Molecules, Compounds, and Chemical Equations

 When two or more elements are combined chemically, they are called a compound .
 Just like letters of the alphabet can be combined to form an endless number of words, elements can be combined to form an endless number of compounds.
 Life would not possible without the diversity of compounds.

 Hydrogen, a diatomic and colorless gas
 Oxygen, a diatomic and colorless gas
 Water, a clear liquid which is essential for most life forms

 Most elements are not found by themselves in nature, rather they occur in compounds like NaCl or
MgO
 Some occur as diatomic molecules:
N
2
, O
2
, F
2
, Cl
2
, Br
2
, and I
2
.
 Others like P and S are polyatomic.
 Only the metals of Gold, Copper, and
Silver can sometimes be found in their pure state.

 Compounds are held together by one of two types of bonds.
 Ionic bonds are the electrostatic attraction between cations and anions.
 Covalent bonds are when two atoms share at least a pair of electrons.

 A chemical formula is used to describe our compounds.
 Uses subscripts after each element in the formula.
 Subscript of “1” is implied if no subscript.
 Empirical formula = simplest whole number ratio of the elements in compound.
 Molecular formula = actual number of atoms bonded together.
 Structural formula = shows how the atoms are bonded together with lines representing the bonds.

 A model may be used to represent a molecule.
 Ball-and-stick
 Space-filling
 No model is completely accurate!

• Atomic elements = elements whose particles are single atoms
• Molecular elements = elements whose particles are multi-atom molecules
• Molecular compounds = compounds whose particles are molecules made of only nonmetals
• Ionic compounds = compounds whose particles are cations and anions

Propane – contains individual C
3
H
8 molecules
Table salt – contains an array of Na + ions and Cl ions

 Classifying each as either an Atomic element,
Molecular element, Molecular compound, or Ionic compound.



A) Formaldehyde, CH
2
O
B) Red phosphorous, P
4
C) Cobalt, Co
 D) Magnesium chloride, MgCl
2

 Ionic compounds – are substances that contain both cations and anions.
 Formulas are always empirical – lowest whole number ratio of cations and anions.
 Writing ionic formulas – charges must cancel to yield a neutral species.
 Most common charges for cations = +1, +2, or +3
 Most common charges for anions are -1, -2, or -3.

-1
-2
-3
M , X +1 +2
MX
M
2
X
M
3
X
MX
2
MX
M
3
X
2
+3
MX
3
M
2
X
3
MX

 Learning Check
 Write the correct formula between the ions of…
 A) Ca and Br
 B) K and S
 C) Li and N
 D) Al and O
 E) Mg and P

 Chemical Nomenclature - the systematic method of naming chemical compounds.
 Different system for naming Ionic and Molecular compounds.
 Ionic = cation (usually a metal) with an anion (usually a non-metal.
 Molecular = two or more non-metals or metalloid with a non-metal.
 Identify compound BEFORE going to the rules!!!

 Learning Check…
 Identify as Ionic or Molecular
 A) Na
2
S
 B) PCl
3
 C) SiH
4


D) FeBr
E) ZnO
3
 F) CCl
4

 An ion composed of two or more elements with a net charge.


Only one is a cation – NH
4
+ = ammonium ion.
All the rest are anions.
 Any time a polyatomic ion is present, then the compound is ionic.



Ex) Na
2
SO
4
, MgCO
3
, NH
4
Cl
When more than one polyatomic ion is needed to balance charges, then use (poly) x
.
Ex) Al(NO
3
)
3
, (NH
4
)
2
S

 Two subcategories:
 Metals with only one valence (mostly the main group metals)
 Metals with more than one possible valence charge.

1.
Metals with only one valence
Name the metal first
2.

Name the non-metal second and change its suffix to
ide.
For polyatomic ions, they ALWAYS keep their same name – DO NOT CHANGE TO ide ending.

 What is the name of…
 A) NaBr
 B) MgSO
4
 C) K
2
S
 D) Li
3
PO
4


1.
2.

Metals with more than one valence
Name metal first followed by its valence in Roman
Numerals and in ( ). This means that you will have to figure out the charge for that metal based on what it is bonded to.
Name the non-metal second and change the suffix to ide.
As before, polyatomics remain the same.

 What is the name of…
 A) Cu
2
O
 B) Fe
2
S
3
 C) Mn(NO
3
)
2
 D) Co
3
(PO
4
)
2


1.
2.
3.

Remember, these contain only non-metals or a metalloid with a non-metal.
Name the first non-metal in the formula.
Name the second non-metal in the formula and change the suffix to ide.
Add prefixes for all subscripts - except if the first one is a “1”.
1 = mono, 2 = di, 3 = tri, 4 = tetra, 5 = penta, and 6 = hexa.

