“Structure of Matter” Covalent Bonds Ch. 6. Matter Matter is anything that has mass and occupies space. Matter is made of atoms which are the smallest

“Structure of Matter”Covalent Bonds

Ch. 6

Matter Matter is anything

that has mass and occupies space.

Matter is made of atoms which are the smallest particles that have the properties of an element.

Matter Pure substances are

any matter that has a fixed composition and definite properties.

Cannot be broken down by physical changes.

There are about 100 million pure substances that have been identified Out of these pure substances, only 118 of them are elements, the rest are compounds

Matter

Matter Elements are

substances that cannot be broken down into simpler substances.

Compounds are substances made of atoms of more than one element bound together. Every compound is

made up of a chemical formula

Chemical Formulas

A chemical formula tells us: the type of atoms present the number of atoms present the type of compound

Chemical Formulas Example: table salt:

Sodium Chloride Chemical formula:

NaCl Count the atoms

present: 1 Na atom 1 Cl atom

Chemical Formulas Sometimes there are subscripts present.

A subscript is a small number that is in a chemical formula. Example - water: H2O

2 H atoms 1 O atom

Subscript

Chemical Formulas Sometimes there are parentheses with a subscript. The subscript only

applies to the atoms within the parentheses. Example - calcium hydroxide (kidney stones): Ca(OH)2.

1 Ca atom 2 O atoms 2 H atoms

Chemical Formulas Sometimes there are subscripts in the parentheses.

Multiply the subscript outside the parentheses by the subscript of each element within the parentheses. If no subscript is present assume that it is 1.

Example - calcium nitrate: Ca(NO3)2

1 Ca atom 2 N atoms 6 O atoms (3 oxygens x 2 = 6)

Structure of Matter Nuances in molecular structure can affect its

properties. Chemical formulas can be visually represented using chemical structures which can show bond length, bond angles and atomic sizes.

Structure of Matter The structure of a compound affects its

properties. Example: strong bonds = high melting points.

Types of Molecular Structures Network

Structures: Structure: strong,

rigid structure Bond Strength:

strong Boiling and

Melting Points: high

Types of Molecular Structures Ionic network

structures: Structure: regularly

shaped crystals Bond Strength:

strong Boiling and Melting

Points: high

Types of Molecular Structures Molecular

structures: Structure:

molecules weakly bonded together.

Bond Strength: weak

Boiling and Melting Points: low

Types of Molecular Structures Molecular structures

typically experience two types of attractive force:

The attraction between molecules is called intermolecular force.

It is rarely as strong as intramolecular force which is inside the molecule.

Atomic Bonds Atoms form atomic

bonds to become more stable.

Atoms become more stable by filling their valence shell or at least meeting the octet rule by getting 8 valence electrons.

Atomic Bonds There are three main types of chemical

bonds used by atoms to fill their valence shell:

Covalent Metallic Ionic

“Bond,Chemical Bond”

Covalent Bonds In covalent bonds, nonmetal atoms meet the

octet rule by sharing one or more pairs of electrons.

The shared electron pair is called a bonding pair and represented by a line on a Lewis structure.

Covalent Bonds

Chlorine forms a

covalent bond with

itself.

Covalent Bonds

Each chlorine atom wants to gain one electron to achieve an octet.

Covalent Bonds


Covalent Bonds


Covalent Bonds

The octet is achieved by each atom sharing the electron pair in the middle.

Covalent Bonds

This is the bonding pair.

Covalent Bonds

It is a single bonding pair so it is called a single bond.

Covalent Bonds

Single bonds are abbreviated with a dash

Covalent Bonds

This is now a chlorine molecule.

Covalent Bonds

Oxygen is also a diatomic

molecule (a molecule

with 2 of the same

element bonded

together).

Covalent Bonds

How will oxygen bond?

Covalent Bonds


Covalent Bonds


Covalent Bonds

Since each oxygen has 6 valence, they would each need to gain 2 more electrons to

be stable.

Covalent Bonds

Both pairs of electrons are shared.

Covalent Bonds

6 valence electrons + 2 shared electrons = full octet

Covalent Bonds

Two bonding pairs, making a double bond.

Covalent Bonds

The double bond can be shown as two dashes.

Covalent Bonds

This is now an oxygen molecule.

Covalent Bonds Elements can share up to three pairs (6 electrons).

Single Bond (2e)

Double Bond (4e)

Triple Bond (6e)

Covalent Bonds Equal sharing of electrons

creates nonpolar covalent bonds.

Ex. Ethane, C2H6

Unequal sharing of electrons is called polar covalent bonds and can lead to molecules having a positively and negatively charged side.

Ex. Water, H20

Covalent Bonds The slight charges on a polar molecule can cause a loose atomic

bond called polar or hydrogen bond.

Covalent Bonds Nomenclature Naming binary covalent

compounds: Two nonmetals Name each element End the last element in –ide Add prefixes to show more

than 1 atom or 1 atom on the second element.

# of Atoms Prefix

1 mono-

2 di-

3 tri-

4 tetra-

5 penta-

6 hexa-

7 hepta-

8 octa-

9 nona-

10 deca-

Covalent Bonds Nomenclature CO

carbon monoxide

CO2

carbon dioxide

PCl3

phosphorus trichloride

CCl4

carbon tetrachloride

N2O dinitrogen monoxide

# of Atoms Prefix

1 mono-

2 di-

3 tri-

4 tetra-

5 penta-

6 hexa-

7 hepta-

8 octa-

9 nona-

10 deca-

Covalent Bonds Nomenclature dihydrogen monoxide

H2O

nitrogen dioxide NO2

carbon tetrahydride CH4

# of Atoms Prefix

1 mono-

2 di-

3 tri-

4 tetra-

5 penta-

6 hexa-

7 hepta-

8 octa-

9 nona-

10 deca-

Metallic Bonds Metallic bonds are metal to metal

bonds formed by the attraction between positively charged metal ions and the electrons around them.

