Unit 5 – Bonding & Nomenclature

Ionic Bonds, Covalent Bonds, and Writing Formulas

In our last unit…•We talked all about the ATOMATOM– Development of the atomic model, subatomic

particles•HOWEVER, we were only talking about atoms as one individual unit

In this unit we will be looking at how atoms combine (or BONDBOND) with other

atoms to form compounds!

Electrons are Arranged in Shells– As we move down a family, the elements gain an

electron shell.• The first shell can hold up to 2 electrons.

(Row/Period 1)• The second and third shells can hold up to 8

electrons each. (Periods 2 & 3)• The fourth and fifth shells can hold up to 18

electrons each. (Periods 4 & 5)• The sixth and seventh shell can hold up to 32

electrons each. (Periods 6 & 7)

The Valence Electrons Give Atoms Their Properties

• The outer most shell is called the valence shell.• The electrons in the valence shell are called the

valence electrons.• Atoms tend to gain or lose electrons to end up

with a full valence shell. (8 e-)– Except H2 and He which want a duet (2 e-)

• The number of valence electrons correspond to the digit in the ones place of the group number for elements in groups 1, 2, 13-18.

Electron Dot Diagrams Show Valence Electrons

• An electron dot diagram depicts an atom as its symbol and its valence electrons.– Ex: Carbon

Carbon has four electrons in its valence shell (carbon is in group 14), so we place four dots representing those four valence electrons around the symbol for carbon.

Electron Dot Diagrams Show Valence Electrons

• Electrons are placed one at a time in a clockwise manner around the symbol in the north, east, south and west positions, only doubling up if there are five or more valence electrons.– Example: Chlorine (7 valence electrons b/c it is in

group 17)

Paired and Unpaired Electrons• As we can see from the chlorine example, there are

six electrons that are paired up and one that is unpaired.

• When it comes to bonding, atoms tend to pair up unpaired electrons.– They do this in two ways that we will discuss in the

chapter.

1. A bond that forms when one atom gives an unpaired electron to another atom is called an ionic bond.

2. A bond that forms when atoms share unpaired electrons between each other is called a covalent bond.

Ionic Bonding

• Ionic Bond–Bond between metal and nonmetal due

to “electrostatic interactions” –Electrons are transferred from metal to

nonmetal forming a positively charged metal ion (cation) and negatively charged nonmetal ion (anion)

6.1 Ions and Ion Formation• If one or more valence electrons are lost or gained,

the balance between positive and negative particles in the atom is upset.– When an atom gains electrons, there are more negative

charges in the atom than positive charges, so it becomes negative• Example – Oxygen gains two electrons so it gains a charge of -2.

– When an atom loses electrons, there are more positive charges in the atom than negative charges, so it becomes positive.• Example – Magnesium loses two electrons so it gains a charge

of +2.

Atoms Do What is Easiest For Them!

Atoms tend to do the least to gain a filled valence shell.•Elements in groups 1, 2 and 13 tend to lose electrons to wind up with a filled valence shell because it only involves the removal of 1, 2, or 3 electrons, not the addition of 7,6, or 5 electrons respectively.•Likewise for elements in groups 15,16, and 17, only in their case, they tend to gain 3,2, or 1 electrons, not lose 5,6, or 7 electrons respectively.

Examples

• Sulfur

• Sodium

Get Out Your Periodic Tables!• What are the charges for groups 1, 2, 13, 15,

16, 17, and 18?– Group 1:

– Group 2:

– Group 13:

– Group 15:

– Group 16:

– Group 17:

– Group 18:

Atoms That Gain Or Lose Electrons Are Called Ions

(we saw this last chapter!)

• An ion is formed when the number of electrons differs from the number of protons in an atom. – This is accomplished by gaining or losing

electrons, not by adjusting the number of protons!

– The charge of an ion is denoted as a superscript in the atom’s isotope notation.• Ex: A Nitrogen ion:

6.2 Ionic Compounds

• When an atom that tends to lose electrons comes in contact with an atom that tends to gain them, an electron transfer occurs.– As a result, two oppositely charged ions are

formed.– The two oppositely charged ions are attracted to

each other via the electrostatic attractions.– These attractions create an IONIC BONDIONIC BOND

Ionic bonds Result from a Transfer of Valence Electrons

+ -

• Ionic Compounds - All chemical compounds containing ions (one cation, one anion)

• Ionic compounds typically are formed from elements on opposite sides of the periodic table.

