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Chapter 6 Chemical Bonds
6.1 Ionic Bonding & Naming
6.2 Covalent Bonding & Naming
Metallic Bonds
6.1 Ionic Bonding
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-
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:
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
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.
Common Ions
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!!!
Formation of Ions• Cation= ion with a positive charge
(has lost e-)
–Cations are pawsitive
• Anion= ion with a negative charge (has gained e-)
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.”
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
Ionic Bond Formation
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.
Ionization Energy
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.
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
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?
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
Example - NaCl
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.
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!
• 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
6.2 Covalent 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!
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
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.
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.
IONIC/COVALENT BONDS
• Ionic bonds– Metal to
nonmetal– Electrons are transferred
• Covalent bonds– Nonmetal to
nonmetal– Electrons are shared.
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)
Greek Prefixes
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
• Now, for one last type of bond…and I don't mean James Bond....
• Waaayyy more daring and exciting are the infamous...
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
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!
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.
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.
So what happens if we add more valence electrons?
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)!
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
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!
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)!
If a metal is struck, the ions simply shift to new positions, but they are still connected by metallic bonds