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Unit 5 Chemical Compounds Chapter 6

Objective 1 ď Ž

ď Ž

Describe how a compound differs from its component elements.

A compound is two or more simple substances and can be broken down into simple substances.

Objective 2 ď Ž

ď Ž

Explain what a chemical formula represents.

It shows what elements a compound contains.

Objective 3 ď Ž

ď Ž

State a reason why chemical bonding occurs.

To obtain a complete set of valence electrons.

Objective 4 ď Ž

Describe the three types of chemical bonds.

There are three types of bonds   

Ionic Covalent Metallic

Bonding  



How can you tell what type of bond is going to form?

You use what is called an electronegativity table. It is on the back of Review Sheet, Part 1.

Electronegativity- is the tendency for an atom to attract a shared pair of electrons.

Objective 5 Use electronegativity differences to determine the chemical bond that will form.

Look at an electronegativity table. What do you notice about the numbers?

Electronegativity Table ď Ž ď Ž

Which element has the highest value What does this mean?

Electronegativity Difference     

Table 1 Look up Li write down it’s electronegativity value.

Look up Cl write down it’s electronegativity value.

Find the difference and record. (Subtract the smaller number from the bigger number) Repeat until you have finished all of the first two columns.

Elements Li-Cl C-H H-H N-H C-H K-Br Na-S

EN Difference I or C? 0.98,3.16 2.18 Don’t do 2.55, 2.20 0.35 this column yet

Third column: Ionic or Covalent? 





If the difference is less than 1.67 the compound will be covalent. That compound would share electrons.

If the difference is greater than 1.67, the compound will be ionic. That compound would form oppositely charged ions that would attract to each other.

Complete the rest of Table 1.

Electronegativity Difference Covalent 0

<

Ionic 1.67

Objective 6 Recognize stable electron configurations.

Ionic Bonding 

  

All elements combine to obtain a stable electron configuration.

Metals will lose electrons.

Nonmetals will gain electrons.

Please highlight this in your notes packet.

Combining Atoms 



The first two atom we want to combine are Lithium + chlorine Let’s got through the reasoning to determine the ionic compound formed.

Li-Cl  

 

Lithium is 1s22s1.

If lithium could lose one electron it would be like the noble gas helium.

Chlorine is 1s22s22p63s23p5.

When lithium loses it’s one electron it gives it to chlorine.

Li-Cl  



So lithium will form a positive ion or cation with a charge of +1.

When chlorine receives the one electron it will form a negative ion or anion with a charge of –1.

The attraction of the oppositely charged ions will then form the compound LiCl, lithium chloride.

Li-Cl ď Ž

We can represent this with what is called a dot diagram.

Combining Atoms Lithium + Chlorine Copy the next slide.

ď&#x192;

Compound formed Lithium chloride

[ ]

Li + Cl

Li + Cl

+

LiCl

-

K-Br Potassium

+

Bromine

ď&#x192;

Potassium bromide

[ ]

K + Br

K + Br

+

KBr

-

Mg-Cl Magnesium

+

Chlorine

ď&#x192;

Magnesium chloride

Mg + Cl Cl

[ ]

Mg +2 Cl +2

MgCl2

-

Na-Cl Sodium

+

Chlorine ď&#x192;

Sodium chloride

Na + Cl

[ ]

Na + Cl +

NaCl

-

Review Sheet, part I 

 

In this homework be sure to include a word equation and formulas.

Show electron transfer.

Place only negative ions in brackets.

Draw the dot diagram for:  

Sulfur Phosphorus

Objective 7 Describe how an ionic bond forms and how ionization energy affects the process.

Conclusions: ď Ž ď Ž

Ionic bonding involves the transfer of one or more electrons from the metal to a nonmetal.

Energy changes are involved. a.

b.

ď Ž

Ionization energy-the amount of energy needed to remove an electron from an atom.

Electron affinity-the energy change when an electron is added.

The element that loses electrons will form a positive ion. +1, +2, +3 for example.