• For all ionic compounds, positive and negative charges must balance out.

– Formula must be neutral overall!

Ionic Bonds Result from a Transfer of Valence Electrons

Properties of Ionic Compounds

• Positive and negative ions come together to create an organized pattern called a lattice structure.

• An example of a lattice structure for sodium chloride

Properties of Ionic Compounds

• Ionic compounds tend to be brittle• They have high melting and boiling points• They tend to be soluble in water.• They can conduct electricity when dissolved in

H2O

Many of the physical and chemical properties of the components of the compound are changed

when they bond.

6.3 Writing Ionic Formulas

+3 +3

-2 -2 -2

+2 +2+1 +1

-1 -1 -1

+3 +3

-2 -2 -2

6.3 Writing Ionic Formulas

Subscripts Show the Number of Atoms of Each Element in a Compound

• A subscript notes how many atoms of each element are in a subunit of a compound.

• 1 is NEVER used as a subscript.• Ex: NaCl 1 Sodium atom : 1 Chlorine atom Al2O3 2 Aluminum atoms : 3 Oxygen atoms.

• The subscripts determine how many of each atom is required to balance out the positive and negative charges from the ions.

Writing Formulas for Ionic Compounds• We can use the charges of the ions in a compound

to easily determine the formula of the compound.• All you do is crisscross the charges to make them

subscripts (Reduce the subscripts if you can!!)– Ex: magnesium nitride

Mg N+2 -3

23

Don’t forget to drop the positive/negative signs once they’re subscripts!!

PRACTICE, PRACTICE, PRACTICE!!Write formulas for the following ionic compounds:

Magnesium chloride

Calcium bromide

Sodium oxide

Beryllium oxide

Potassium iodide

Polyatomic Ions• There are some ions that are made of more

than one element– These ions are called POLYATOMIC IONSPOLYATOMIC IONS• “poly” meaning more than 1, “atomic” meaning atoms

more than 1 type of atom in the ion

• Examples: (see chart in your notes for more)– Sulfate SO4

2-

– Hydroxide OH-

– Nitrate NO3-

Writing Formulas w/ Polyatomic Ions1. Write both the ions with their accompanying charges.2. Crisscross the charges to make them subscripts.– Use parenthesis around the polyatomic ion if there is more

than one of the ion!

3. Drop the positive and negative signs once they are subscripts.

4. Reduce subscripts if possible.

Example: Barium phosphate

PRACTICE WITH POLYATOMICS!Write formulas:Write formulas:

calcium phosphate beryllium bicarbonate

lithium carbonate

Magnesium sulfate

Barium hydroxide

Don’t forget to reduce subscripts that come from the charges if you can!

Don’t forget to reduce subscripts that come from the charges if you can!

Writing Formulas w/ Transition Metals

• Transition metals – all elements in groups 3-12 and elements in groups 13, 14, and 15 below the stair step line– Can have more than one charge in ion form!• Except Ag+ and Zn2+

• When writing formulas for compounds involving transition metals, you will get the charge on the metal cation from the Roman numeral in parenthesis after the metal’s name

Writing Formulas w/ Transition MetalsWriting Formulas of Compounds w/ Transition Metals:1. Write the metal’s atomic symbol with the value of the

Roman numeral as a positive charge2. Write the value of the anion (negatively charged ion)

with its charge3. Crisscross the charges to get your formula

Example: Lead (IV) hydroxide

Writing Formulas w/ Transition Metals

Examples:

Copper (II) chloride

Tin (IV) sulfate

Gold (I) nitrate

Don’t forget to reduce subscripts that come from the charges if you can!

Don’t forget to reduce subscripts that come from the charges if you can!

On a separate piece of paper:Please write formulas for the following compounds & turn them in to me on a separate piece of paper so I can check them.

1.Sodium fluoride

2.Magnesium phosphate

3.Copper (II) sulfate

4.Tin (IV) phosphide

5.Hydrogen cyanide

6.4 Naming Ionic Compounds• When naming ionic compounds, we take the

name of the metal followed by the name of the non-metal, only we drop the last syllable and add –ide as an ending.

