Download Chemistry 1, Chapter 5

Chemistry 1, Chapter 5
Ions and Ionic Compounds
Section 1, Simple Ions
1. Chemical Reactivity
•
Atoms react to achieve a stable electron
configurations
– Remember, the stable arrangement is a full
outer level of electrons, usually 8
• How reactive an element is depends on its outer
electron configuration
– noble gases are non-reactive because they
already have a full outer level
• The idea that 8 outer electrons (electrons in the s
and p sublevels) is a full level and they make an
atom stable or unreactive is called the octet rule
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There are threes ways that atoms can get
a full outer level
1. they can gain electrons
2. they can lose electrons and expose a full level
below
3. they can share electrons
•
in this chapter we are going to learn about
gaining and losing electrons
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From these three ways to get a full outer
level we can form three types of bonds
1. ionic – which involves the gain and loss of
electrons (between metal and nonmetal this chapter)
2. covalent – which involves the sharing of
electrons (between nonmetals - chapter 6)
3. metallic bonds- which involves sharing of
electrons (between metals - chapter 6)
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Metals tend to lose electrons in reactions
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Remember, metals are usually found in
groups 1-12, some in 13-16, and the
lanthinides and actinides
metals usually have 3 or fewer outer
electrons and it is therefore easier for them
to lose electrons than to gain them
alkali metals are the most reactive metals
•
Nonmetals tend to gain electrons in
reactions
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remember, nonmetals are usually in groups
17 and 18, with some in groups 13-16
nonmetals tend to have 5 or more outer
electrons and it is therefore easier for them
to gain electrons than to lose them
halogens are the most reactive nonmetals
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Valence Electrons
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atoms of many elements become stable by
achieving the outer electron arrangements
of a noble gas
remember, we call these outer electrons
valence electrons
–
we can determine an atoms valence electrons from
the pattern on the periodic table, and from that we
can predict whether the atom will lose or gain
electrons
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group 1 – 1 outer electron (lose 1)
group 2 - 2 outer electrons (lose 2)
group 13 – 3 outer electrons (lose 3)
group 14 – 4 outer electrons (lose or gain 4)
group 15 – 5 outer electrons (gain 3)
group 16 – 6 outer electrons (gain 2)
group 17 – 7 outer electrons (gain 1)
group 18 – 8 outer electrons (already full, so they are nonreactive)
group 3-12, transition metals can have different
configurations but we predict 2 outer electrons for them
(lose 2) except for silver and gold for which we predict 1
(lose 1) unless we are told otherwise
–
When atoms gain or lose electrons they are
making their number of protons and electrons
unequal
•
this means the atoms will no longer remain
neutral
•
when atoms are not neutral they must have a
charge and we call charged atoms ions
•
if an atom loses electrons, it will have more
protons (+) than electrons (-) and its charge
will be positive
• we call positive ions “cations”
• their positive charge is equal to the
number of electrons they lose
• if an atom loses 1 electron it will have a +1
charge, if it loses 2 electrons it will have a +2
charge and so on
•
if an atom gains electrons, it will have more
electrons (-) than protons (+) and its charge
will be negative
• we call negative ions “anions”
• their negative charge will be equal to
the number of electrons they gain
• if an atom gains 1 electron its charge will be
–1, if it gains 2 electrons its charge will be –2,
and so on
•
so we can use the number of valence electrons to predict
how many electrons an atom will lose or gain and therefore
the charge it will take on
- group 1 – 1 outer electron (lose 1) = +1
– group 2 - 2 outer electrons (lose 2) = +2
– group 13 – 3 outer electrons (lose 3) = +3
– group 14 – 4 outer electrons (lose or gain 4) = + or - 4
– group 15 – 5 outer electrons (gain 3) = -3
– group 16 – 6 outer electrons (gain 2) = -2
– group 17 – 7 outer electrons (gain 1) = -1
– group 18 – 8 outer electrons (already full, so they are
non-reactive)
– group 3-12, transition metals can have different
configurations but we predict 2 outer electrons for them
(lose 2 = +2) except for silver and gold for which we
predict 1 (lose 1 = +1) unless we are told otherwise
»
transition metals can often form more than one ion, we predict the
ions as mentioned above unless we are told otherwise
↓
Predict +2,
except Au & Ag = +1
↓
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therefore, we can see that metals form cations
(+) by losing electrons and nonmetals form
anions (-) by gaining electrons
•
notice the nucleus is never changed in this
process
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label and learn the pattern for prediction ionic
charge on the periodic table and your
flashcards
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Atoms and Ions
–
When atoms lose and gain electrons to form ions,
they are still the same elements
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remember, the number of protons is what determines
the identity of an atom
just because an atom is charged, it still has the same
number of protons and therefore the same identity
Ions have different properties than atoms that
made them
Ions are different sized than atoms that made
them
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atoms that lose electrons form ions that are smaller
than their parents
atoms that gain electrons form ions that are larger than
their parents
Section 2, Ionic Bonding and Salts
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Because opposites attract, cations and anions
attract one another
– This is what happens when ionic bonds form
• the word salt actually means an ionic
compound that forms when a metal atom
or a positive radical (group of two or more
atoms acting as a single atom) replaces
the hydrogen of an acid
• examples of salts: sodium chloride,
potassium chloride, magnesium oxide etc.
