Chapter 20: Molecules and Compounds

Chemical Bonds

A chemical compound occurs when 2 different elements bond together so that both become stable.

• Chemical bonds form to lower the energy of the system. • Components of the system become more stable by forming bonds. • Everything wants to be more stable - its easier to lie down than to stand up. • Bonding is nature's way of allowing the elements to lie down.

Why do atoms form chemical bonds?

Atom Stability

* In order for atoms to be stable, they need to have their outer energy level completely filled.

* In most cases, atoms need to have 8 valence electrons in their outer energy level (8 is great).

* Atoms with 8 valence electrons are said to have octets.

The Noble Gases (group 18) naturally have octets and are said to be chemically stable. All other atoms must either lose or gain electrons to become stable.

There are exceptions to the octet rule: atoms with only one energy level need only two electrons to be stable, ex. hydrogen and helium.

Electron Losers & Gainers Atoms want to become stable by doing the

least amount of work possible.

Consider the following analogy:You are sitting at a table all by yourself . 7 of your friends are sitting at the table next to you. Will it be less work for you to join your friendsor for all 7 of your friends to join you? Of course, it takes less work for you to join your friends.

Who gains, who loses?? Is it easier for group one atoms to lose their 1

valence electron or to gain 7 more?

Group 1 = Lose 1 electron Is it easier for group two atoms to lose their 2

valence electrons or to gain 6 more?

Group 2 = Lose 2 electrons Is it easier for group seventeen atoms to lose their

7s electron or to gain 1 more?

Group 17 = Gain 1 electron

In general, metals are electron losers,

and nonmetals are electron gainers.

Metals * All atoms that lose electrons and

form positive ions are called cations.

* All metals form cations• Metals are electron losers

(so they become positive.)

Example: Lithium An atom of lithium has 3 protons and 3 electrons

An atom of lithium is electrically neutral because the (+) = (-), however, it is unstable.

+

++

e

e

eValence electron

This energy level can hold 8 electrons, it needs 7 more to be stable

Lithium will lose its 1 valence electron to become stable. Once Lithium loses its 1 valence electron, the number of (+) = (-)

Now lithium has 1 more (+) than (-). Lithium has become an ion with a charge of 1+. Written as Li1+

++

+

e

e

Nonmetals* All atoms that gain electrons and

form negative ions are called anions.

* All non-metals form anions.

* Non-metals are electron gainers

Example: Fluorine An atom of fluorine has 9 protons and 9

electrons. An atom of fluorine is electrically neutral;

however, it is unstable.

++

+

+

+

++

++

e

e

e

e

e

e

e

This energy level can hold 8 electrons, it needs one more to become stable

e

e

Fluorine will gain 1 electron to become stable. Once fluorine gains an electron, the number of (+) = (-)

Now, fluorine has 1 more (-) than positive. Fluorine has become an ion with a charge of 1- Written as F1-

++

+

+

+

++

++

e

e

e

e

e

e

e

e

e

e

Now, fluorine has a complete outer shell.

Bonding Atoms become stable by transferring or sharing

electrons with other atoms.

Ionic Bonds are formed when atoms transfer electrons (occur between metals and non-metals)

ex. NaCl (salt)

Covalent Bonds are formed when atoms share electrons (occur between non-metals)

ex. F2

Hydrogen Bonds into Helium Hydrogen is a very reactive

element. When it bonds with itself, it becomes helium, which is non-reactive.

Hydrogen reaction was very evident in the zeppelin (blimp) Hindenburg disaster.

The Hindenburg Disaster

Broadcast Recording:

http://www.youtube.com/watch?v=CgWHbpMVQ1U

Ionic Bonds

Occur between a metal and a non-metal. Ionic bonds form when electronsare gained or lost.

Covalent Bonds

                                  

Covalent Bonding and the Octet Rule:Atoms will share electrons so that each atom will fill its valence shell

Equal sharing of electrons between two atoms in a molecule

Hydrate vs Anhydrous Compounds

A hydrate compound is one that has water chemically attached to its ions and written in the formula (H20). If water is attached, the formula is followed by a dot (*):

CaSO4 * 2H20

A anhydrous (without water) compound means that water has been removed, so it reverts back to the original compound/

Molecules

When a neutral particle forms as a result of covalent bonding, it is called a molecule.

Sometimes one atom is much larger than the other in bonding (ex: Oxygen is larger than Hydrogen). So electrons being shared in a molecule are held more closely to the atoms with the larger nucleus.

Polar Molecules

When the larger element in the molecule has a stronger negative hold, then the smaller elements on the ends have a slightly positive charge. A polar molecule has a slightly positive end and a slightly negative end, causing them to attract similar molecules. This is what gives water its unique properties.

Polar Molecule

Nonpolar Molecules

When electrons are equally shared in bonds so that the molecule does not have oppositely charged ends. Molecules may form identical atoms or molecules that are symmetric (CCl4).

A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.

Why should two atoms share electrons?

