COMPOUNDS & MOLES

Unit 5

Overview

Naming Ionic Covalent Acids Simple Organic

The Mole Molar Mass Mole Conversions

Calculations Percent

Composition Empirical Formula Molecular Formula

Why do we name compounds? Think of some common compounds that you

know of H2O = water NaCl = table salt CaCO3 = limestone

Imagine if we had to memorize common names for the millions of known compounds that we had today…IMPOSSIBLE!

Standard system was created to name compounds IUPAC (International Union of Pure and

Applied Chemistry)

Chemical Formulas

Indicate the relative numbers of atoms of each kind in a chemical compound

C8H18Indicates 8

carbon atoms

Indicates 18 hydrogen

atoms

Molecular vs. Structural Formulas Molecular Formula

Lists elements in a compound and how many of each element you have

Example: C2H6O

Structural Formula Shows how atoms are “connected” in the

structure Example CH3CH2OH or CH3OCH3

Monatomic Ions

Ions formed from a single atom

Naming cations Simply give the element’s name Example

Ca+2 = calcium ion Na+1 = sodium ion

Naming anions Drop the ending of the element’s name and add “-ide” Example

F-1 = fluoride ion O-2 = oxide ion

Binary Ionic Compounds

Ionic compound composed of 2 elements

Writing Names Name the cation 1st Name the anion 2nd

Example: NaCl = sodium chloride MgF2 = magnesium fluoride Sr3N2 = strontium nitride

Binary Ionic Compounds

Writing FormulasExample: aluminum oxide

Write the symbols for the ions side by side (cation first)

Al+3 O-2

Criss-cross the charges (use absolute value)Al2 O3

Simplify (divide both numbers by largest common factor)

Al2O3

Binary Ionic Compounds

More examples (name to formula)

Calcium nitride = Ca3N2

Potassium sulfide = K2S

Magnesium oxide = MgO

Polyatomic Ions

Electrically charged group of two or more atoms

Oxyanion – polyatomic anion that contains oxygen

General naming rules Most common oxyanion ends in “-ate”

Example ClO3

-1 = chlorate NO3

-1 = nitrate SO4

-2 = sulfate

Polyatomic Ions

The number of oxygen atoms may be altered giving new endings and prefixes to oxyanions

1 more oxygen = per_______ate Common form = _______ate1 less oxygen = _______ite2 less oxygens = hypo_______ite

Example ClO4

-1 = perchlorate ClO3

-1 = chlorate ClO2

-1 = chlorite ClO-1 = hypochlorite

Notice that the charge of the oxyanion does not change (only the number of oxygen atoms)

Polyatomic Ions

Ionic compounds (contain “ions”)

Writing Name If ion comes first, name the polyatomic ion then

name the anion If the ion comes second, name the cation then

name the polyatomic ion (do not change ending) Examples

NH4Cl = ammonium chloride CaSO4 = calcium sulfate Ba3(PO4)2 = barium phosphate

Polyatomic Ions

Writing Formula Follow same rules as binary ionic

compound, but when charges are criss-crossed, use parenthesis to indicate number belongs to entire polyatomic ion

Example: calcium nitrate

Ca+2 NO3-1 = Ca(NO3)2

Stock System (Ionic Compounds) For elements that form two or more cations

with different charges (example Pb+2 and Pb+4) Uses roman numeral to indicate ion’s charge Transition metals, Sn, and Pb use this system

Writing Formulas Roman numeral indicates charge of the cation

(use that to criss cross) Examples

Copper (II) bromide = CuBr2

Iron (III) sulfide = Fe2S3

Tin (IV) phosphate = Sn3(PO4)4

Stock System (Ionic Compounds) Writing Names Use the anion (known charge) that the cation is

bonded to and solve for the charge of the cation Total positive charge (from cation) must equal total

negative charge (from anion)

Example: VF6

Fluorine has a charge of -1 There are six fluorines bonded to the vanadium 6 × -1 = -6 so the charge of vanadium is 6 Name = vanadium (VI) fluoride

