffix) changes from the -ate to -ic acid
HNO3 Hydrogen and nitrate ions - Nitric acid
If the anion is an –ite the ending (suffix) changes from the -ite to -ous acid
HNO2 Hydrogen and nitrite ions
Nitrous acid
If the ternary acid ends in ide follow the binary acid rules
HCN
Hydrogen ion with cyanide ion
Hydrocyanic acid
Metallic Bonds
Metallic bonding refers to the forces that exist between atoms in a pure metal or in a metal alloy.
Metallic bonds do not involve the sharing of valence electrons as in covalent bonding or the donation of
valence electrons as in ionic bonding. Metallic bonding is used to explain the physical properties of
metals: malleability, ductility, electrical conductivity, and heat conductivity.
Metals have a unique set of properties so chemists have a devised a different model of bonding to
explain metallic bonding. Metals hold onto their valence electrons very weakly. The valence electrons
are free to move from one atom to the next. Valence electrons are free to move throughout the solid. The
model is called the free-electron model. Metals can be described as a group of cations tightly packed
together and are surrounded by valence electrons, which are in continuous motion around the cations.
These freely moving valence electrons spend most of the time between one metal atom and the next.
You can visualize metals as positive ions embedded in a “sea” of valence electrons. After loosing their
valence electrons, the ions share all of the electrons.
This attraction of “free-floating” valence electrons around the cation is called a metallic bond. Because
the valence electrons do not belong to any one particular metal atom metals can be viewed as metallic
cations held together by a common cloud of valence electrons. Metal atoms are held together in the solid
as positive ions floating in a sea of electrons. The loosely held valence electron cloud allows for the flow
of charge. Thus metals are conductors of electricity. The valence electron cloud also isolates the cations
from one another as they “roll” around in the structure. This ensures a great mobility for the cations as
well. So, it is possible to make the cations slide on top of one another so that we can change the shape of
their structure without affecting the composition of the metal.. Metals are malleable that is they can be
hammered into different shapes (can be bent). Metals are ductile that is they can be drawn into wires
(valence electrons allow atoms to slide by each other).
Bonding and Nomenclature
HTHS Chemistry
emetal cation
15
The diagram represents metal atoms that have released their two valence electrons and are
embedded in a sea of electrons
Metals have closest packing patterns and when melted they will mix to form solutions which cooled
solidify to form what we refer to as alloys. Gold jewelry is a common alloy of gold, copper and silver.
The new mixture will not have the same properties as the original respective metals.
Metallic Bonding Review Questions
1. In your own words define a metallic bond.
2. Describe what is meant by the terms ductile, and malleable. Explain why metals exhibit these
properties.
3. Explain why it is possible to bend metals but not ionic crystals?
4. Explain why metals are good conductors of electricity.
Nomenclature
Name the following binary ionic compounds:
Write a formula for the binary ionic compounds:
1) MgS
__________________________
1) magnesium oxide
__________________
2) KBr
__________________________
2) lithium bromide
__________________
3) Ba3N2
__________________________
3) calcium nitride
__________________
4) Al2O3
__________________________
4) aluminum sulfide
__________________
5) NaI
__________________________
5) potassium iodide
__________________
6) SrF2
__________________________
6) strontium chloride
__________________
7) Li2S
__________________________
7) sodium sulfide
__________________
8) CaO
__________________________
8) radium bromide
__________________
9) K2S
__________________________
9) magnesium sulfide
__________________
10) LiBr
__________________________
10) aluminum nitride
__________________
Bonding and Nomenclature
HTHS Chemistry
Name for the following, use the Stock system.
