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CRHS Academic ChemistryUnit 6 - Nomenclature

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Notes, Homework, Exam Reviews and Their KEYS located on CRHS Academic Chemistry Website: https://cincochem.pbworks.com

6.1 Introduction to Bonding

Chemistry

https://cincochem.pbworks.com/

of 20 Unit 6 Notes

There are millions of known chemical substances. Without a universal, international system, naming them reliably would

be __impossible___. The system used in naming substances is called chemical nomenclature.

Chemical ____nomenclature____ is a set of ____rules_____ used to generate systematic _______________ for chemical

compounds.

The governing body responsible for chemical nomenclature is the International Union of Pure and Applied Chemistry

(___IUPAC__). There are two major divisions to the IUPAC system:

1. Organic nomenclature – contain carbon, usually bonded with hydrogen, oxygen, nitrogen, and sulfur. Organic

chemistry and naming organic compounds is a topic for a college chemistry course.

2. Inorganic nomenclature – everything else is inorganic and is divided into three categories…

__Covalent__ (a.k.a. molecular) compounds

__Ionic____ compounds

__Acids_____

This unit covers inorganic nomenclature...

Overview of Chemical Bonding

A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that

binds the atoms together.

Valence electrons are the electrons in the outermost energy level (S and P sublevels)

Atoms form compounds by gaining, losing or sharing electrons.

When atoms bond, the valence electrons are redistributed in a way that makes the atoms more stable.

Octet Rule – Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an

octet of electrons in its highest occupied energy level.

Type of Bonds

The way in which electrons are redistributed determines the type of bond.

1. Covalent Bonds – Chemical bonding that results from the sharingof electron pairs between two atoms.

Unit 6 Notes of 20

2. Ionic Bonds – Chemical bonding that occurs from the electrical attraction between positively charged cations and

negatively charged anions.

Ionic Bonds

Ionic bonds are formed between a CATION and an ANION when electrons are gained or lost due to forces between

the elements.

o The CATION has lost one or more electrons and the ANION has gained one or more electrons.

o Compounds formed with ionic bonds between a metal cation and a nonmetal anion are called

_____ionic_____ compounds.

o Examples of ionic compounds:

NaCl – sodium chloride MgO – magnesium oxide RbS – rubidium sulfide

o Ionic compounds can also be formed between metals cations and ___ polyatomic ___ anions. A list of

polyatomic ions and their charges are found in the inside back cover.

o Examples of polyatomic ionic compounds (PT):

Na2SO4 – sodium sulfate CaClO3 - calcium chlorate AgNO3 – silver nitrate

o _Acids__ are compounds that form from the H+ cation and a nonmetal anion or a polyatomic anion.

o Examples of acids:

H2SO4 – sulfuric acid HCl – hydrochloric acid

Covalent Bonds

Covalent bonds are formed between two nonmetals and involve the __sharing__ of electrons between atoms.

o Examples of covalent compounds:

CO2 – carbon dioxide N2O5 - calcium chlorate AgNO3 – silver nitrate

of 20 Unit 6 Notes

Review of Periodic Table and Oxidation Numbers

We studied the historical development, organization and theory of the Periodic Chart in Unit 5. In this unit we utilize

the following key facts about the Periodic Chart that are needed to name chemical compounds.

Metals – tend to _lose electrons (e–) to form CATIONS (+ charge)

Cation – atom that has _lost______ one or more electron(s)

Nonmetals – tend to __gain___ electrons (e–) to form ANIONS (- charge)

Anion – atom that has ___gained______ one or more electron(s)

Metalloids – have properties of metals and nonmetals and are found along and to either side of the zig-zag line

(except _Al__).

Groups 1,2, and Al metals and Group 15-18 nonmetals have defined oxidation numbers. Oxidation numbers indicate

that element’s ability to donate electrons or to accept electrons.

_ Negative _____ oxidation numbers indicate that electrons are accepted by that element. _Nonmetals_ have

negative oxidation numbers.

__Positive_ oxidation numbers indicate that electrons have been donated. __Metals__ have positive oxidation

numbers.