Conclusions, cont. ď Ž

ď Ž

The alkali metals, alkaline earth metals and aluminum will form positive ions equal to the group number.

The element that receives the transferred electron(s) will form a negative ion. -1, -2 or -3 for example.

Electrostatic force ď Ž

ď Ž

The attraction of these oppositely charged ions is what forms the ionic bond and a compound has been formed.

This is the electrostatic force.

Properties of Ionic Compounds 1. High melting point.

2. Nonconductor in the solid state.

3. Conductor when melted or dissolved in solution. Ionic compounds are electrolytesconductors of electricity in aqueous solution.

4. Brittle, or shatter when hit with a hammer.

Objective 8 Write the chemical formula for ionic compounds.

Ionic Compounds 

 

Binary compounds -a compound made from only two elements.

Look at your List of Common Ions.

An ionic compound can be made from any positive ion from the front combined with any negative ion on the back.

Ionic Compounds Table 2 Letâ&#x20AC;&#x2122;s combine Li+ with IThe sum of the charges in a compound is zero. So one of each will work.

LiI

Ionic Compounds Li+

ClO32. LiClO3

ClO23. LiClO2

Ionic Compounds Li+

C2O4-2

need 2

need 1

+2

-2

Formula 4. Li2C2O4

Ionic Compounds Li+ C2O4-2 Formula need 2 need 1 4. Li2C2O4 +2 -2 SO3-2 need 2 need 1 5. Li2SO3 PO4-3 need 3 need 1 +3 -3 6. Li3PO4

Ionic Compounds Li+

SO3-2

need 2

need 1

PO4-3 need 3 +3 -3

need 1 6. Li3PO4

Formula 5. Li2SO3

Ionic Compounds Cr+2 Ineed 2 7. CrI2 ClO3Need 1 +2 -2

need 2 8. Cr(ClO3)2

9. Cr(ClO2)2 Must use parenthesis on polyatomic ion.

Ionic Compounds Cr+2 C2O4-2 need 1 need 1 10. CrC2O4 SO3-2 need 1 need 1 11. CrSO3 PO4-3 need 3 need 2 12. Cr3(PO4)2 +6 -6 Be sure to use parenthesis only on the polyatomic ion. Notice there is no parenthesis on the Cr.

Ionic Bonding ď Ž

Continue to complete the rest of Table 2 and check your work on the slides that follow.

Table 2 14.

KI KClO3

15.

KClO2

16.

K2C2O4

17.

K2SO3

18.

K3PO4 no parenthesis when there is just one.

13.

Table 2 19.

AlI3

20.

Al(ClO3)3

21.

Al(ClO2)3

22.

Al2(C2O4)3

23.

Al2(SO3)3

24.

AlPO4

Table 2 25.

ZnI2

26.

Zn(ClO3)2

27.

Zn(ClO2)2

28.

ZnC2O4

29.

ZnSO3

30.

Zn3(PO4)2

Summary 





The positive ion is always written first in the compound.

There will be no charges written in the compound. Do not do this Al+3 PO4 Be sure you have used parenthesis appropriately.

Complete electron transfer for ď Ž

Aluminum-oxygen being combined.

Lab Ionic Compounds. ď Ž

ď Ž

Predict the formula for ionic compounds given the combining ions.

Use the foam ion cutouts to help you to write the formula for the ionic compounds.

Objective 9 Name ionic compounds.

Letâ&#x20AC;&#x2122;s name the compounds we made in Table 2.

Rules for Naming 1. Name the positive ion first.

2. If the positive ion has more than one charge, use a Roman numeral in the name. For example there is Cr+2 and Cr+3.

So you have chromium (II) and chromium (III).

3. Name the negative ion second.

Table 2: use your ion sheetâ&#x20AC;Ś 1.

2.

3.

4.

5.

6.

Lithium iodide Lithium chlorate Lithium chlorite Lithium oxalate Lithium sulfite Lithium phosphate

Table 2 Use Roman numeral with metals that have more than one charge.

7.

Chromium (II) iodide 8.