Stock System (Ionic Compounds) Example 2: Sn3N2

The charge of nitrogen is -3 There are 2 nitrogen atoms 2 × -3 = -6 There are 3 tin atoms that add up to a charge of +6 +6 ÷ 3 = -2 so each tin atom has a charge of +2 Name = tin (II) nitride

Exception: some transition metals only have one charge (nickel, silver, zinc, etc.) so the roman numeral is omitted

Prefixes

Used in naming covalent compounds

Indicate how many of each atom you have

Number Prefix

1 mono-

2 di-

3 tri-

4 tetra-

5 penta-

6 hexa-

7 hepta-

8 octa-

9 nona-

10 deca-

Binary Covalent Compounds

Writing Names Name the cation followed by the anion (-ide ending) Use prefixes to indicate how many of each atom you

have

Examples: P4Br10 = tetraphosphorous decabromide Si2O5 = disilicon pentoxide

Note If an o or a are doubled, drop the o or a of the prefix Never use mono- on cation (only on anion)

Binary Covalent Compounds

Writing Formulas Prefix indicates how many of each atom

you have Do not criss-cross numbers

Examples: Trinitrogen octachloride = N3Cl8 Arsenic tetrabromide = AsBr4

Summary

Is it ionic?

Is the cation a transition metal, Sn, or Pb?

Use Roman numerals

Name cation then anion

(write it like it is)

Use prefixes (covalent)

NOYES

YES NO

When writing names of formulas…

Acids

Binary acid – contains two elements (one usually hydrogen and the other usually a halogen)

Oxyacid – acids that contain hydrogen, oxygen, and a third element (usually a nonmetal) Usually hydrogen and a polyatomic ion

Acids

Naming binary acids Use form of

hydro_____ic acid

Examples: HF = hydrofluoric acid HCl = hydrochloric acid

Acids

As the number of oxygen atoms changes in oxyacids, so does the name (just like the oxyanions)

1 more oxygen = per_______ic acidCommon form = _______ic acid1 less oxygen = _______ous acid2 less oxygens = hypo_______ous acid

Example HClO4 = perchloric acid HClO3 = chloric acid HClO2 = chlorous acid HClO = hypochlorous acid

Carbon

• Basis for all life.

• Study of carbon compounds is called organic chemistry.

• Can form single, double and triple bonds.

• Long carbon chains can be produced.

• Will bond with many other elements.

• A HUGE number of compounds is possible (organic compounds)

Naming Simple Organic Compounds

Organic compounds containing only carbon and hydrogen are called hydrocarbons

Alkane – all carbons form single bonds Alkene – carbons form double bonds Alkyne – carbons form triple bonds

Whether a compound is an alkane, alkene, or alkyne determines the suffix (ending) in the name of the hydrocarbon

Naming Simple Organic Compounds

Prefix CarbonsMeth- 1Eth- 2Prop- 3But- 4 Pent- 5Hex- 6Hept- 7Oct- 8Non- 9Dec- 10

Number of carbons determines prefix

used in name

Naming Simple Organic Compounds

Examples CH4 = methane C2H6 = ethane

propane

propene

propyne

The Mole

The amount of a substance that contains as many particles as there are atoms in exactly 12 g of 12C SI unit of amount of a substance Abbreviated “mol”

Counting unit just like a “dozen” 1 dozen donuts is the same

amount as 1 dozen books 1 mole of hydrogen atoms is

the same amount as 1 mole of sodium atoms

Avogadro’s Number

6.022×1023 of anything is a mole Named after Italian scientist Amadeo Avogadro Experimentally determined number of atoms in 12 grams of

12C

How big is 602,200,000,000,000,000,000,000? One mole of donut holes would cover the Earth 5 miles deep

in the donut holes One mole of pennies stacked on top of each other would

reach from the Earth to the moon 7 times If you started counting when you were born and never

stopped until the day you died, you would never come close to reaching 6.022×1023

Avogadro’s Number

1 Liter of water contains 55.5 moles of H2O

A 5 lb bag of sugar contains 6.6 moles of sugar

How can that be?!