Groups 3-12 are called ____transition_ _metals_____ and have __variable____ , but positive, oxidation numbers.

Weak Metals just to left of Zig Zag Line also have variable positive oxidation numbers.

1 18

+1

2 13 14 15 16 17

0

+2+3

+2 or +4 or-4or?

-3 -2 -13 4 5 6 7 8 9 10 11 12

+?Tend to have more than 1 oxidation number

Many are +2

Common Oxidation States

Unit 6 Notes of 20

+3 or +4

+3

Review Practice: Fill in the missing information

Group Number Group Name Oxidation Number Gain or Lose Electrons

How Many Electrons Gained or Lost?

Anion or Cation

1 Alkalai Metals +1 lose 1 Cation

2 Alkali Earth Metals +2 lose 2 Cation

13 ---- +3 lose 3 Cation

14 ---- +4,-4 Lose/Gain 4 -----

15 ---- -3 gain 3 Anion

16 ---- -2 gain 2 Anion

17 Halogens -1 gain 1 Anion

18 Noble Gases 0 Neither gain nor

lose electrons

0 -----

3-12 Transition Metals +1, +2, +3, +4 Gain 1, 2, 3, 4 Cation

of 20 Unit 6 Notes

6.2 NAMING AND WRITING FORMULAS OF IONIC COMPOUNDS - GROUP 1-3 METAL CATIONS

___Ionic_____________ compounds are made of a _____metal_______________ cation (+) and a

_______nonmetal____________ anion (–).

First element is from the __left____ of the zig-zag line (METAL)

Second element is from the RIGHT of the zig-zag line (NONMETAL).

The ___metal_______ is always written first.

The Ionic Naming Rules

1) Confirm the compound is ionic

2) The number of atoms _does_not__matter___ and there are __No _ __ PREFIXES.

3) Name the ___cation____________ (metal).

4) Name the ___anion____________ (nonmetal), using the root of the element and changing the ending to

___-ide____.

Examples: NaBr sodium bromide

Practice:

K3N potassium nitride

Ca3P2 calcium phosphide

SrO strontium oxide

LiCl lithium chloride

STEPS to Writing Formulas1. Identify the first term. This will be a __metal___ cation (+) and written exactly as it is on the periodic table.

2. The second term is a _____nonmetal_______ anion (–) and its ending has been changed to “___ide______”.

Determine the element name from the root of the word.

Unit 6 Notes of 20

3. Write the cation and anion by finding their __charges_____.

4. Use the ___criss-cross_____ method.

5. __Reduce______ subscripts to lowest whole numbers.

Example:

Barium Iodide: Ba I → Ba+2 I–1 → Ba1I2 → BaI2

Practice

magnesium fluoride MgF 2

aluminum iodide AlI 3

calcium sulfide CaS

lithium oxide Li 2O

beryllium phosphide Be 3P2

francium sulfide Fr 2S

of 20 Unit 6 Notes

6.3 NAMING AND WRITING FORMULAS OF IONICS COMPOUNDS – POLYATOMIC IONS

Ionic compounds also form with groups of nonmetals acting as a cation or anion called Polyatomic Ions. (ex. NO3 –1,

nitrate)

Naming Rules for Ionic Compounds with Polyatomic Ions

1) __Never___ change the name of a polyatomic ion (PI)

2) This is ionic bonding, so NO PREFIXES.

3) If the PI comes _first_, name it, then name the anion by changing the ending to “ide”. (Ex. NH4Cl - ammonium

chloride)

4) If the PI comes ___last___, name the cation and then the PI. No changes necessary. (Ex. Na2SO4 – sodium

sulfate)

5) All but one PI are ANIONS (–). Ammonium, NH4+ is a cation.

Examples: Mg(MnO4)2 magnesium permanganate

NH4Br ammonium bromide

Mixed Practice –indicate I (nonmetal anion) or IP (polyatomic ion). Name it.

1) NaNO3 IP sodium nitrate

2) BaCl2 I barium chloride

3) NH4ClO4 IP ammonium perchlorate

4) LiCl I lithium chloride

The most common PI’s list is found at the back of your packet. You will always be given this list, but you need to know how to use it!