• all of these salts are ionic compounds that
are electrically neutral
• anions actually attract several cations and
vice-versa, in this way many ions are pulled
together into a tightly packed structure
•this structure gives any salt its crystalline
structure
•the smallest crystal of table salt that you
can see still has more than a billion billion
sodium and chloride ions
•
adding and removing electrons requires energy
• remember ionization energy is the energy
required to remove an electron and electron
affinity is the energy required to add an electron
onto a neutral atom
• this is only a part of the process involved in
forming an ionic bond
•
more energy is released when ionic bonds
are formed than is required to make the
ions
• since energy is released the overall process is
exothermic and spontaneous, even though
some parts require energy to occur
• also, since energy is released when these bonds
are formed, then energy is required to break these
bonds
Ionic Compounds
• remember that cations are positive and
anions are negative, but when they come
together to form an ionic bond they have
no overall charge
– the ratio of the charges cancel each other out
• ionic compounds do not consist of
molecules
– remember that ionic compounds are formed
when anions and cations attract each other to
form a crystal
• elements in group 1 and 2 reacting with
groups 16 and 17 will almost always form
ionic compounds and not molecular
(covalent) compounds
• ionic bonds are very strong which gives ionic
compounds certain properties
– most ionic compounds have high melting and boiling points,
and they are usually solid at room temperature
– liquid and dissolved salts conduct electrical currents
• to conduct an electric current, a substance must satisfy 2 conditions
– they must contain charged particles
– those particles must be free to move
• ionic solids, such as salts, do not usually conduct electrical current
because the ions are not free to move
• however, when they melt or dissolve their ions are free to move and
are therefore excellent conductors
• there is a small class of ionic compounds that can allow charges to
move through their crystals
– ex. Zirconium oxide
– Salts are hard and brittle
– how to identify ionic compounds (characteristics)
• solid at room temperature
• hard and brittle
• high melting and boiling points
• good conductors of electrical current in the liquid
state (melted) or dissolved in water
• Salt Crystals
– the ions in salts form repeating patterns
• not all salts have the same crystal structure
• the crystals of all salts are made of simple repeating
units, called a crystal lattice
– crystal structures depend on the sizes and ratios of
ions
– Salts have ordered packing arrangements
– the smallest repeating unit in a crystal lattice pattern
is called a unit cell
Section 3, Names and Formulas of Ionic Compounds
• Naming Ionic Compounds
– naming salts (ionic compounds) is very easy
• simple ionic compounds made of just two
elements are called binary compounds, we
will learn about these first
• rules for naming simple ions
– cations use the same name as their parent atom
• ex.: K+ is the potassium ion, Zn+2 is the zinc ion
– when elements, such as transition metals, form
more than one type of ion, the ion name includes a
roman numeral to indicate its charge
• ex.: Copper has 2 common ions, one with a +1 cahrge
and one with a +2 charge; so, we name them copper (I)
and copper (II) with symbols Cu(I) and Cu (II)
– anions also get names from their elements, but we
change their endings to “ide”
• examples: Chlorine forms an anion with a -1 charge, its
symbol is Cl – and we call it a chloride ion. Oxygen
forms an ion with a -2 charge, its symbol is O-2 and we
call it an oxide ion.
