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Bonding and Reactions
A compound is two or more elements
chemically combined.
Compounds are formed from elements in
chemical reactions.
The forces that hold atoms together are called
bonds. Atoms form bonds in two main ways:
(1) An ionic bond is a bond formed by the
transfer of electrons (between oppositely
charged ions).
Example: Na+1 + Cl-1 -- NaCl
(1) Covalent bonding is bonding involving a
sharing of electrons. A bond is a shared
pair of electrons.
Example:
C
+
4H
---
H
C
H
A molecule is the smallest part of a covalent
compound that retains all of its
characteristics.
The term formula unit is used to describe the
smallest part of an ionic compound that
retains all of the characteristics of the
compound.
A formula is used to represent a single
molecule or formula unit of a compound.
Example: H2O – 2 atoms of hydrogen with
one atom of oxygen
A subscript is a small number indicating the
number of atoms or ions of each element that
are present in a compound. The subscript one
is understood.
An atom that gains or loses electrons is in a
charged condition and is called an ion.
Atoms that form negative ions gain electrons
(anions).
Example: O-2
Atoms that form positive ions lose electrons
(cations).
Example: Na+1
An ion made of more than one atom is a
polyatomic ion.
Examples: OH-1, SO4-2, NO3-1 (Refer to the
complete list on your worksheet or in your
text. You must learn the names, formulas, and
charges.)
When an atom or group of atoms has an
overall charge, that charge is called an
oxidation number. Oxidation numbers must
be used to write a formula correctly.
Rules For Determining Oxidation Numbers
(1) Metals (positive) in groups IA to IIIA
have an oxidation number equal to the
group number.
Example: Ca is in group IIA so it is +2.
(2) Many B group metals have more than
one charge and must be memorized.
Example: Sn can be +2 or +4.
(3) Nonmetals in groups IVA to VIIA form
ions with negative charges equal to the
group number – 8.
Example: O is in groupVIA , so 6-8 = -2
(4) Polyatomic groups – refer to chart on
worksheet or in your text
Atoms and ions combine in definite whole
number ratios.
If the charges on the ions forming a
compound add to zero, the formula is written
correctly.
In a binary compound (made of 2 elements)
the positive ion is written first and the
negative ion is written second.
Example: Ca+2 + Br –1 ----- CaBr2
Br –1
Ca+2 + Br –1 --- CaBr2
(calcium bromide)
Example: Al +3 + S -2
Al2S3
(aluminum sulfide)
Non-binary compounds contain at least 3
elements.
Polyatomic ions must be written in
parenthesis, when expressing more than one.
Example: Al +3 + SO4-2
Na+1 + NO3-1
Al2(SO4)3
NaNO3
Naming Compounds
(1) For binary compounds:
(A) Write the name of the element with
the positive charge.
(B) Add the name of the negative
element modified to end in -ide.
Examples:
NaCl – sodium chloride
Al2S3 – aluminum sulfide
Roman numerals are used in naming
compounds that contain elements with more
than one oxidation number.
Examples:
Cu2S – copper (I) sulfide
CuS – copper (II) sulfide
FeF2 – iron (II) fluoride
FeF3 – iron (III) fluoride
PbCl2 – lead (II) chloride
PbCl4 – lead (IV) chloride
An older system for naming cations having
more than one possible charge uses an –ic
ending for the ion with the higher charge and
an –ous ending for the ion with the lower
charge.
In addition, the Latin names for some of the
elements are used.
Examples:
stannous chloride – SnCl2
cupric oxide – CuO
cobaltous fluoride – CoF2
auric chloride – AuCl3
Occasionally, covalent compounds are named
using a system in which Greek prefixes are
used to indicate the number of atoms present.
Mono – 1
Di – 2
Tri – 3
Tetra – 4
Penta – 5
Hexa – 6
Hepta – 7
Octa – 8
Nona – 9
Deca – 10
Hints:
(1) The prefix mono- may be omitted for
the first element (PCl3 – phosphorus
trichloride).
(2) For oxides, the ending “a” or “o” in the
prefix is sometime omitted (N2O4 –
dinitrogen tetroxide).
Examples:
CO2 –carbon dioxide
CO - carbon monoxide
CCl4 – carbon tetrachloride
SO3 – sulfur trioxide
Some compounds have names ending in –ide
and are not binary. These compounds contain
polyatomic groups.
