Transcript Chapter 8 - Ionic Compounds

Ionic
Compounds
and
Metals
Chapter 7
Vocabulary

Ch. 7.1
Chemical bond
 Cation
 Anion


Ch. 7.2
Ionic bond
 Ionic compound
 Binary compound
 Crystal lattice
 Electrolyte
 Lattice energy

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Objectives
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Define a chemical bond
Describe how ions form
Identify ionic bonding and the characteristics of
ionic compounds
Name and write formulas for
binary compounds
 Polyatomic ion containing compounds
 Compound with metals that have multiple oxidation
states
 Hydrates

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4
Forming Compounds

The Octet Rule: Atoms become stable by having 8
electrons in their outer energy level (2 for smaller
atoms)
 Similar arrangement of valence electrons for element
in the same group. (Ex: ns1 for alkali metals)
 Electron arrangements determines chemical
properties
 Presents a model of chemical stability
 When they have gotten 8 electrons they have
achieved NOBLE GAS CONFIGURATION
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Forming Compounds
atoms collide with enough energy
 outer electrons move to achieve a stable
octet of valence electrons (noble gas
configuration).
 atoms formed a new compound
 Total # of e- must remain the same.
Electrons are particles of matter.
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Forming Compounds

Transfer of Electrons

..
..
Na· + ·Cl:  [Na]+ + [:Cl:]˙˙
˙˙
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Lattice Energy

Ionic compounds consist of a crystal lattice of
positive and negative ions, which is a repeating
pattern of ions.
= Na+
= Cl-
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Lattice Energy

The lattice energy is the energy required to
separate the ions from their crystalline solid
state.
Increases with decreasing ionic size.

Increases with increasing ionic charge.
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Lattice
energy
decreases
as anion
gets
bigger!!
Lattice
energy is
large due to
+2 and -2
Charges;
decreases as
ions get
larger.
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Lattice Energy
Physical Properties of ionic compounds:
 The melting point increases as lattice energy
increases.
 Solutions are electrolytes, which conduct an
electric current.
 Force applied to crystals causes them to shatter.
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Vocabulary

Ch. 7.3
 Formula Unit
 Monoatomic ion
 Oxidation Number
 Polyatomic ion
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Names and Formulas

Rules for naming binary compounds
1. Name the cation first. This is usually a metal.
Use element name.
2. Then name the anion using the 1st syllable of
the of the element name and then end in –ide.
Example: Potassium Chloride (KCl)
Magnesium Oxide (MgO)
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Binary Compound Naming
Practice
Formula
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Name
NaF
Al2S3
CaO
Mg3N2
CaH2
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Predicting Charge on Ions
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Can you predict the charge on an ion depending
on where it is in the periodic table?
Yes, of course. Group numbers can help
predict many charges.
For groups 1 & 2 – they lose electrons and become
positive
 Group 13 – loses 3 electrons and are generally
positive
 Groups 15 – 17 gain electrons and become negative
in ions.
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Predicting Charge on Ions

The charge on a monoatomic ion is the
Oxidation number

Again, oxidation number for many elements in
the main group (Groups 1 & 2, 13-18) can be
predicted by looking at the group number.
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Rules For Writing Ionic
Compound Formulas
1. Write cation symbol first. (Same as we did in
naming them.)
2. Write anion symbol second.
3. Add subscripts such that the sum of the
charges is zero.
4. Write the simplest ratio of ions. This is the
formula unit.
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Writing Ionic Formula

Example: What is the formula for Aluminum
Oxide?
What is charge on aluminum when it is an ion?
 What is charge on oxygen when it is an ion (oxide)?
 How do we get the charges to balance?
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Method of “Cross-Multiplying”
What would be the subscripts
to make the Sum = 0?
Drop the signs and make
the numbers subscripts
Use it backward—knowing compound
Figure out charges
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Examples of Writing Formulas
(Binary Compounds)
Name of Compound
 Lithium oxide
Lithium ions and oxide ions
Li+
O2 Calcium chloride
Calcium ions and chloride ions
Ca2+
Cl
Formula of Compound
Magnesium nitride
magnesium ions and nitride ions
Mg2+
N322
Practice Writing Binary
Compound Formulas
Name of Compound
 Lithium iodide
 Beryllium chloride
 Calcium oxide
 Sodium oxide
 Strontium sulfide
 Calcium phosphide
Formula of Compound
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Compounds of Metals with Multiple
Oxidation States (Multivalent Metals)
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Includes metals in Groups 3 through 12 of the
periodic table
And the representative metals in Groups 13 &
14 (Commonly Sn & Pb, but also Ga, In, and Tl)
They have more than one
Oxidation State
 So they can form more than one type of
positive ion.
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Compounds of Multivalent Metals

