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What is an ionic compound?. Composed of a metal and a nonmetal Electrically non-conductive as a solid Conductive as molten liquids or in solution. Ions. Cation : A positive ion Mg 2+ , NH 4 + Anion : A negative ion Cl - , SO 4 2 -

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what is an ionic compound
What is an ionic compound?
  • Composed of a metal and a nonmetal
  • Electrically non-conductive as a solid
  • Conductive as molten liquids or in solution
slide2
Ions
  • Cation: A positive ion
  • Mg2+, NH4+
  • Anion: A negative ion
  • Cl-, SO42-
  • Ionic Bonding: Force of attraction between oppositely charged ions.
formation of an ionic compound
Formation of an ionic compound
  • Combines a metal and a nonmetal through the transfer of electrons
  • Example: A compound made from potassium and chlorine
slide4

Potassium has one valence electron and tends to lose it to become a cation with a charge of +1

K

Cl

Chlorine has seven valence electrons and tends to gain one to become an anion with a charge of -1

slide5

K

Cl

slide6

Potassium, being an alkali metal, has an oxidation number of +1

-

+

K

Cl

Chlorine, being a halogen, has an oxidation number of -1

slide7

When these two ions combine to form an ionic compound, they will combine in such a way that the overall charge of the compound is zero. In this case, the smallest ratio that adds up to zero is one potassium ion and one chloride ion.

-

+

K

Cl

slide8

When these two ions combine to form an ionic compound, they will combine in such a way that the overall charge of the compound is zero. In this case, the smallest ratio that adds up to zero is one potassium ion and one chloride ion.

KCl

Note that the metal ion is ALWAYS written first, and that when there is only one ion of an element, it is not necessary to place a subscript “1” next to that element.

slide9

Calcium has two valence electron and tends to lose both of them to become a cation with a charge of +2

Ca

Cl

Chlorine has seven valence electrons and tends to gain one to become an anion with a charge of -1

slide10

Since calcium has two electrons to give, but chlorine can only accept one, those electrons must go to two separate chlorine atoms.

Cl

Ca

Cl

slide11

-

Calcium, being an alkaline earth metal, has an oxidation number of +2

Cl

2+

Ca

-

Cl

Chlorine, being a halogen, has an oxidation number of -1

slide12

When these two ions combine to form an ionic compound, they will combine in such a way that the overall charge of the compound is zero. In this case, the smallest ratio that adds up to zero is one calcium ion and two chloride ions.

-

2+

Ca

Cl

slide13

When these two ions combine to form an ionic compound, they will combine in such a way that the overall charge of the compound is zero. In this case, the smallest ratio that adds up to zero is one calcium ion and two chloride ions.

CaCl2

Note that the metal ion is ALWAYS written first, and that when there is only one ion of an element, it is not necessary to place a subscript “1” next to that element.

slide14

Hopefully by this point you can see that it is not necessary to draw out the dot structures before writing the formula for an ionic compound. The dot structures provide us with the oxidation number, and the oxidation numbers determine the ratio of ions in the compound.

But since we can find the oxidation numbers of the representative elements from the periodic table, we can skip the step of drawing out the dot structures.

beryllium and oxygen
Beryllium and oxygen
  • Beryllium is an alkaline earth metal and has an oxidation number of +2 (it loses two electrons)
  • Oxygen has an oxidation number of -2 (it gains two electrons)
slide16

Be2+

O2-

The smallest ratio that adds up to zero is one beryllium ion to one oxide ion:

BeO

lithium and nitrogen
Lithium and nitrogen
  • Lithium is an alkali metal and has an oxidation number of +1 (it loses one electron)
  • Nitrogen has an oxidation number of -3 (it gains three electrons)
slide18

Li+

N3-

The smallest ratio that adds up to zero is three lithium ions to one nitride ion:

Li3N

aluminum and sulfur
Aluminum and sulfur
  • Aluminum has an oxidation number of +3 (it loses three electrons)
  • Sulfur has an oxidation number of -2 (it gains two electrons)
slide20

Al3+

S2-

The smallest ratio that adds up to zero is two aluminum ions to two sulfide ions:

Al3S2

writing ionic compound formulas
Writing Ionic Compound Formulas

Example: Barium nitrate

1. Write the formulas for the cation and anion, including CHARGES!

( )

2. Check to see if charges are balanced.

Ba2+

NO3-

2

Not balanced!