Atoms are packed tightly together to the point where outermost energy levels overlap.

This allows electrons to move freely from one atom to the next making them great conductors of electricity.

Ionic Bonds An ion is a charged atom or

molecule. It is charged because the number of electrons do not equal the number of protons in the atom or molecule. Atoms with ADDED electrons

are negative (anions). Atoms with LESS electrons

are positive (cations).

Ionic Bonds The normal charge of an

ion can be quickly determined using the oxidation number of an element. The oxidation number of

an atom is the charge that atom would have if the compound was composed of ions.

Ionic Bonds To find oxidation number:

All elements with a valence number less than four will lose all of their electrons to achieve a full valence or the octet rule. Example:

Beryllium has 2 e- Loses the 2 e- Gains a charge of +2

Ionic Bonds To find oxidation number:

All elements with a valence number greater than four will gain electrons until they have achieved a full valence or the octet rule. Example:

Nitrogen has 5 e- Gains 3 e- Gains a charge of -3

Ionic Bonds Examples:

Oxygen – Group 16 -2

Calcium – Group 2 +2

Aluminum – Group 13 +3

Chlorine – Group 17 -1

Ionic Bonds Ionic bonds are bonds formed by the

attraction between oppositely charged ions. Electrons are transferred from one element to

another.

Potassium (metal – cation) needs to lose 1 valence electron to drop down to a full valence shell. Fluorine (nonmetal – anion) only needs 1

electron to complete its valence shell.









Once the transfer is complete, the potassium will have a +1 charge (K+) and the fluorine will have

a -1 charge (F-).

Once the transfer is complete, the potassium will have a +1 charge (K+) and the fluorine will have

a -1 charge (F-).

The ionic bond is formed because of the electrostatic forces between the positive and negatively charged ions and the new overall

charge is 0.

Magnesium (metal – cation) needs to lose 2 valence electron to drop down to a full valence shell. Iodine (nonmetal – anion) only needs 1

electron to complete its valence shell, but Mg can give to two different atoms.







Once the transfer is complete, the magnesium will have a +2 charge (Mg2+) and each iodine will

have a -1 charge (I-).

Once the transfer is complete, the magnesium will have a +2 charge (Mg2+) and each iodine will

have a -1 charge (I-).

Ionic Bonds Ionic bonds form strong

network structures with high melting and boiling points.

When melted or dissolved in water ionic compounds conduct electricity because ions are free to move.

Ionic Bonds Nomenclature.• Name the cation (metal).

• If the first ion is a transition element other than zinc, cadmium, or silver, you must use a Roman Numeral with the name – we’ll discuss this later.

• Name the anion (nonmetal) by changing the suffix to -ide.

Examples

NaClName the metal ionSodium

Name the nonmetal ion, changing the suffix to –ide.

Chloride

CaO

Calcium Oxide

Al2S3

Aluminum Sulfide

MgI2

Magnesium Iodide

BaNa2 You should recognize a problem with this oneThis is two metals – not a binary ionic compoundThe name of this is BananaBanana (JOKE – haha)

What is the name of this compound:

HIJKLMNO?

WATER – “H” to “O”

You have to admit – that was funny!

Ionic Bonds Nomenclature. To go backwards from

the name to the formula you can use the “Swap and Drop” method.:

1. Write the symbols for each ion.

2. Determine the oxidation number of each ion.

3. Swap and Drop

4. Reduce (if necessary).

5. Rewrite

Ionic Bonds Nomenclature. To go backwards from

the name to the formula you can use the “Swap and Drop” method.:



3. Swap and Drop


5. Rewrite

Polyatomic Ions A polyatomic ion is a group of covalently

bonded atoms that have lost or gained an electron. (Example: Nitrate NO3

- and Ammonium NH4

+). Oppositely charge polyatomic ions can form

compounds. (Example: Ammonium nitrate NH4NO3).

Polyatomic Ions Naming of these

compounds follows the same rules as binary ionic compounds.

The most important part is recognizing there is a polyatomic ion present.

Polyatomic bonds To go from the formula

to the name:

1. Name the cation.

2. Name the anion.

Polyatomic bonds To go from the formula

to the name:

1. Name the cation.

2. Name the anion.

Polyatomic bonds To go from name to formula:



3. Swap and Drop


5. If a subscript greater than one is added to the polyatomic ion use parentheses.

6. Rewrite

Polyatomic bonds To go from name to formula:



3. Swap and Drop


5. If a subscript greater than one is added to the polyatomic ion use parentheses.

6. Rewrite

Transition metals are cations that have variable charges that makes them hard to name.

We use Roman numerals to indicate the charge of a transition metal.

Example: copper (II) oxide – charge of copper is +2 titanium ( IV) sulfide – charge of titanium is +4

Transition Metal Ionic Compounds

To go from formula to name you need to determine the Roman numeral for your transition metal:

1. If there are subscripts present use the reverse “Swap and Drop.”

2. Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY.


To go from formula to name you need to determine the Roman numeral for your transition metal.

1. If there are no subscripts, simply give the transition metal the equal and opposite charge to the nonmetal.

2. Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY.


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“Structure of Matter” Covalent Bonds Ch. 6. Matter Matter is anything that has mass and occupies space. Matter is made of atoms which are the smallest