Unit 6 Notes of 20

Steps to Writing Formulas of Ionic Compounds with Polyatomic Ions

STEPS1. Identify the cation (+), and the anion (–)

a. the cation, anion, or BOTH may be polyatomic ions.

2. Write the cation and anion by finding their ___charges_____

a. Place parentheses around the PI to remind you to ___NEVER CHANGE___________ a polyatomic ion!

3. Use the ____criss-cross_________ method.

4. ___Reduce______ subscripts to lowest whole numbers.

Examples:

Ammonium Bromide: NH4 Br → (NH4)+1 Br–1 → (NH4)1Br1 → NH4Br

Magnesium Nitrate: Mg NO3 → Mg+2 (NO3)–1 → Mg1(NO3)2 → Mg(NO3)2

Mixed Practice: Write I, IP(polyatomic). Write the formula.

sodium cyanide IP NaCN

radium fluoride I RaF2

barium oxide I BaO

aluminum nitrite IP AlNO 2

calcium sulfite IP CaSO 3

aluminum dichromate IP Al2(Cr2O7)3

ammonium silicate IP ( NH4)2SiO3

of 20 Unit 6 Notes

6.4 NAMING AND WRITING FORMULAS OF IONIC COMPOUNDS - TRANSITION METAL CATIONS

Periodic table __transition_______________________ metals (and nearby poor metals) also form ionic compounds with

nonmetals and PI’s.

Transition metals are __metals_________, which means they can exhibit __>1_____ oxidation state, or charge:

ex. Fe+2, Fe+3.

They are named according to their cation charge using the __Stock___________ system.

The Stock System is a _________________ numeral used to indicate the _____________________ of a

multivalent ion.

Roman Numerals 1 through 6

1 = I 2=II 3=III 4=IV 5=V 6=VI

Naming Rules for Ionic Compounds with Transition Metals

1) From the chemical formula, find the charge of the _anion____.

2) _Multiply________ the anion charge by its subscript to calculate the overall negative charge of the compound.

3) The cation(s) __equals_____ this charge to make the compound’s net charge equal zero (neutral).

4) Divide the overall positive charge by the subscript on the cation to find the charge on each metal cation.

5) Write the Roman numeral of the numerical charge in parentheses _after____________ the transition metal name.

6) Normal __ionic___ naming rules apply.

7) Example Fe2O3

Fe2O3 Fe2O3 Fe2O3 Iron (III) Oxide

Examples:

+6 by 2 (atoms of Fe) = each Fe has a +3 charge, so…

–2

–6

–2

–6+6

The 3 oxygens have a total–6 charge

–2+3

+6 –6

Unit 6 Notes of 20

CuO copper(II)oxide

FeS iron(II)sulfide

TiCl4 titanium (IV) Chloride

Ni(OH)3 nickel (III) hydroxide

Mixed Practice – indicate I, or IP, or IT (transition metal) or IPT (polyatomic & transition metal). Name it.

1) VCl3 IT vanadium (III) choride

2) BaSO3 IP barium sulfite

3) FeBr2 IT iron (II) bromide

4) CuSO4 IPT copper (II) sulfate

5) NaI I sodium iodide

Steps to Writing Formulas for Ionic Compounds with Transition metals

1) Identify the transition metal and the anion.

2) Form the cation and anion by finding their _charge_______

a. cation charge is the ____Roman Numeral________________________

__________________________________________

3) Use the ___criss-cross_________________________ method to form the compound.

___Reduce__________________ subscripts to lowest whole numbers

Example:

Manganese(IV) oxide: Mn O → Mn+4 O–2 → Mn2O4 → MnO2

Mixed Practice: indicate I, or IP, or IT (transition metal) or IPT (polyatomic & transition metal). Write the formula.

tin (IV) iodide IT SnI 4

mercury (II) sulfide IT HgS

lithium chloride I LiCl

of 20 Unit 6 Notesiron (III) oxide IT Fe 2O3

manganese (II) chloride IT MnCl 2

chromium (VI) phosphate IPT Cr 3(PO4)4

copper (III) permanganate IPT Cu(KMnO 4)3

Extra Mixed Practice: indicate I, or IP, or IT (transition metal) or IPT (polyatomic & transition metal). Write the formula.

