CHE 120

chapter 3 - Chemical Formulas & Nomenclature


Chemical Formulas

A chemical formula tells you which elements are present and how many atoms of each element are present

Subscripts - number or atoms

Ex: H2O -> 2H, 1O

  • if there is no subscript then there is only 1 atom


Ionic vs. Covalent

Ionic: metal + nonmetal

Covalent/molecular: nonmetal + nonmetal


Ionic compounds form large repeating structures called formula units

Covalent compounds exist as indivisual molecules

Ones that naturally occur as pairs: H2, N2, F2, Cl2, Br2, I2


Allotropes are different forms of the same element

Ex: O2 and O3 form oxygen


Naming Binary Covalent Compounds

Binary = exactly 2 elements

covalent = 2 nonmetals

Prefixes to show how many atoms there are

1 =   mono

2 =   di

3 =   tri

4 =   tetra

5 =   penta

6 =   hexa

7 =   hepta

8 =   octa

9 =   nona

10 = deca

Naming Rules

  1. Name the first element

  2. Add a prefix only if there is more than 1

  3. Name the second element with -ide

  4. ALWAYS use a prefix for the second element

Ex: C = Carbon, 2 O = dioxide, Carbon dioxide


Naming Acids

If formula starts with H, it is probably an acid

Binary Acids

H + one other nonmetal

Naming pattern: hydro + root + ic acid

Ex: HCl -> hydrochloric acid

Oxyacid rules

Look at the end of the ion

-ate -> ic acid

Ex: nitrate -> nitric acid

-ite -> ous acid

Ex: nitrite -> nitrous acid

MEMORIZE:

ate -> ic

ite -> ous

ide -> hydro...ic


Nomenclature

when given a formula, first figure out what type of compound it is

metal present? -> usually ionic

two nonmentals? -> covalent

starts with H? -> usually acid


ionic naming

cation + anion

NO prefixes

Ex: NaCl -> sodium chloride, Ca,Cl2 -> calcium chloride


covalent naming

use prefixes

Ex: CO -> carbon monoxide, CO2 -> carbon dioxide


Metals with multiple charges

use roman numerals

Ex: FeCl2 -> iron(ll) chloride, FeCl3 -> iron(lll) chloride


The Mole

mole = giant counting unit

1 mole = 6.022 x 10^23 particles

this number is called Avogardo's number

Use moles to count:

  • atoms

  • molecules

  • formula units


Important conversion

moles <-> particles

use:

6.022 x 10^23 - mol

so:

particles -> divide by 6.022 x 10^23 -> moles

moles -> multiply by 6.022 x 10^23 -> particles


Moles of atoms in a compound

look at the subscripts

Ex: 1 mol C12H22O11 contains:

  • 12 mol C

  • 22 mol H

  • 11 mol O


Percent Composition

the percentage by mass of each element in a compound

Formula:

% element = (mass of element / total mass of compound) x 100

Ex: NaCl

Na = 22.99 g/mol

Cl = 35.45 g/mol

total: 22.99+35.45=58.44 g/mol

then:

%Na = 22.99 / 58.44 x 100

%Cl   = 35.45 / 58.44 x 100

ALL ELEMENT PERCENTAGES SHOULD ADD UP TO AROUND 100


Molecular Formulas

Empirical vs. molecular


Empirical formula

= simplest whole-number ratio


Molecular formula

= actual number of atoms on a molecule

Ex: glucose

molecular -> C6H12O6

empirical -> CH2O

(since C6H12O6 can be divided by 6)


Key equation

molecular formula = empirical formula x whole number

to find muliplier: molar mass of compound / molar mass of empirical formula

Ex:

if

  • compound = 104 g/mol

  • empirical formula = CH = 13 g/mol

104 / 13 = 8

CH x 8 = C8H8


Combustion Analysis

used to determine the formula of a compound containing C and H (and sometimes O)

the compound is burned

  • carbon -> CO2

  • hydrogen -> H2O

Basic Process

  1. find moles of CO2 -> find moles of C

  2. Find moles of H2O -> find moles of H

  3. if O is present, find its mass by subtraction

  4. convert each element to moles

  5. find simplest whole-number ratio -> empirical formula

  6. of molar mass is given, find molecular formula