Examples:
OH-1 – hydroxide
CN-1 – cyanide
O2-2 – peroxide
In naming a compound containing more than
2 elements, write the name of the first element
or polyatomic group (NH4+1). Then write the
name of the second element (-ide ending) or
the polyatomic group (do not change the
ending).
Polyatomic ions that contain oxygen end in
-ite or –ate.
Examples:
CuSO4 –copper(II) sulfate
ZnCO3 – zinc carbonate
NH4Cl – ammonium chloride
BaOH – barium hydroxide
Chemical equations – are used to represent
chemical reactions.
The starting substances are called reactants.
The substances formed are called products.
Example: hydrochloric acid and sodium
hydroxide react to form sodium chloride and
water
HCl(aq) + NaOH(aq) - NaCl(aq) + H2O(l)
phase labels (physical state symbols)
(g) – gas
(l) – liquid
(cr) – crystalline solid
(aq) – aqueous (water solution)
(s) – solid
There are 3 steps to writing an equation
correctly:
(1) Determine what the reactants and
products are.
(2) Assemble the parts of the equation:
(a) Write the formulas correctly.
(b) Connect reactants and products
using an arrow to show the direction
of the reaction.
------- means yields or produces
(c) Reactants are written to the left,
products on the right.
(3) Balance the equation by changing the
coefficients (stoichiometric coefficients).
Never change subscripts.
Example:
Propane reacts with oxygen to give carbon
dioxide and water.
(1) reactants – propane and oxygen
(2) products – carbon dioxide and water
(3) formulas/equation –
C3H8(g) + O2(g) ---- CO2(g) + H2O(l)
(4) balance – to get the same number of
atoms of each element on both sides of
the equation
C3H8(g) + 5 O2(g) ---- 3 CO2(g) + 4 H2O(l)
Practice: Balance the following
(1) __ Al2O3 -- __ Al + __ O2
(2) __ Na2S + __ O2 -- __ Na2O + __ SO2
(3) __ C + __ H2O -- __ H2 + __ CO
(1) 2 Al2O3 -- 4 Al + 3 O2
(2) 2 Na2S + 3 O2 -- 2 Na2O + 2 SO2
(4) C + H2O -- H2 + CO (balanced)
Four Types of Reactions
(1) Single displacement – one element
displaces another in a compound
M + M1Nm -- MNm + M1
Nm + MNm1 --MNm + Nm1
Examples:
Cl2 + 2 KBr -- 2 KCl + Br2
Na + AgCl -- NaCl + Ag
(2) Double displacement – the positive and
negative portions of 2 compounds are
interchanged
MNm + M1Nm1 --- MNm1 + M1Nm
Examples:
PbCl2 + Li2SO4 -- 2 LiCl + PbSO4
ZnBr2 + 2 AgNO3 -- Zn(NO3)2 + 2 AgBr
(3) Decomposition – substances break into
simpler substances when energy (heat,
light, electricity) is supplied
A -- B + C + D
Examples:
H2CO3 -- H2O + CO2
2 KClO3 -- 2 KCl + O2
(4) Synthesis – 2 or more substances
combine to form one new substance
B + C + D -- A
Examples:
2 H2 + O2 -- 2 H2O
NH3 + HCl -- NH4Cl
Law of Conservation of Matter
Matter can neither be created nor
destroyed, but can be converted from one
form to another. Therefore, the mass of the
reactants in an ordinary chemical reaction
must equal the mass of the products.
All chemical reactions involve changes in
energy. The law of conservation of energy
states that the total energy in a system is
constant.
If the energy of the products is less than the
energy of the reactants, energy is released
during the reaction. This is an exothermic
reaction.
reactants -- products + energy
example: the combustion of propane
If the energy of the products is greater than
the energy of the reactants, energy must be
absorbed during the reaction. This is an
endothermic reaction.
Example: instant ice pack
Molecular mass is the sum of the atomic
masses for all elements (atoms) in a
compound.
Examples:
NaCl
1 Na – 23
1 Cl – 35.5
58.8 AMU
Al(OH)3
1 Al – 27
3 O – 48 (3 x 16)
3 H – 3 (3 x 1)
78 AMU