– Examples: NaCl: Sodium Chloride Al2O3: Aluminum Oxide

PRACTICE, PRACTICE, PRACTICE!Write names for the given ionic compounds:

LiBr

AlCl3

BeO

Ra3N2

Naming Ionic Compounds with Polyatomic Ions

• The rules for naming polyatomic ions are the same rules for naming normal ionic compounds except you use the polyatomic ion name for the second word

Example: MgSO4 – magnesium sulfate

LiNO3 – lithium nitrate

Ba(OH)2 – barium hydroxide

PRACTICE WITH POLYATOMICS!

Name the compounds:Name the compounds:

RbNO3

NaHCO3

Mg(OH)2

Naming Formulas w/ Transition Metals

• Transition metals – all elements in groups 3-12 and elements in group 14 beneath the stair step line– Can have more than one charge in ion form!• Except Ag+ and Zn2+

• When naming compounds involving transition metals, you need to include a roman numeral in parenthesis to indicate the charge of the ion

Naming Compounds w/ Transition MetalsNaming Compounds Containing Transition

Metals:1. Name the metal from its symbol2. Determine the charge on the metal by multiplying the

subscript on the anion by the charge the anion normally has (see box below)

3. Write the charge as a roman numeral in parenthesis4. Write the name of the anion either as an element with

an “–ide” ending or as the unchanged name of the polyatomic ion

Charge of metal = |charge of anion × anion subscript|

Naming Compounds w/ Transition Metals

Examples:

CrCl6

Fe(OH)2

NiO

Covalent BondingCovalent Bonds–Bonds in which e- are shared

between two atoms –Most common type

6.5 Covalent Bonds Result from a Sharing of Valence Electrons

• Elements that have a partially filled valence shell can complete them by sharing electrons with another element

• The mutual attraction for shared electrons is called a covalent bond (co- signifies sharing, -valent signifies valence electrons)

Shared Electrons Complete Shells

F F

• A substance made up of atoms which are held together by covalent bonds is a covalent compound.–They are also called molecules.

6.5 Covalent Bonds Result from a Sharing of Valence Electrons

Drawing Electron Dot Diagrams for Molecules

• Chemists usually denote a shared pair of electrons as a straight line.

F F• Sometimes the nonbonding pair of electrons

are left off of the electron dot diagram for a molecule

Nonmetals Tend to Form Covalent Bonds

• A covalent bond is formed when two atoms that tend to gain electrons come into contact with each other.

• Hydrogen tends to form covalent bonds because it has a fairly strong attraction for an additional electron.

• The number of covalent bonds an atom can form is equal to the number of additional electrons it can attract.

H H

Examples

OH

H

H2

H2O

More Examples

CH HH

H

CH4

NH3

H HHN

Multiple Bonds• Sometimes two atoms share more than just one pair of

electrons which results in a double bond or triple bond.

O2 O O

CO2 C OO

Writing Formulas for Covalent Compounds(2 non-metals bonded together)(2 non-metals bonded together)

1. Write the symbol for the first element listed• If there is a prefix on the first element, write the

prefix value as a subscript attached to the first element’s symbol

2. Write the symbol for the second element listed• If the prefix on the second element is “di-” or

greater, write the prefix value as a subscript attached to the second element’s symbol

Ex: Diphosphorous monoxide

Writing Formulas for Covalent Compounds(2 non-metals bonded together)(2 non-metals bonded together)

Examples:

Diphosphorous hexoxide

Carbon tetrachloride

Nitrogen trihydride

Naming Covalent Compounds (2 non-metals bonded together2 non-metals bonded together)

Given the formula:1. Write the names of the 2 elements2. Attach a Greek prefix (see chart) to the beginning of

the first element only if it has a subscript greater than or equal to 2

3. Attach a Greek prefix to the second element no matter what!

4. Change the ending of the second element’s name to an –ide ending

Ex: Se2O3

Naming Covalent Compounds (2 non-metals bonded together2 non-metals bonded together)

Examples:

P2O6

CCl3

NO

In review…Chemical Bonds– Force that holds atoms together– It’s all about the electrons (e-)

Types of Chemical Bonds– Ionic• Bond between metal and nonmetal due to “electrostatic

interactions” • Electrons are transferred from metal to nonmetal

In review…Covalent Bonds– Bonds in which e- are shared – Most common type

Metallic Bonds– Atoms are bonded

to one another (not to other elements)

– Positive ions in a “sea” of negative charge (e-)– Example: Cu, Ag