F F+

7e- 7e-

F F

8e- 8e-

F F

F F

Lewis structure of F2

lone pairslone pairs

lone pairslone pairs

single covalent bond

single covalent bond

9.4

8e-

H HO+ + OH H O HHor

2e- 2e-

Lewis structure of water

Double bond – two atoms share two pairs of electrons

single covalent bonds

O C O or O C O

8e- 8e-8e- double bonds

Triple bond – two atoms share three pairs of electrons

N N8e-8e-

N N

triple bond

or

9.4

Ionic and Covalent Song

http://www.youtube.com/watch?v=BCYrNU-7SfA

Polyatomic Ion

Poly means many, so polyatomic is many atoms. A polyatomic ion is a positively or negatively charged, covalently bonded group of atoms. For example, baking soda is hydrogen carbonate HCO3.

Oxidation numbers

Oxidation numbers indicate the number of electrons an atom will lose or gain to become stable.

Elements in the same family have the same oxidation number.

The transition metals have variable oxidation numbers.

All metals have a (+) oxidation numbers (meaning they lose electrons to stable)

All nonmetals have a (-) oxidation numbers (meaning they gain electrons to be stable)

What’s Next?

Now you are ready to use this information to write chemical formulas.

Chemical Formulas A chemical formula is a shorthand

method that uses symbols to tell which elements are in a compound, and their ratios. Example:

3CaF2 means that there are

3 Molecules of Calcium Fluoride There are a total of 9 atoms in

the 3 molecules

Rules for Writing Formulas:

The sum of the oxidation numbers for each element must equal zero. All compounds must have a neutral charge.

The ion with the positive oxidation number is written first. The ion with the negative oxidation number is written second.

Do not include the charges of the oxidation numbers in the formula.

Writing formulas for Binary Compounds

Definition: compounds made of two different elements.

Steps:

1) Write the symbol of the element with the positive oxidation number as a superscript.

2) Write the symbol of the element with the negative oxidation number as a superscript.

3) Criss-cross the oxidation numbers so they become subscripts. Leave the (+) and (-) signs behind.

Examples:

1) Sodium and Oxygen

Determine Ox. #s: Na 1+ O 2-

Criss-Cross: Na 1+1+ O 2-2-

Na22O11

The one is not needed.

The correct formula is: Na22O

2) Magnesium and Chlorine

Determine the Ox #s: Mg 2+ Cl 1-

Criss Cross: Mg 2+2+ Cl 1-1-

Mg11Cl22

The one is not needed.

The correct formula is: MgCl22

Criss-Cross with Subscripts Remember, you never change the

subscript. If you use the criss-cross method and there is an existing subscript, DO NOT CHANGE IT. It just means that you will need to put that element or compound in parentheses and put the new subscript outside of the parentheses.

Examples

Al+3 + CO32- →Al2 (CO3)3

NH4 + CO32- → (NH4)2CO3

Notice that, if there is already a subscript, you put the new subscript outside of a parentheses.

Breaking Down Formulas

When you have a compound that you need to break down, then just “undo” the criss-cross method. Subscripts stay in place if they are next to an element. If the subscript is outside of parentheses, then they represent the oxidation number of the opposite element or compound.

Criss- Cross:POLYATOMIC IONS!

Before you do the criss-cross method, always check the oxidation number of polyatomic ions (page 619.) These have to be written first, before doing the criss-cross method.

Cation Anion

Ex: K3PO4 K PO4

then K PO43-

FIND THE OXIDATION NUMBER FOR

POTASSIUM (K) ON THE PERIODIC TABLE K 1+ PO4

3-

Breaking Down Formulas Examples:

Cation (+) Anion (-) NaOH Na OH

KNO3 K NO3

Ca(OH)2 Ca+2 OH

Because the 2 is outside the parentheses, it has to be the oxidation number of the first element or compound.

Writing Names of Binary Ionic Compounds

You can write the name of a binary formula by

1) Write the name of the positive ion

2) If the positive ion can form more than one oxidation number, then you write the charge using roman numerals in parentheses.

Example copper can be +1, +2 or +3, so the formula would say Cu(I) Cu (II) or Cu (III). This is only the case if the positive ion can have more than 1 oxidation number. Otherwise, you just write the positive ion first.

Writing Names of Binary Ionic Compounds

3) Write the root name of the negative ion, The root is the first part of the element’s name. For example: Chlorine = chlor-

Oxygen = ox-

4) Add the ending ide to the root. For example: BaF2 is Barium Fluoride.

The subscript does not influence the name of ionic compounds.

Writing Names of Polyatomic Ionic Compounds1) You would always write the

name of the positive ion first.

2) Look on page 619 for the endings of polyatomic ions. You would write this ending after the name of the positive ion.

http://www.youtube.com/watch?v=wWUYHHo-zB0

Conclusion

Valence electrons determine the reactivity of an atom.

All atoms seek full outer shells (typically, an outer shell of eight electrons).

Atoms form ionic or covalent bonds to achieve stability.

Oxidation numbers assist in the development in the chemical formulas.