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Chapter 6 Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Chapter 6 Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

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Page 1: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Chapter 6 Chemical Bonds

6.1 Ionic Bonding & Naming

6.2 Covalent Bonding & Naming

Metallic Bonds

Page 2: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

6.1 Ionic Bonding

Page 3: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Stable Electron Configurations

•When the highest occupied energy level of an atom is filled with electrons, the atom is STABLE and NOT LIKELY TO REACT.

–ex: the Noble gases (8 valence e-)

•Chemical properties of an element depend on the number of valence e-

Page 4: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 5: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Representing Electron Configurations

•Use electron dot diagram or a Lewis Structure to focus on the valence e-

•Model of an atom in which each dot represents a valence e- (symbol in the center represents the nucleus and all the other e- in the atom)

–ex:

Page 6: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 7: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Ionic Bonds • Elements that do not have complete

sets of valence e- tend to react

• By reacting, they achieve e- configurations similar to those of noble gases (stable configurations)

• Achieve stable e- configurations through the transfer of e- between atoms– ex: NaCl

Page 8: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Formation of Ions • Ion: an atom that has a net positive or

negative electric charge– Atoms that have the same number of protons as

electrons have NO charge. They are considered neutral!

• Ions form when an atom gains or loses an e- (number of protons is no longer equal to the number of e-)

• Charge is represented by a plus or a minus sign– Atoms will transfer electrons to achieve a stable

electron configuration. This is what creates ions.

Page 9: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 10: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Common Ions

Page 11: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Formation of Ions

• Question: When is an atom considered to have a stable electron configuration?

 • Answer: When the highest occupied

energy level is filled with 8 valance electrons!!!

Page 12: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Formation of Ions• Cation= ion with a positive charge

(has lost e-)

–Cations are pawsitive

• Anion= ion with a negative charge (has gained e-)

Page 13: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Some Humor… A sodium ion walks into a class. It says to the

teacher, “I’ve lost an electron.”

The teachers says, “Are you sure?”

The sodium ion says, “I’m positive.”

Page 14: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Formation of Ionic Bonds• Particle with negative charge attracts

particle with positive charge• Chemical Bond: force that holds atoms or

ions together as a unit• Ionic Bond: force that holds cations and

anions together• Forms when e- are transferred from one

atom to another

Page 15: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Ionic Bond Formation

Page 16: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Ionization Energy• Ionization Energy- amount of energy

used to remove an electron– It allows electrons to overcome the attraction

of the protons in the nucleus – Trend- ionization energy increases from left

to right and decreases from top to bottom

• It takes more energy to remove an electron from a nonmetal than a metal– Example: It takes more energy to remove an

electron from Na than K.

Page 17: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Ionization Energy

Page 18: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 19: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Ionization Energy

• Question: how does ionization energy relate to the reactivity of the alkali metals?

• Answer: low ionization energy = easier to remove electrons. – The easier it is to remove an electron the

more reactive an atom is.

Page 20: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

 DRAWING IONIC COMPOUNDS

• FIRST, YOU HAVE TO THINK...! • HOW MANY ATOMS OF CHLORINE DOES

SODIUM NEED TO REACT WITH TO BE STABLE??? 

• THEN FOLLOW THE DIRECTIONS BELOW...1. Draw electrons (dots)2. Move e- to most stable configuration3. Draw arrows to show new location4. Write the new charge of the ions formed.

• Draw ionic bond for sodium & chlorine• Draw ionic bond for magnesium and chlorine

Page 21: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Chemical Formulas• Chemical formula: a notation that shows

what elements a compound contains and the ratio of the atoms or ions of those elements in the compound– Examples: H20, NaCl, C6H1206

– Identify the number of each type of atom in the space below:

• H20 How many hydrogen? Oxygen?

• NaCl How many sodium? Chlorine?

• C6H1206 How many carbon? Hydrogen? Oxygen?

Page 22: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Properties of Ionic Compounds• Properties of ionic compounds can be

explained by the strong attractions among ions within a crystal lattice:– High melting points– Poor conductor of electric current as a solid– Good conductor of electric current when melted– Will shatter

Page 23: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Example - NaCl

Page 24: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 25: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Naming Binary Ionic Compounds

• Binary compound- a compound made of only two elements

• Naming binary compounds is easy!!!  1. The first part is the name of the cation (+) metal

with NO change to the name • Sodium stays sodium

2. The second part is the name of the anion (-) nonmetal use part of the nonmetal name with the suffix “IDE” at the end• Chlorine becomes chloride

3. EXAMPLE Sodium Chloride

• Refer to Figure 16 on page 171 for a list of the names and charges of the 8 common anions. 

Page 26: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Writing Formulas Ionic Compounds

• STEPS: 1. Place the symbol of the cation first!

2. Follow this symbol with the symbol for the anion

3. Use subscripts to show the ratio of the ions in the compound

4. Because all compounds are neutral, the total charges on the cations and anions must add up to zero!

Page 27: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

• EXAMPLE: What is the name of the compound that is formed when sulfur and sodium react? 