For Example:
Copper can exist as Cu+ and Cu2+
Iron can exist as Fe2+ and Fe3+

Exceptions: Zn and Ag.
Zinc only forms a Zn2+ ion
 Ag only forms a Ag+ ion
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Compounds of Multivalent Metals


How do chemists distinguish the names of compounds
formed these metals?
We use Roman numerals in parenthesis after the name
of the element.
Copper Ion Chloride Ion
Formula
Name
Copper (I)
Chloride
Copper (II)
Chloride
Cu+
Cl-
CuCl
Cu2+
2Cl-
CuCl2
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Compounds of Multivalent Metals
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Compounds of Multivalent
Metals
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Formulas of Multivalent Metals
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How to write a formula containing one of these
metals. First, look at the name.
Example: if it is Manganese (III) it means that
it’s Mn3+
The roman numeral is the oxidation number
of the ion, not how many there are!!!!
Then, you can balance charges as before.
Example: MnCl3
Manganese is 3+ and chloride is always 1-.
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Writing Names of Multivalent Metal
Compounds
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If your given a formula, how do you figure out
the name?
Determine the charge on the metal ion.
Look at the negative ion.
Figure out what was needed to make the
compound neutral
Example: FeCl3
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Compounds of Multivalent Metals
Given name or formula
 Iron(II) oxide

MnF3

Nickel(II) chloride

PbS2
Determine formula or name
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What is the oxidation number of
Zinc when it’s an ion?
+
1
+
2
a)
or
+
b) 2 always
+
c) 1 always
d) 1 or 2
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What is the formula for iron(II)
oxide?
a)
b)
c)
d)
Fe2O
FeO2
Fe2O2
FeO
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What is the correct name of
MnF3?
a)
b)
c)
d)
Manganese (I) fluoride
Magnesium (I) fluoride
Manganese (III) fluoride
Magnesium (III) fluoride
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What is the correct formula for
nickel(II) chloride?
a)
b)
c)
d)
Ni2Cl
NiC2
NiCl2
Ni2Cl2
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What is the name of PbS2?
a)
b)
c)
d)
Lead (IV) sulfide
Lead (I) sulfide
Lead (II) sulfide
Lead (III) sulfide
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Polyatomic Ions
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Ions can contain more than one element.
An ion with two or more different elements is a
Polyatomic Ion
In polyatomic ions, the atoms are covalently
bonded and the atoms share electrons.
Individual atoms have no charge, but the group
has an overall charge.
Example: SO42- (Sulfate)
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Common Polyatomic Ions
Name of Ion
Ammonium
Hydrogen carbonate
(bicarbonate)
Hydrogen sulfate
Acetate
Nitrite
Nitrate
Cyanide
Hydroxide
Formula
NH4+
HCO3-
Charge
1+
1-
HSO4C2H3O2NO2NO3CNOH-
111111-
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Common Polyatomic Ions
Name of Ion
Dihydrogen phosphate
Permanganate
Carbonate
Sulfate
Sulfite
Oxalate
Monohydrogen phosphate
Dichromate
Phosphate
Formula
H2PO4MnO4CO32SO42SO32C2O42HPO42Cr2O72PO43-
Charge
112222223- 39
Compounds Containing Polyatomic
Ions
1. When there is more than one polyatomic ion,
treat it as if it were a single ion by keeping it
together as a unit using parenthesis.
2. Write a subscript outside the parenthesis to show
how many units there are of polyatomic ion.
3. Remember that the sum of the charges must
equal zero.
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Examples of Writing Formulas
(w/Polyatomic Ions)
Name of Compound
 Calcium nitrate
Calcium ions and nitrate ions
Ca2+
NO31 Sodium carbonate
Sodium ions and carbonate ions
Na+
CO32 Lead(II) Phosphate
lead 2+ ions and phosphate ions
Pb2+
PO43-
Formula of Compound
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Naming Polyatomic Compounds
from formulas
1. Positive ion (usually a metal) is named
first as before.
2. Negative polyatomic ion name as listed in
the chart is second.
Note: If you see more than one element after
the metal the rest is a polyatomic ion.
Look for that group in the chart.
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Polyatomic Ion Practice
Given name or formula: Determine formula or name
 NaHCO3
 Lithium Acetate
 Ca(CN)2
 Copper(II) Hydroxide
 K2Cr2O7
 Fe(HCO3)3
 Ammonium Chloride
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What is the formula for
zinc arsenate?