Atoms and molecules are so tiny that when we use units of moles (6.022×1023) it puts the particles into measurable quantities

Molar Mass

1 mole of hydrogen atoms = 1 mole of sodium atoms

BUT… 1 mole of hydrogen atoms DOES NOT

have the same mass as 1 mole of sodium atoms Individual atoms have different masses

They are the same amount but not the same mass

Molar Mass

The periodic table tells us the mass of 1 mole of any atom

It’s the same as the average atomic mass/relative atomic mass (decimal number on the table)

Molar Mass – mass of 1 mole of an atom or compound Units are “grams/mole” or “g/mol”

Molar Mass

To find the molar mass of a compound, add the molar masses of all atoms in a compound Also called formula mass or molecular mass

(compounds only)

Example: CO2 (1 atom of C and 2 atoms of O)

1 atom C x 12.011 = 12.011 2 atoms O x 15.9994 = 31.9988

Molar mass = 44.010 g/mol

AtomsMolecule

sMole Grams

Mole Relationships

6.02 x 1023Molar Mass

To go between units of grams, moles and atoms (or molecules) use conversions! 6.022×1023 is how many atoms or molecules

are in 1 mole of any substance The molar mass is how many grams are in

one mole of any substance

Mole Conversions

How many grams are in 5.0 moles of calcium?

5.0 mole × = 200.39 g

How many atoms are in 2.1 moles of xenon?

2.1 moles × = 1.26×1024 atoms

40.078 g1 mole

6.022×1023 atoms1 mole

Mole Conversions

There is no way to go straight from grams to atoms or molecules in one step Must use moles as the intermediate step

How many atoms are in 9.8 g of Pb?

9.8g × × = 2.8×1022 atoms

1 mol207.2 g

6.022×1023 atoms1 mole

Mole Conversions

When a conversion includes a compound, it will use the word molecules when a conversion includes an element, it will use the word atoms There are still as many molecules in a mole as there are

atoms

How many grams are in 3.4×1022 molecules of H2O? First solve for molar mass of H2O

(H2O molar mass = 18.02g/mol)

3.4×1022 molecules × × = 1.0 g

1 mole6.022×1023 molecules

18.02 g1 mole

Percent Composition

Percentage by mass of each element in a compound

Example: What is the percent composition of BaSO4?

Ba = 1 × 137.3 = 137.3(137.3/233.4) ×100= 58.8% Ba

S = 1 × 32.1 = 32.1 (32.1/233.4) ×100= 13.8% S

O = 4 × 16.0 = 64.0 (64.0/233.4) ×100= 27.4% O

233.4

Molar Mass part ÷ totalMultipl

y by 100

Total molar mass

Empirical Formula

Smallest whole-number ratio formula of a compound Simplest formula

What is the empirical formula of a compound that is 27.0% sodium, 16.5% nitrogen, and 56.5% oxygen by mass? Assume that you have a 100 gram sample

Na 27.0/22.99 = 1.17 /1.17 = 1

N 16.5/14.01 = 1.18 /1.17 = 1

O 56.5/16.00 = 3.53 /1.17 = 3

EmpiricalFormula = NaNO3

Molar MassDivide by

smallest

number

Empirical Formula

When numbers are too far to round, you may need to multiply all values by the same factor to make all numbers whole

What is the empirical formula of a compound that contains 40.6g of calcium and 9.5g of nitrogen?

Ca 40.6/40.1 = 1.01 /0.69 = 1.5 × 2 = 3

N 9.5/14.01 = 0.69 /0.69 = 1 × 2 = 2EmpiricalFormula = Ca3N2

too far to round

double both numbers to get whole numbers

Molecular Formula

Indicates actual number of atoms of each element in a compound Multiple of empirical formula

An empirical formula can be the molecular formula, but the molecular formula is not always the empirical formula

CH4

Empirical Formula

C3H12

Molecular Formula

Molecular Formula

If the molecular mass is known, you can solve for the molecular formula

The molar mass of a compound with empirical formula of CH2O is 180.12 g/mol. What is the molecular formula of this compound?

Molar mass CH2O = 30.02g/mol = 6

Molecular Formula = CH2O × 6 = C6H12O6

180.1230.02