Ionic Bonding

Valence Electrons, Lewis Dot Structures,

and Electronegativity

Valence Electrons

• valence electrons – the outermost electrons.

6 protons = C = carbon

Inner electrons = 2 e– not available for bonding

Outer electrons = 4 e–available for bonding

4 electrons in valence shell

Lewis Dot Structures

• Lewis dot structures are a convenient way to show how many valence electrons an atom has.

• Example: Draw the Lewis dot structure for hydrogen.

H

Lewis Dot Structures

dots = number of valence electrons.

The maximum number of dots is 8.

Look for the number at the top of the column (e.g. 5A).

Exception: Helium only has 2 dots. He

Ne

More Lewis Structure PracticeDraw the Lewis structure for oxygen.

Draw the Lewis structure for magnesium.

Draw the Lewis structure for chlorine.

O

Mg

Cl

Even More Lewis Structure PracticeDraw the Lewis structure for carbon.

Draw the Lewis structure for potassium.

Draw the Lewis structure for phosphorus.

C

K

P

Electronegativity

• electronegativity – how much an atom wants to keep hold of its electrons.

• ionization energy – the energy required to remove an electron from an atom.

Lower electronegativity

Greater electronegativity

Metals

Nonmetals

Role Models: The Noble Gases

• An atom’s electrons are at their most stable when they reorganize their electrons to more closely resemble the electron configuration of a noble gas.

• All atoms want to have stable electron configurations.

Role Models: The Noble Gases

• All atoms wish their electrons were like the noble gases’ electrons.

Example: Beryllium and Oxygen

4 protons = Be = beryllium

The 2 valence e–

Now beryllium’s matches that of helium

8 protons = O = oxygen

The 6 valence e– Now oxygen’s matches that of neon

Same Thing, but in Lewis Dot Structure

Be O

Nonmetals

Metals

Semi-metals or metalloids

Three General Bonding Types

• Metal with Nonmetal- form ionic compounds

• Metal with Metal- form metallic compounds

• Nonmetal with Nonmetal- form covalent compounds

Metal with Nonmetal Bonding

• Ionic compounds – the metal gives all of its valence e– to the nonmetal.

• Known as – Salts, ions

Ions• ion – an atom that gained or lost electrons to

become more like a noble gas.

+ Metals lose electrons to become positively charged ions. We call them cations (cat-ions) the “t” looks like a “+”. [e.g. 2A +2]

– Nonmetals gain electrons to become negatively charged ions. We call them anions (an-ions) “n” for negative “–”. [e.g. (8 – 6A) × -1 -2]

Writing the Charges• Write out the ion that sodium forms.

Na+

• Write out the ion that chlorine forms.Cl–

• Write out the ion that magnesium forms.Mg2+

• Write out the ion that oxygen forms.O2–

So where do the electrons go?

• Usually atoms that become cations give their electrons to anions.

• Now both the cations and anions resemble noble gases, however now both have net charges.

Be O2+ 2–

Basic Electrical Charge Laws

+ and – : Attract(pull

together)

Naming (aka nomenclature)• Metals keep their names unchanged.

(e.g. sodium, aluminum, calcium) • Transition metals have their charge shown as roman

numerals in parenthesis after the name.Fe2+ iron (II) Cu1+ copper (I)Fe3+ iron (III) Cu2+ copper (II)

• Nonmetals have the last one or two syllables of their names altered with an –ide ending.

fluorine fluoride nitrogen nitride chlorine chloride oxygen oxide

Naming (aka nomenclature)

• Metals keep their names unchanged. (e.g. sodium, aluminum, calcium)

• If there are more than one possible charge for a metal (the transition metals), the charge will be specified in roman numerals after the name.

Fe2+ iron (II) Cu1+ copper (I)Fe3+ iron (III) Cu2+ copper (II)

Naming Continued

• Nonmetals have the last one or two syllables of their names altered with an –ide ending.

• Examples: carbon carbide fluorine fluoridenitrogen nitride chlorine chlorideoxygen oxide bromine bromidesulfur sulfide iodine iodidephosphorus phosphide

Naming Continued• Now put the metal and nonmetal ion names

together and you get the name for the ionic compound.