Note: You have to type this in.
Hit the ‘Shift’ key (lower right) to get capital
letters.
If you need parenthesis, hit sym 4 for ‘(‘ and sym
5 for ‘)’.
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Ch. 7.4 Vocabulary

Ch. 10.5

Hydrate (page 351)

Ch. 7.4
Electron Sea model
 Delocalized electron
 Metallic Bond
 Alloy
 Interstitial alloy
 Substitutional Alloy

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Compounds of Hydrates (Ch. 10.5)
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Hydrate – when there is a specific amount of
water in the crystal of a compound.
The water molecules are part of the crystal
structure of the ionic compound.
Many compounds become hydrates by
absorbing water from the air. The water then
becomes part of its structure.
Many of them are used as drying agents
(desiccants).
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Compounds of Hydrates

When writing a formula for a hydrated
compound use a dot (•) followed by the number
of water molecules.
Example: CaCl2•2H2O

This means that this hydrated chloride
compound has 2 molecules of water for each
formula unit of calcium chloride.
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CaCl2•2H2O
Water
-
Cl-
Cl-
Ca+2
-
Water
Compounds of Hydrates
CaCl2•2H2O
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This means that this hydrated chloride compound has 2
molecules of water for each formula unit of calcium
chloride.
You would call the compound
Calcium chloride dihydrate.
Rule: Follow the regular name of the compound
with the word “hydrate.” The prefix before
“hydrate” tells you how many water molecules
there are.
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Prefixes for Hydrates
Molecules of Water
1
2
3
4
5
6
7
8
9
10
Prefix
MonoDiTriTetraPentaHexaHeptaOctaNonaDeca-
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Hydrate Compound Practice
Given Name or Formula

MgSO3•6H2O

Strontium oxalate
monohydrate
Ni3(PO4)2•7H2O
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Determine the Formula or Name
Lead(II) acetate
trihydrate
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Compounds of Hydrates
Hydrates can also lose water molecules,
usually by heating.
 They may even have a different color than
the hydrated compound.


When they lose their water molecules, they
become anhydrous (without water).
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Ch. 7.4 - Metal Bonding – Electron
Sea Model
 There are not enough electrons for the
metal atoms to be covalently bonded
to each other.
 We
use a delocalized model for
electrons in a metal
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Delocalized Electrons in a Metal


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

The metal nuclei are seen to exist in a sea of
electrons.
No electrons are localized between any two
metal atoms.
Therefore, the electrons can flow freely through
the metal.
Without any definite bonds, the metals are easy
to deform (and are malleable and ductile).
The more delocalized e-, the harder and stronger
the metal.
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Delocalized Model for Electrons in a
Metal
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Melting Point
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Malleability
Metals, above, compared to ionic compounds, below.
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Alloys


Alloys are homogeneous mixtures, metal
solutions, really.
Substitutional alloys:
atoms must have similar atomic size,
 elements must have similar bonding characteristics.


Interstitial alloys:
one element much smaller than the other in order to
fit into the interstitial sites, e.g. a nonmetal.
 The alloy is much stronger than the pure metal
(increased bonding between nonmetal and metal).
 Example steel (contains up to 3% carbon).
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Alloys
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Common Alloys
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