 What is the name of…
 A) CS
2
 B) PCl
5
 C) AsBr
3
 D) N
2
O
4
 E) N
2
O

Compound
Ionic or
Molecular?
Ionic Molecular
Metal
Fixed charge
Metal
Variable
Charge

 What is the name of…
 A) NiCl
2
 B) K
2
CrO
4
 C) Cl
2
O
 D) SF
4
 E) (NH
4
)
2
SO
4

 Acids are molecular compounds that release an H + when added to water.
 Formula starts with an “H”.
 Binary acids = H with one other element.
 Hydro + base name of non-metal + ic + acid
 HCl = hydrochloric acid
 Oxoacids = HXO y
 Polyatomic ion ends in –ate, then suffix changes to –ic
 Polyatomic ion ends in –ite, then suffix changes to -ous

 What is the name of…
 A) HBr
 B) HNO
3
 C) H
2
SO
3
 D) H
3
PO
4

 A chemical formula also has a quantitative aspect.
 A Formula Weight for an element or compound is found using the periodic table.
 Formula weights can refer to a single element’s weight or an ionic compound.
 Molecular weight refers to a molecular compound’s weight.
 Weights from periodic table should be rounded to the nearest 0.1 amu at the bare minimum!

 The molar mass of any compound is equal to the sum of the atomic weights expressed in grams.
 Ex) The molar mass of CO2 is 44.0 grams.
 Thus, one mole of CO2 = 44.0 grams.
 1 mol CO2 = 6.02 x 10 23 molecules.
 Calculate the molar mass for…
 A) PCl
3
 B) C
6
H
12
O
6
 C) Fe
2
(SO
4
)
3

 A molar mass can be used to convert grams to moles or moles to grams.
 Calculate:
 A) the moles present in 2.85g of CO
2
 B) the grams present in 0.552 moles of NH
3
 C) the number of molecules present in 0.255g of H
2
O
 D) the number of O atoms in 3.00g of C
6
H
12
O
6

 A formula weight can be used to calculate the mass percentage of any element in the formula by:
Mass % of A = formula weight of compound
 This is one place to also test nomenclature!
 Find the mass percent of each element in the compound Calcium nitrate.


1.
We can use moles to find an empirical (simplest) formula from mass percentages by:
Assume a 100 gram sample (%
 grams).
2.
3.
4.
Convert grams of each element to moles use the formula weights.
Divide each mole amount by the smallest one.
Using a multiplier to eliminate fractions like: 0.25, 0.33,
0.50, 0.67, and 0.75.

 A compound contains 17.6% Na, 39.7% Cr, and 42.7%
O by mass. What is its empirical formula?

 An empirical formula may not be the actual formula since molecular formulas do not have to be the lowest whole number subscripts.
 The multiplier, n, can be found if we know the overall molecular weight of the compound.
n = molecular weight empirical weight
 A compound is contains 40.9% C, 4.6% H, and 54.5%
O by mass with a molecular weight of about 176g/mol.
What is the molecular formula for this compound?

 Combustion of a 0.8233 g sample of a compound containing only carbon, hydrogen, and oxygen produced the following:
CO
2
= 2.445 g
H
2
O = 0.6003 g
 Determine the empirical formula of the compound

 Chemical reactions are represented in a concise method by a chemical equation.
 Ex) 2 H2(g) + O2(g) 2 H2O(l)
Reactants
Products

 Ex) 2 H2(g) + O2(g) 2 H2O(l)
Phase Symbols
Coefficient

 A subscript in a chemical formula tells us how many of each type of atom are in the compound.
 Ex) C
6
H
12
O
6
 Subscripts cannot be altered!!!
 Atoms can be created nor destroyed in a chemical reaction.
 Thus, we balance a reaction by adding coefficients in front of each substance.

 Balance by inspection.
 Use a tally sheet.
 Start with elements that occur once on each side.
 Combustion – do C, then H, then O.
 Balance each reaction…
 __ Al + __ Cl
2

__ AlCl
 __ H
3
PO
3

__ H
3
PO
4
+ __ PH
3

 __ Na
3
PO
4
+ __ CaBr
2

__ Ca
3
(PO
4
)
2
+ __ NaBr
 __ C
3
H
8
+ __ O
2

__ CO
2
+ __ H
2
O
 __ C
4
H
8
O + __ O
2

__ CO
2
+ __ H
2

 In the 18 th century, compounds from living things were called organic; compounds from the nonliving environment were called inorganic
 Organic compounds easily decomposed and could not be made in the 18 th -century lab
 Inorganic compounds very difficult to decompose, but able to be synthesized

 Today we commonly make organic compounds in the lab and find them all around us
 Organic compounds are mainly made of C and H, sometimes with O, N, P, S, and trace amounts of other elements
 The main element that is the focus of organic chemistry is carbon


Carbon atoms bond almost exclusively covalently
 compounds with ionic bonding C are generally inorganic
 When C bonds, it forms four covalent bonds
 4 single bonds, 2 double bonds, 1 triple + 1 single, etc.
 Carbon is unique in that it can form limitless chains of C atoms, both straight and branched, and rings of C atoms

 There are two main categories of organic compounds, hydrocarbons and functionalized hydrocarbons
 Hydrocarbons contain only C and H
 When all C-C bonds are single bonds, they are called alkanes.
 Formula for any alkane is:
C n
H
2n+2
.



CH
4
= Methane
C
2
H
6
= Ethane
 C
3
H
8
= Propane

 C
5
H
12
= Pentane

C
4
H
10
= Butane


C
6
H
14
= Hexane
C
7
H
16
= Heptane
C
8
H
18
= Octane

Family Name ends w/ Generic Example
Alcohol
Ether
Aldehyde
Ketone
Carboxylic Acid
Ester
Amine
Name