• Examples: LiF lithium fluoride NaCl sodium chloride KBr potassium bromideMgS magnesium sulfide CuI copper (I) iodide CuO copper (II) oxide FeN iron (III) nitride

Sodium Chloride – NaCl

Na+ Cl– Cl–

Cl–Cl–

Cl–Cl– Na+

Na+

Na+

Na+

Na+

+ 6– 6

0

Magnesium Chloride

Mg2+ Cl–

Mg2+

Mg2+

Mg2+

Mg2+

Mg2+ Cl–

Cl–

Cl–

Cl–

Cl–

+ 12– 6

0 + 6– 12

Cl–

Cl–

Cl–

Cl–

Cl–

Cl–

Magnesium Chloride

Mg2+

Cl–

6 x

12 x

MgCl2

Mg6Cl12

6 112 2

=

Magnesium Chloride

Mg2+

Cl–

Cl–MgCl2

Formula Unit (f.u.) – the smallest amount of an ionic compound that still has the same ratio of ions as in the formula.

Empirical Formula

Iron (III) Oxide

Fe2O3

Fe3+ O2–

criss-cross

2 x (+3) =3 x (–2) = –6

+6

0

Sodium Cloride

Na1Cl1

Na+ Cl–

criss-cross

NaCl

Magnesium Sulfide

Mg2S2

Mg2+ S2–

criss-cross

MgS

2 12 1

=

Sodium Oxide

Na2O

Na+ O2–

criss-cross

REMEMBER

Fe2O3

Fe3+ O2– Top right corner: Charge

Bottom right corner: How many atoms/ions

Polyatomic Ions

• Poly – many• Atomic – having to do with atoms

• Polyatomic ions – ions made from multiple atoms

• List on p.257

Polyatomic Ions (List on p.257)Ion name FormulaAcetate CH3COO– Ammonium NH4

+

Carbonate CO32–

Chromate CrO42–

Cyanide CN–

Dichromate Cr2O72–

Hydroxide OH–

Ion name Formula

Nitrate NO3–

Nitrite NO2–

PermanganateMnO4

–

Peroxide O22–

Phosphate PO43–

Sulfate SO42–

Sulfite SO32–

Thiosulfate S2O3

2–

Magnesium Nitrate

Mg (NO3)2

Mg2+ NO3–

criss-cross

1 x Mg 2 x N 6 x O

REMEMBER

Mg (NO3)2

Mg2+ NO3– Top right corner:

Charge

Bottom right corner: How many atoms/ions

Sodium Chloride – NaCl

Na+ Cl– Cl–

Cl–Cl–

Cl–Cl– Na+

Na+

Na+

Na+

Na+

Sodium Chloride – NaCl

Sodium Chloride – NaCl

Sodium Chloride – NaCl

Sodium Chloride – NaCl

Sodium Chloride – NaCl

Magnesium Chloride (MgCl2)

Mg2+ Cl–

Mg2+

Mg2+

Mg2+

Mg2+

Mg2+ Cl–

Cl–

Cl–

Cl–

Cl–

Cl–

Cl–

Cl–

Cl–

Cl–

Cl–

Magnesium Chloride (MgCl2)

Magnesium Chloride (MgCl2)

Calcium Fluoride (CaF2)

Calcium Fluoride (CaF2)

Calcium Fluoride (CaF2)

Calcium Fluoride (CaF2)

Calcium Fluoride (CaF2)

Calcium Fluoride (CaF2)

NeFONCBBe

He

Li

H

Kr

ArCl

Br

XeI

SPSiMg Al

Ca

Na

K

4 e– in valence shell

Ions• ion – an atom that gained or lost electrons.

• metal ions lose electrons to become more positively charged. (e.g. 2A +2)

• Nonmetal ions gain electrons to become more negatively charged. (e.g. 8 – 6A –2)

Ions• ion – an atom that gained or lost electrons.

• metal ions lose electrons to become more positively charged. (e.g. 2A +2)

• Nonmetal ions gain electrons to become more negatively charged. (e.g. 8 – 6A –2)

Ions - again• ion – an atom that gained or lost electrons to

become more like a noble gas.

+ Metal ions = cations (cat-ions) the “t” looks like a “+”.

– Nonmetal ions = anions (an-ions). “–”