• to name ionic compounds we use the names
of the ions present in the compound
– we write the name of the cation first followed by
the anion
– example: NaCl is sodium chloride. ZnS is zinc
sulfide. K2O is potassium oxide. CuCl2 is copper
(II) chloride. Al2S3 is aluminum sulfide.
• Writing Ionic Formulas
– remember, compounds have no overall
charge
• this means that the charges from the
cations and anions must cancel each other
out mathematically
• steps for writing formulas of binary compounds
– write the symbol and charges for the cation and anion
– find the lowest common multiple of the charges to see how
many of each ion you need to cancel the charges
– example: magnesium nitride is made of Mg+2 and N-3 , we
need to figure out how many of each ion we need to get the
charges to cancel. Find the least common multiple of 3 and
2 which is six, so we need to have 6 positive charges and 6
negative charges to balance out. How can we get 6
positives; each Mg has 2 positives, so we need 3 of them.
How can we get 6 negatives; each N has 3 negatives, so
we need 2 of them. We then write the formula to reflect how
many of each ion we have. Mg3N2 . Notice that we no
longer write the charges, and the numbers of each ion are
written as subscripts.
• remember, ionic compounds are made of crystal
lattices (not molecules), therefore, their formulas
show us the smallest ratio of charges needed to
be neutral
• there is a shortcut to writing ionic formulas; the
cross-over method
– we take the charge from the cation and write it as the
subscript of the anion and vice-versa
– If the subscripts can be reduced mathematically,
reduce them
– Ex.1
Mg+2 and Cl-1 ,
MgCl2
– Ex. 2
Mg+2 and O-2 , Mg2O2, MgO
Polyatomic Ions
• Up till now, we have been talking about simple
ions. This means ions made from one charged
atom.
– we also call these monatomic ions
• But there are ions that are made of more than
one atom. These are called polyatomic ions.
– polyatomic ions are a charged group of bonded atoms
that act together as one atom
– they usually act the same way simple ions do
– they can either be cations or anions
– see chart on page 178 for common polyatomic ions,
you will have to memorize some of these
• the names of polyatomic ions can be complicated
– many polyatomic ions contain oxygen; endings such as “ite”
and “ate” indicate the presence of oxygen
– many polyatomic ions differ only by the number of oxygen
atoms present
• if there are only 2 forms; the form with less oxygen ends in “ite”, the
form with more oxygen ends in “ate”
• if there are more than 2 forms; we can use the prefix “hypo” for the
least number of oxygen and “per” for the most
• examples: nitrate (NO3-) and nitrite (NO2-) ,or
• hypochlorite (ClO-) and chlorite (ClO2-) and chlorate (ClO3-) and
perchlorate (ClO4-)
– the prefix “thio” means replaced an oxygen with a sulfur
• example: sulfate (SO4-) and thiosulfate (S203-)
– most of these you just have to memorize (there symbols and
charges)
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Naming compounds with polyatomic ions
in them
– we name them the same as binary
compounds except if the compound ends in
the polyatomic ion (anion) we do not use
and “ide” ending, but rather keep the ending
of the polyatomic ion
– example: NH4Cl is ammonium chloride and
CaCO3 is calcium carbonate
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Writing a formula for a compound containing a
polyatomic ion
– we also write formulas for compounds with
polyatomic ions the same as we do for binary
compounds
– the only thing we need to remember is that the
polyatomic ion acts as one atom, so if we need to
add a subscript to a polyatomic ion we must put
parentheses around it to show that the subscript
goes with all of its atoms
– If you don’t use parentheses, you will be changing
the formula and therefore the compound
• example: to write the formula for magnesium
hydroxide (Mg+2 and OH-), we can see that
we need 2 OH’s to cancel the magnesium.
When we write the formula, we put
parentheses around the hydroxide and then
add the subscript so we keep the identities of
the ions the same
– Mg(OH)2 is correct, if we do not include the
parentheses we get MgOH2 which is incorrect
• *Read your chapter summary on page 182.