Chromium (II) chlorate 9.

Chromium (II) chlorite 10. Chromium (II) oxalate 11. Chromium (II) sulfite 12. Chromium (II) phosphate

Table 2 ď Ž

Got the idea. No need to name the rest unless you are having trouble.

Table 2 13.

14.

15.

16.

17.

18.

Potassium iodide Potassium chlorate Potassium chlorite Potassium oxalate Potassium sulfite Potassium phosphate

Table 2 19.

20.

21.

22.

23.

24.

Aluminum iodide Aluminum chlorate Aluminum chlorite Aluminum oxalate Aluminum sulfite Aluminum phosphate

Table 2 25.

26.

27.

28.

29.

30.

Zinc iodide Zinc chlorate Zinc chlorite Zinc oxalate Zinc sulfite Zinc phosphate

Objective 10 



Given the name, write the chemical formula.

Let’s go the other way!!!

Given the name, write the formula Name 1. Copper (I) chloride

2. Copper (II) chloride

Ions Formula Combined

Given the name, write the formula Name

Ions Formula Combined

1. Copper (I) chloride

Cu+ Cl-

2. Copper (II) chloride Cu+2 Cl-

Given the name, write the formula Name

Ions Formula Combined

1. Copper (I) chloride

Cu+ Cl-

2. Copper (II) chloride Cu+2 Cl-

CuCl

CuCl2

Name

3. Calcium hydroxide

4. Calcium carbonate

5. Ammonium sulfate

Ions Combined

Formula

Name

Ions Combined

3. Calcium hydroxide

Ca+2

OH-

4. Calcium carbonate

Ca+2

CO3-2

5. Ammonium sulfate

NH4+ SO4-2

Formula

Name

Ions Combined

Formula

3. Calcium hydroxide

Ca+2

OH-

Ca(OH)2

4. Calcium carbonate

Ca+2

CO3-2

CaCO3

5. Ammonium sulfate

NH4+ SO4-2

(NH4)2SO4

Name 6. Silver Sulfate

7. Lithium fluoride

8. Iron (II) sulfate

Ions Combined Formula

Name

Ions Combined Formula

6. Silver Sulfate

Ag+ SO4-2

7. Lithium fluoride

Li+

8. Iron (II) sulfate

Fe+2 SO4-2

F-

Name

Ions Combined Formula

6. Silver Sulfate

Ag+ SO4-2

Ag2SO4

7. Lithium fluoride

Li+

LiF

8. Iron (II) sulfate

Fe+2 SO4-2

F-

FeSO4

Name 9. Sodium oxide

10. Sodium nitride

Ions Combined Formula

Name

Ions Combined Formula

9. Sodium oxide

Na+ O-2

10. Sodium nitride

Na+ N-3

Name

Ions Combined Formula

9. Sodium oxide

Na+ O-2

Na2O

10. Sodium nitride

Na+ N-3

Na3N

Lab Ionic Compounds Homework: WS I.

Objective 11 ď Ž

Describe how covalent bonds form and the attractions that keep atoms together in molecules.

Objective 12 ď Ž

Draw dot diagrams for covalent compounds.

 



C-H In covalent bonding there is the sharing of one or more electron pairs. There is NO ELECTRON TRANSFER! The compound that forms stays together because the atoms that are bonding share the electrons to help obtain a stable electron configuration.

C-H   



Carbon is 1s2 2s2 2p2 it has four valence electrons.

Hydrogen is 1s1 and has one valence electron.

Notice that if you put them together to make CH4, carbon will have 8 electrons around it and is stable and each hydrogen will have one pair. That fills the 1s level for hydrogen and it too is stable.

Continue with the remaining covalent examples.

CH4

H2

NH3

Na2S  

Ionic, not covalent. It is an exception.

Draw showing electron transfer.

PH3

I2

CO   

If single bonds do not work, try double.

If double fails, try triple.

Never double bond to a halogen.