16
Use the Stock system to write the formula for:
1) CuS
_____________________________
1) iron(II) chloride
________________
2) Fe2O3
_____________________________
2) copper(I) sulfide
________________
3) FeI2
_____________________________
3) lead(IV) iodide
________________
4) Cu2S
_____________________________
4) tin(II) fluoride
________________
5) SnCl2
_____________________________
5) chromium(III) oxide
________________
6) CuO
_____________________________
6) iron(II) nitride
________________
7) PbF2
_____________________________
7) cobalt(III) phosphide
________________
8) CuCl2 _____________________________
8) iron(III) chloride
________________
9) CuBr
_____________________________
9) copper(II) sulfide
________________
10) PbO
_____________________________
10) lead(II) bromide
________________
Ternary Compounds
Ternary Compounds
Write the correct chemical name for:
Write the correct chemical formula for:
1) AlPO4
_____________________
1) potassium phosphate
___________
2) CaCO3
_____________________
2) aluminium hydroxide
___________
3) Mg(OH)2
_____________________
3) sodium hydrogen carbonate
___________
4) K2SO4
_____________________
4) calcium nitrate
___________
5) Na2CO3
_____________________
5) sodium carbonate
___________
6) NH4NO3
_____________________
6) tin(II) carbonate
___________
7) Sn(NO3)2
_____________________
7) lead(II) sulfate
___________
8) FePO4
_____________________
8) iron(III) carbonate
___________
9) PbCO3
_____________________
9) potassium hydroxide
___________
10) Cu2CO3
_____________________
10) ammonium sulfate
___________
Bonding and Nomenclature
Write the correct formula for:
HTHS Chemistry
17
Write the correct name for:
1) dinitrogen monoxide
_____________
1) N2O3
_____________________________________________
2) nitrogen trifluoride
_____________
2) NI3
_____________________________________________
3) carbon dioxide
_____________
3) SF6
_____________________________________________
4) diphosphorous pentoxide
_____________
4) CO
_____________________________________________
5) sulfur dioxide
_____________
5) P2O5
_____________________________________________
6) silicon tetrachloride
_____________
6) SO2
_____________________________________________
7) silicon dioxide
_____________
7) SiO2
_____________________________________________
8) dinitrogen pentasulfide
_____________
8) N2S5
____________________________________________
9) carbon monoxide
_____________
9) CO2
_____________________________________________
10) nitrogen monoxide
_____________
10) PF3
____________________________________________
Write the formula for each of the acids:
Nitric acid
Hydrocyanic acid
Ethanoic acid (Acetic acid)
Hydrobromic acid
Sulfurous acid
Hydrochloric acid
Phosphoric acid
Nitrous acid
Hydrofluoric acid
Hydroiodic acid
Phosphorous acid
Carbonic acid
Sulfuric acid
Nitrous acid
Name each of the following acids:
H3PO4(aq)
HCl (aq)
H2SO4(aq)
HNO2(aq)
HI (aq)
CH3COOH(aq)
HF (aq)
H3PO3(aq)
HCN (aq)
H2CO3(aq)
H2SO3(aq)
HNO3(aq)
HBr (aq)
Bonding and Nomenclature
HTHS Chemistry
18
All Compounds
Binary compound (2 elements)
Metal – non-metal
(Ionic compounds)
Hydrogen – non-metal
Ternary compound (> 2 elements)
(Ionic compound - with a polyatomic
ion)
Non-metal – non-metal
(Molecular compounds)
Metal with polyatomic
ion
Cation constant charge
Cation variable charge
Name 1st element with proper prefix (di, tri
…Never mono)
Name 2nd element (root + ide ending) with
proper prefix (mono, di, tri ..)
Metal with constant charge
Name the cation
Metal with variable charge (Stock system)
Name the cation - Roman numeral shows charge
Name the polyatomic anion
Name cation (element name)
Name anion (root + ide ending)
Stock System
Name cation (Roman numeral charge)
Name anion (root + ide ending)
In aqueous phase name as acid
Use the prefix hydro
change ide ending to ic acid
As compound hydrogen + (Root + ide
ending)
In aqueous phase H with polyatomic anion, name as acid
anion ending -ate changes to -ic acid
anion ending -ite changes to -ous acid