Unit 4 Notes

ChemistryMr. Nelson

2009

Chemical Bonds• Three basic types of bonds

– Ionic• Electrostatic attraction between

ions

– Covalent• Sharing of electrons

– Metallic• Metal atoms bonded to several

other atoms

Why do atoms bond?

• Why DON’T some atoms bond?– The noble gases – why?

• Why do other atoms bond, then?– They are more chemically stable when bonded

How do atoms bond?

• The octet rule– The octet rule, or rule of eight, says that an atom will

strive for a full s and p subshell

– Atoms will either lose or gain electrons to get 8 in the outer shell

– NOTE: when an atom loses or gains electrons, it’s nucleus remains the same – only the outer electron shell has changed!!!

Bonding and energy changes

• Energy is the ability to do work

• Stability is a measure of inability to do work– So, the lower the energy, the more stable something is!

• When atoms bond, the process favors stability (lower energy). Things will never go from a stable to an unstable state on their own!

Electrons, bonding, and IONS

• When they do this, they get a CHARGE, because protons (+) and electrons (-) are no longer equal. They are now IONS

• Positive and negative IONS come together and balance each other out in IONIC BONDS.

Cations and Anions

• Remember: + + +A “plussy cat”

An “antion”

Ionic Bonding

• Sodium wants to GIVE an electron, Chlorine wants to GET an electron.

Ionic Bonding

• The low ionization energy of sodium and the high electronegativity of chlorine is one reason this works so well.

•

Energetics of Ionic Bonding

As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.

Energetics of Ionic Bonding

But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

Energetics of Ionic Bonding• There must be a third

piece to the puzzle.

• What is as yet unaccounted for is the electrostatic attraction between the newly-formed sodium cation and chloride anion.

Lattice Energy

• This third piece of the puzzle is the lattice energy:– The energy required to completely separate a mole of a

solid ionic compound into its gaseous ions.

Lattice Energy

• Lattice energy, then, increases with the charge on the ions.

• It also increases with decreasing size of ions.

Energetics of Ionic Bonding

By accounting for all three energies (ionization energy, eletronegativity, and lattice energy), we can get a good idea of the energetics involved in such a process.

Energetics of Ionic Bonding

• These phenomena also helps explain the “octet rule.”

• Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

Naming ions

• Monatomic ions = – One atom ions

• Polyatomic ions =– Many atom ions

• Naming monatomic ions– To name positive ions, just add the word “ion”– To name negative ions, drop the last part of the word,

and add “-ide ion”

Naming monatomic ions

Rubidium loses an electron to become Rb+

Rubidium ion

Calcium loses two electrons to become Ca2+

Calcium ion

Chlorine gains an electron to become Cl-

Chloride ion

Oxygen gains two electrons to become O2-

Oxide ion

Nitrogen loses three electrons to become N3-

Nitride ion

Compounds made of two monatomic ions

• These are called BINARY COMPOUNDS

• You always put the positive part first and the negative part last: Na+ + Cl- NaCl

• Names = name of the positive ion + name of the negative ion: Sodium Chloride

Examples: Name the following

Write the formulas of the following:

Back to ions: Writing Ionic Formulas

• The nomenclature (naming system):1. Write the symbols for the ions side by side. Write

the cation first.Al3+ O2-

2. Find the lowest common multiple that will make the charges on each ion cancel out

Al3+ O2-

3. Check the subscripts for the lowest whole number ratio of ions. Then write the formula.

Al2O3

d-block naming

• Write the electron configuration for Iron.

• Predict the oxidation number

d-block• The d-block (yo) has its own rules

– Metals in the d-block have variable charges– All d-block metals must get a ROMAN NUMERAL to

indicate the charge

– EXAMPLE: copper (I) chloride is made of Cu1+ and Cl-

– EXAMPLE: copper (II) chloride is made of Cu2+ and Cl-

– Don’t use roman numerals if you don’t have to

Examples

• Write the formulas for– Tin(II) iodide

– Cobalt(III) chloride

Working backward• If you are given the formula you need to calculate

the charge of the d-block metal.

• Assume the anion did not change its charge (they are very consistent)

• Example: FeO, to write the name we need the charge of iron.

A few more examples

• PbS2

• MnBr3

• Cu3P2

More d-block (old school)

• Roman numerals are the ‘new’ way.• The ‘old’ way is based on Latin names

– Two endings• ic• ous

• ic is for the highest charge• ous is for the lower charge

– Example• Ferric = Iron(III)• Ferrous = Iron(II)

Old School

• The only three I expect you to know are:

• Tin (Sn)– Stannic = Tin(IV)– Stannous = Tin(II)

• Copper (Cu)– Cupric = Copper(II)– Cuprous = Copper(I)

Polyatomic ions

• When two or more ions are clumped together it is a polyatomic ions.

• They usually end with –ates or -ites

Nomenclature of Oxyanions

• They are not standard!– Example Sulfate vs Phosphate

• Nomenclature examples– Perchlorate– Chlorate Nitrate– Chlorite Nitrite– hypochlorite

Writing formulas for compounds with polyatomic ions

• Polyatomic ions should ALWAYS be treated like BOY BAND. Don’t ever break it up!

• If you need more than one polyatomic ion to balance a charge, use PARENTHESES ( )

Polyatomic ions

• Naming compounds that contain polyatomic ions:– The steps are the same:– the name of the first ion + the name of the second:

• NH4+ = ammonium ion (polyatomic)

• Cl- = chloride ion (monatomic)

• NH4Cl = ammonium chloride

Example

• Write the formulas for:

• potassium perchlorate

• tin(IV) sulfate

• Iron(II) chromate

• ammonium sulfate

Ionic vs. Metallic bonds

• In an IONIC BOND, the electrons of one atom (that wants to lose electrons) are donated to the electrons of another atom (that wants to gain electrons). The charges on each ion balance each other out and equal ZERO.

• In a METALLIC BOND, all the atoms are the same (all copper, for example) and the electrons don’t belong to any one atom. They move around a lot – that’s why electricity is conducted.

Metallic Bonds

• A “sea” of mobile outer electrons.

• Low ionization energies means the atoms don’t hold electrons well.