CO2 ď Ž

Note: two double bonds formed

O2 яБо

Double bond

SCl2

N2 ď Ž

Triple bond

Covalent Bonds    

Single bond- the sharing of one electron pair.

Double bond- the sharing of two electron pairs.

Triple bond- the sharing of three electron pairs.

Covalent bonds occur to help each element attain a full energy level (H – 2, as opposed to an octet

Hydrocarbons 

We will make the following compounds: C2H6, C2H4, and C2H2



Always combine the carbon to carbon first Then attach the hydrogens.



C 2H 6

C 2H 4 яБо

Double bond

C 2H 2 

Triple bond ≡

Lab Covalent Compounds. ď Ž

ď Ž

Use the bag of atoms and plastic dots to help you to form the covalent compounds.

Draw the dot diagram, determine if the bond is polar or nonpolar and if the molecule is polar or nonpolar.

Objective 14 ď Ž

Name and determine the chemical formulas for molecular compounds.

Naming Covalent Compounds.  



Typically covalent compounds are composed of two or more nonmetals combined.

To name the compound, say the name of the first element and use Greek prefixes to indicate how many of each element you have.

The second element ends in –ide.

The Greek prefixes must be memorized.

1 Mono2 Di3 Tri4 Tetra5 Penta-

6 hexa7 Hepta8 Octa9 Nona10 Deca-

Greek prefixes Never use mono on the first word, always use mono on the second!

Compound Name CO carbon monoxide CO2 carbon dioxide ď Ž

SO2 sulfur dioxide N2H4 dinitrogen tetrahydride SO3 sulfur trioxide P2O5 diphosphorus pentoxide (drop the a in penta)

You try: N2O3 PCl3 SiO2 CCl4 I2O7 BeH2

Answers Dinitrogen trioxide Phosphorus trichloride Silicon dioxide Carbon tetrachloride Diiodine heptoxide (drop the a) Beryllium dihydride

Properties of Covalent Compounds 1.

2.

3.

ď Ž

Low melting point.

Formed by the sharing of electron pairs.

Nonconductor in the solid state, when dissolved in water or when melted. Covalent compounds are non-electrolytes.

Volatile (have an odor or forms fumes).

Example: Isopropyl alcohol (rubbing alcohol), mothballs.

Objective 15 ď Ž

Describe the structure of bonds in metals.

Metallic Bonding 





Metal atoms achieve stable electron configurations by losing electrons.

But what happens if there are no nonmetal atoms available to accept the electrons?

The metal atoms become cations surrounded by a “pool of shared electrons.”

Metallic Bonding 

 

A metallic bond is created by the attraction between the positive metal cation to the electrons lost.

These electrons are often referred to as “delocalized electrons.”

The metal is still neutral in charge, in effect, because the number of protons and electrons are still equal.

Metallic Bonding: The Electron Sea Model

Objective 16 ď Ž

Relate the properties of metals to their structure.

Properties of Metals 1.

2.

3.

4.

Conductor of electricity and heat.

Ductile Malleable Luster

Objective 17 ď Ž

Distinguish between ionic, covalent and metallic bonding. How the bonds are formed, the forces that hold them together and properties.

Summary: Types of Bonds 

Ionic    

Metal and nonmetal Large electronegativity difference > 1.67 Forms fomula units Transfer of electrons held together by electrostatic attraction

Properties of Ionic Compounds  

   

Solids are nonconductors Liquids, gases and aqueous solutions are conductors.

They are electrolytes High melting point and boiling point Low volatility Low vapor pressure

Summary: Types of Bonds 

Covalent    

Two nonmetals combined Small electronegativity difference < 1.67 Forms molecules Sharing of one or more electron pairs

Properties of Covalent Compounds     

Nonconductors of electricity Non-electrolytes Low melting and boiling points Volatile High vapor pressure

Summary: Types of Bonds 

Metallic  



Composed of metal atoms Delocalized electrons are free electrons among the cations Electron Sea Model

Properties of Metallic bonding     

Conductors of heat and electricity High melting and boiling point Malleable Ductile Luster

The end.