barium fluoride I BaF

aluminum phosphate IP Al 2(PO4)3

sodium acetate IP NaOC 2H3O2

lead (II) oxide IT PbO

sodium hydroxide IP NaOH

calcium bromide I Ca(OH) 2

potassium dichromate IP K 2Cr2O7

phosphorus tribromide

manganese (III) oxide IT Mn 2O3

sodium nitride I Na 3N

iodine trichloride

tin (II) oxide IT SnO

Unit 6 Notes of 20

6.5 NAMING AND WRITING FORMULAS OF ACIDS

Acid – a compound that produces ___hydrogen ions______, H+, when dissolved in water. ALL acid formulas begin with

__H_______.

Naming Binary Acids - hydrogen + 1 nonmetal; example: HCl

STEPS1) Begin with “_hydro_______________________”.

2) Use the ____root______________________ of the element name.

3) Change the ending to ____ic_______________.

4) End with ____acid______________________.

Examples: HBr hydrobromic acid

HI hydroiodic acid

Naming Ternary Acids – hydrogen + a polyatomic ion, ex: H2SO3

STEPS1) ____Never _ _____________________ use “hydro”.

2) Use polyatomic ion name ___without________ ending.

3) Change to __ic______ if polyatomic ion ends in __ate_____

4) Change to ___ous_____ if polyatomic ion ends in ___ite_______

5) End with _____acid_____________.

Examples: H3PO3 phosphoric acid

H2CO3 carbonic acid

Mixed Practice – indicate I, or IP, or IT (transition metal) or IPT (polyatomic & transition metal), Ab or At (Acid). Name it.

1) KBr I potassium bromide

2) CaCO3 P calcium carbonate

I ate an icky

white mouse

of 20 Unit 6 Notes3) HNO3 At nitric acid

4) MnO2 IT manganese (IV) oxide

5) Cu(NO3)2 IPT copper (II) nitrate

6) RhCl4 I rhodium chloride

7) H3PO4 At phosphoric acid

8) HNO2 At nitrous acid

9) HCl Ab hydrochloric acid

Writing Formulas of Binary Acids

An acid name ends in the word ________acid______________!

An acid formula begins with ______H______________!

Binary Acids - If it begins with “_____hydro___________”, it is binary, meaning hydrogen bound to one

______nonmetal___________.

STEPS

1. Identify the __nonmetal_____________________ in the acid name.

2. Use oxidation state trends to find the ______charges____of hydrogen (always __+1____) and the nonmetal.

3. Write “H” and the symbol of the nonmetal.

4. Write the charge as a ____superscript____________ of each ion. This forms the cation (+) and an anion (–)

5. _____Bring________________________ (just like IONIC BONDING) the number of each ____superscript____

down to the subscript of the opposite ion. Omit the__1’s_________!

6. ____Reduce_____________ subscripts to lowest whole numbers ratio.

Example:

Hydroselenic Acid→ H Se → H+1 Se–2 → H2Se1 → H2Se

Unit 6 Notes of 20

Writing Formulas of Ternary Acids

Ternary acid - If it does not begin with “hydro”, it is ternary, meaning hydrogen bound to a _______polyatomic

anion_______________.

STEPS

1. Identify the ____ polyatomic ion . ___.

a. If it ends in “-___ic___” acid, change the end to “-___ate___”, look up on the common polyatomic ion list.

b. If it ends in “-___ous____” acid, change the end to “-____ite_____”, look up on the common polyatomic ion

list.

2. Write the cation and anion by finding their ___charge__________

3. _Criss-Cross________the numbers only

4. ___Reduce________ subscripts to lowest whole number ratio.

