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 Cations are pawsitive 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... Draw electrons (dots) Move e- to most stable configuration Draw arrows to show new location 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!!!   The first part is the name of the cation (+) metal  with NO change to the name Sodium  stays sodium 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 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: Place the symbol of the cation first! Follow this symbol with the symbol for the anion Use subscripts to show the ratio of the ions in the compound 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?  Identify the cation and anion. Na = Sodium, S = Sulfur How many electrons does sodium lose to become stable? 1 What is the overall charge on sodium after it loses electrons? +1 so, Na+ How many electrons does sulfur accept to become stable? 2 What is the overall charge on sulfur after it gains these electrons? –2 so, S2- 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 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: 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 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: Write the symbols for the elements in the order they appear in the name Prefixes indicate the number of atoms in each element in the molecule 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