1. Identify the cation and anion. Na = Sodium, S = Sulfur2. How many electrons does sodium lose to become stable? 1 3. What is the overall charge on sodium after it loses

electrons? +1 so, Na+ 4. How many electrons does sulfur accept to become stable? 2 5. What is the overall charge on sulfur after it gains these

electrons? –2 so, S2- 6. To have a neutral compound of sodium and sulfur ions,

what ratio of ions will create an overall charge of zero? TWO sodiums for every ONE sulfur

7. Use subscripts to show the appropriate ratio of atoms: Na2S

Page 28: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

6.2 Covalent Bonds

Page 29: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Review

• Question: do nonmetals have high or low ionization energies?

•  Answer: high•  Question: so, if it is difficult to

remove, or transfer, electrons between nonmetals, then how do they form bonds?

•  Answer: they have to share their electrons!

Page 30: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Covalent Bond

• Covalent Bonds- a chemical bond in which 2 atoms share a pair of valance electrons.– When two atoms share a pair of electrons,

the bond is called a single bond.– Refer to Figure 9 on page 166 to see how to

model covalent bonds.

• Molecule- atoms joined by covalent bonds

Page 31: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 32: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 33: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Covalent Bonds• Question: What keeps the atoms in

molecules together?• Answer: The attraction between the

shared electrons and the protons in the nucleus hold them together.

Page 34: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 35: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

The oxygen atom in water and the nitrogen atom in ammonia are each surrounded by

eight electrons as a result of sharing electrons with hydrogen atoms.

Page 36: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

IONIC/COVALENT BONDS

• Ionic bonds– Metal to

nonmetal– Electrons are transferred

• Covalent bonds– Nonmetal to

nonmetal– Electrons are shared.

Page 37: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Naming Molecular (Covalent) Compounds

• Rules to naming:1. Most metallic element appears first in the name

• These are further to the left on the table• If in same group – more metallic is found closer to the

bottom

2. Name of second element is changed to end in the suffix – “ide” (as in carbon dioxide)

• Example: two compounds contain nitrogen and oxygen N2O4 & NO2

– The names reflect the actual number of atoms – using Greek prefixes

– Dinitrogen tetraoxide & mononitrogen dixoxide (however, prefix mono is often not used for first element nitrogen dioxide)

Page 38: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Greek Prefixes

Page 39: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Writing Molecular (Covalent) Formulas

• Rules to writing a molecular formula:1. Write the symbols for the elements in the order

they appear in the name2. Prefixes indicate the number of atoms in each

element in the molecule3. Prefixes appear as subscripts in the formulas

• If there is NO prefix for an element – there is only one atom

• Write formula for diphosphorous tetrafluoride:– Di- indicates two phosphorous & tetra- indicates

four fluorine

– Formula would be P2F4

Page 40: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

• Now, for one last type of bond…and I don't mean James Bond....

• Waaayyy more daring and exciting are the infamous...

METALLIC BONDS

Page 41: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Metallic Bonds

• REMEMBER: metals achieve stable configurations by losing electrons

 • Question: In most cases, which class of

elements receive these electrons?•  Answer: nonmetals  

Page 42: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Metallic Bonds

• Question: What happens when there are no nonmetal atoms available to accept these electrons? 

• A: There is a way for metal atoms to lose and gain electrons at the SAME TIME!

Page 43: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 44: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Metallic Bonding• In a metal, the valance electrons are free to

move around the atom.  – This movement basically changes the metal atom

into a cation surrounded by a pool of shared electrons.

• A metallic bond is the attraction between a metal cation and the shared electrons that surround it.  – The cations in a metal form a lattice that is held

in place by strong metallic bonds between the cations and the surrounding valance electrons.

Page 45: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 46: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Not all are created equal…

• However, not all metals are created equally!

• The metallic bonds in some metals are stronger than in other metals. 

• The more valance electrons an atom can contribute to the shared pool, the stronger the metallic bonds will be.

Page 47: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds
Page 48: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

So what happens if we add more valence electrons?

Page 49: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Let’s look at the PERIODIC TABLE a moment….

• Question: how many valence electrons do the alkali metals contribute?

•  Answer: 1– So, the bonds in alkali metals are

relatively weak. – Because the bonds are weak, alkali

metals, like sodium and potassium, are soft enough to cut with a knife! (Plus, they have low melting points)!

Page 50: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Explaining Properties of Metals

• So what does this mean for the other properties of metals???

• The mobility of electrons within a metal lattice explains some of the properties of metals.

•  Question: what are two of the most important properties of metals?

•  Answer: malleability and conductivity

Page 51: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Explaining Properties of Metals

• For conductivity, remember, it’s the ability to allow a flow of electric charge…

• Metals have a built-in supply of charged particles that flow from one location to another (the pool of shared electrons ) • An electric current can be carried through a

metal by the free flow of the shared electrons! 

Page 52: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

Explaining Properties of Metals• For malleability, remember it is the ability

to be hammered without shattering… – the lattice (structure) in a metal is quite

flexible. If a metal is struck, the ions simply shift to new positions, but they are still connected by metallic bonds.

– Since the electrons still surround the ions, the metallic bond between the ions and the electrons is not broken.  

• This also explains why metals are ductile (can be drawn into wires without breaking)!

Page 53: Chapter 6  Chemical Bonds 6.1 Ionic Bonding & Naming 6.2 Covalent Bonding & Naming Metallic Bonds

If a metal is struck, the ions simply shift to new positions, but they are still connected by metallic bonds