Example:

Carbonic Acid → H CO3 → H+1 (CO3)–2 → H2(CO3)1 → H2(CO3)

Practice:

hydrofluoric acid Ab HF

sodium sulfide I Na 2S

chlorous acid At HClO 2

hydroselenic acid Ab H 2Se

sulfuric acid At H 2SO4

acetic acid At HC 2H3O2

hydrosulfuric acid Ab H 2S

“ic” I “ate” it Carbonate

of 20 Unit 6 Notes

6.6 NAMING AND WRITING FORMULAS OF COVALENT COMPOUND

Binary ___covalent_________ compounds are made of TWO___nonmetals_______(Ex. P2O5) bound together. Nonmetals

are ___right____ of the zig-zag line (except _hydrogen___).

Covalent Compound Naming Rules

1) Confirm the compound is __covalent________.

2) Check the number of atoms of each element by the ___subscript___. If there is no subscript, there is one atom.

Ex. P2O5 – __2__ phosphorus, _5___oxygen

3) Name the _first_____ element using the prefix (page bottom) with the element name. (Ex. P2O5 –

diphosphorus…)

NEVER use “_mono___” on the first element.

4) Name the ___second________ element. Choose the prefix, write the root of the element, and change the

ending to “___ide_________”. (Ex. P2O5 – diphosphorus pentoxide…)

Practice:1) SF6 sulfur hexaflouride

2) CS carbon sulfide

3) S3N2 trisulfur dinitrogen

Writing Formulas for Covalent Compounds

STEPS1. Check the elements in the compound. Confirm they are both ____nonmetals___________________________.

2. Use the ______ prefixes ________________________ to determine how many _______________ of each

element are in the compound.

Mono = 1 Di = 2 Tri = 3 Tetra = 4 Penta = 5

Hexa = 6 Hepta = 7 Octa = 8 Nona = 9 Deca = 10

Unit 6 Notes of 20

3. Use _____subscripts_____________________ to indicate how many atoms of each element are in the formula

4. Example: carbon dioxide has __1__ carbon(s) and __2_ oxygen(s) = C1O2 = CO2 (1 subscripts are understood)

5. Example: nitrogen trifluoride NF3

Practice:1) selenium tetrafluoride SeF 4

2) octaphosphorus pentabromide P 8Br5

3) heptanitrogen hexachloride N 7Cl6

of 20 Unit 6 Notes

Mixed Practice

Formula Name Type of Compound

P2O5 diphosphorus pentaoxide C

H3N hydronitric acidAb

KClO2 potassium chlorite IP

FeCl2 Iron (II) chloride IT

Ba(OH)2 barium hydroxide IP

(NH4)2Se ammonium selenideIP

H3N hydronitric acid Ab

CCl4 carbon tetrachloride C

PbI2 lead (II) iodide IT

HNO2 nitrous acid At

HClO hypochlorous acid At

CaSO3 Calcium sulfite IP

CaSO4 calcium sulfate IP

Ra3(PO4)2 radium phosphate IP

Na3N sodium nitrideI

Li2O lithium oxide I

H3PO4 phosphoric acid At

(NH4)2S ammonium sulfide IP

Fe2O3 iron (III) oxide IT

Mg3P2 magnesium phosphide I

KCl potassium chloride I

BaSO4 barium sulfate IP

H3PO3 phosphorous acid At

PI5 phosphorous pentiodide C

SO2 sulfur dioxide C

NO nitrogen monoxide C

Unit 6 Notes of 20

Prefixes for Covalent Compounds

Mono = 1 Di = 2 Tri = 3 Tetra = 4 Penta = 5

Hexa = 6 Hepta = 7 Octa = 8 Nona = 9 Deca = 10

Common Polyatomic Ions List

Name Ion Name Ionacetate C2H3O2

– or CH3COO– hypochlorite ClO–

ammonium NH4+ nitrate NO3

–

carbonate CO32– nitrite NO2

–

chlorate ClO3– perchlorate ClO4

–

chlorite ClO2– permanganate MnO4

–

chromate CrO42– phosphate PO4

3–

cyanide CN– phosphite PO33–

dichromate Cr2O72– silicate SiO3

2–

hydrogen carbonate HCO3– sulfate SO4

2–

hydroxide OH– sulfite SO32–

Roman Numerals for Transition Metals 1-6

1 = I 2 = II 3 = III 4 = IV 5 = V 6 = VI

20 Common Polyatomic Ions

of 20 Unit 6 Notes

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