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

writing ionic compound formulas1
Writing Ionic Compound Formulas

Example: Ammonium sulfate

1. Write the formulas for the cation and anion, including CHARGES!

( )

NH4+

SO42-

2. Check to see if charges are balanced.

2

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

writing ionic compound formulas2
Writing Ionic Compound Formulas

Example: Iron(III) chloride

1. Write the formulas for the cation and anion, including CHARGES!

Fe3+

Cl-

2. Check to see if charges are balanced.

3

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

writing ionic compound formulas3
Writing Ionic Compound Formulas

Example: Aluminum sulfide

1. Write the formulas for the cation and anion, including CHARGES!

2. Check to see if charges are balanced.

Al3+

S2-

2

3

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

writing ionic compound formulas4
Writing Ionic Compound Formulas

Example: Magnesium carbonate

1. Write the formulas for the cation and anion, including CHARGES!

Mg2+

CO32-

2. Check to see if charges are balanced.

They are balanced!

writing ionic compound formulas5
Writing Ionic Compound Formulas

Example: Zinc hydroxide

1. Write the formulas for the cation and anion, including CHARGES!

( )

2. Check to see if charges are balanced.

Zn2+

OH-

2

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

writing ionic compound formulas6
Writing Ionic Compound Formulas

Example: Aluminum phosphate

1. Write the formulas for the cation and anion, including CHARGES!

2. Check to see if charges are balanced.

Al3+

PO43-

They ARE balanced!

naming ionic compounds
Naming Ionic Compounds
  • 1. Cation first, then anion
  • 2. Monatomic cation = name of the element
  • Ca2+ = calciumion
  • 3. Monatomic anion = root + -ide
  • Cl- = chloride
  • CaCl2= calcium chloride
naming ionic compounds continued
Naming Ionic Compounds(continued)

Metals with multiple oxidation states

  • - some metal forms more than one cation
  • - use Roman numeralin name
  • PbCl2
  • Pb2+is cation
  • PbCl2 = lead(II) chloride
naming ionic compounds examples
Naming Ionic Compounds:Examples

Na2SO4

sodium sulfate

Fe(NO3)2

iron (II) nitrate

aluminum chloride

AlCl3

Mg3N2

magnesium nitride

(NH4)3PO4

ammonium phosphate

formulas
Formulas

Formulas for ionic compounds are ALWAYS empirical (lowest whole number ratio).

Examples:

NaCl

MgCl2

Al2(SO4)3

K2CO3

covalent bonding
CovalentBonding
  • Metals can only give electrons, and produce positively charged ions.
  • Nonmetals can gain or lose electrons to complete their octets.
  • The chemical bond formed from electron sharing between atoms is called the covalent bond.
covalent bonding in h 2
Covalent Bonding in H2
  • Atoms of hydrogen ( 1H : 1) have one valence electron in their first electron shell.
  • Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In H2 molecule, which hydrogen atom will donate one electron to which one?
  • But both hydrogen atoms can attain helium like electron arrangement (2He: 2) by sharing their valence electrons.
  • Because the hydrogen molecule is a combination of equally matched atoms, the atoms will share each other\'s single electron, forming one covalent bond.
slide36
When two hydrogen atoms are near each other, the two electrons and the two protons repel each other. But there is an attraction between the proton of one hydrogen atom and the electron of the other. This attraction force brings the atoms closer together.
  • When the two nuclei are kept a certain distance away from each other, the attractive forces become grater than the repulsive forces, and a stable chemical bond forms.
slide37
In this situation, the distance between two nuclei is called bond length or bond distance.
  • The hydrogen atoms in H2 molecule is in a less energetic and therefore more stable state than hydrogen atoms alone.
  • During the formation of the bond, some amount of energy is released. This energy is called the bond energy. As the magnitude of bond energy increases the strength of the bond increases.
slide39
The electrons shared between two identical nuclei such as in H2, Cl2, F2, O2 or N2 are located at equal distances between two nuclei. Neither of the atoms gains or loses electrons, or the positive charge center and the negative charge center in the molecule coincide.
  • Molecules have no positive or negative poles; is called nonpolar covalent bonding. H2, Cl2, F2, O2 , P4, S8 ... all have nonpolar covalent bonds.
covalent bonding in hci
Covalent Bonding in HCI
  • To complete their valence shells, hydrogen and chlorine both need one electron.
  • Since both have high tendencies to accept electrons, we do not expect one of them to donate one electron to the other.
  • But they may share a pair of electrons
slide41
Chlorine has a much greater attraction for the shared electron pair than hydrogen atom.
  • Such bonds are called polar covalent bonds.
  • The unequal sharing of electrons in a bond leads to what is referred to as a dipole.
slide42
The presence of polar covalent bonds does not necessarily mean the molecule will be polar. Some molecules contain two or more dipoles that cancel to give a nonpolar molecule.
  • For example, in CO2 the two oxygen are attached to carbon by polar covalent bonds. The oxygen atoms having greater tendency to attract electrons pull more strongly the shared electrons that form the double covalent bonds. Thus, two polar double covalent bonds are present in CO2, but because of the symmetry of the dipoles the molecule itself is nonpolar.
slide43
When determining whether a molecule is polar or nonpolar, it is important to consider the geometry of the molecule.
  • The tendency of some of the atoms to attract shared electrons is as follows:
  • F>O>CI>Br>N>S>I>C>H
  • Whenever we have diatomic molecules consisting of two different elements, the molecule is generally polar.
  • For example, the polarity order in the molecules HF, HCI, HBr, and HI is HF > HCI > HBr > HI.
covalent bonding in h 2 o
Covalent Bonding in H2O
  • Oxygen with the electron configuration of 8O: 2-6 needs two electrons to have eight electrons in its outer most shell.
  • So two hydrogen atoms and one oxygen atom must combine to form water.
  • Each bond between O and H atoms is a polar covalent bond. Water is a bent molecule.
  • In water, the polar covalent bonds lead to dipoles in which the centers of positive and negative charge do not coincide. This makes water a polar molecule.
naming nonmetal nonmetal compounds
Naming Nonmetal - Nonmetal Compounds
  • When two nonmetallic elements combine, no ions are formed. However it is generally possible to consider one of the two elements to be the more positive.
  • Both in writing the name and the formula of a binary compound, the more positive element appears first, followed by the more nonmetallic one.
slide46
In naming these compounds use the following pattern.
  • Number of the first nonmetal (in Latin), name of the first nonmetal, number of the second nonmetal (in Latin), ionic name of the second nonmetal.
naming binary covalent compounds examples
Naming Binary Covalent Compounds:Examples

N2S4

dinitrogen tetrasulfide

NI3

nitrogen triiodide

XeF6

xenon hexafluoride

CCl4

carbon tetrachloride

P2O5

diphosphorus pentoxide

SO3

sulfur trioxide

naming compounds practice

SiF4

silicon tetrafluoride

Naming Compounds: Practice

two nonmetals  covalent  use prefixes

Na2CO3

sodium carbonate

metal present  ionic  no prefixes

Na  group I  no Roman numeral

N2O

dinitrogen monoxide

two nonmetals  covalent  use prefixes

K2O

potassium oxide

metal present  ionic  no prefixes

K  group I  no Roman numeral

Cu3PO4

copper (I) phosphate

metal present  ionic  no prefixes

Cu  not group I, II, etc.  add Roman numeral (PO4 is 3-, each Cu must be 1+)

CoI3

cobalt (III) iodide

metal present  ionic  no prefixes

Co  not group I, II, etc.  add Roman numeral (I is 1-, total is 3-, Co must be 3+)

PI3

phosphorus trioxide

two nonmetals  covalent  use prefixes

NH4Cl

potassium oxide

NH4 polyatomic ion present  ionic  no prefixes

writing formulas practice
Writing Formulas: Practice

carbon tetrafluoride

CF4

prefixes  covalent  prefixes indicate subscripts

Na3PO4

sodium phosphate

metal  ionic  balance charges  3 Na1+ needed for 1 PO43-

copper (I) sulfate

Cu2SO4

metal present  ionic  balance charges 2 Cu1+ needed for 1 SO42-

aluminum sulfide

Al2S3

metal present  ionic  balance charges 2 Al3+ needed for 3 S2-

dinitrogen pentoxide

N2O5

prefixes  covalent  prefixes indicate subscripts

ammonium nitrate

NH4NO3

polyatomic ion present  ionic  balance charges 

1 NH41+ needed for 1 NO31-

lead (IV) oxide

PbO2

metal present  ionic  balance charges 1 Pb4+ needed for 2 O2-

iron (III) carbonate

Fe2(CO3)3

metal present  ionic  balance charges 2 Fe3+ needed for 3 CO32-

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