Download Unit 3: Bonding and Nomenclature Content Outline: Chemical

Unit 3: Bonding and Nomenclature
Content Outline: Chemical Bonds (3.1)
I.
Chemical Bonds
A. These are mutual attractions between atom’s valence shell electrons to create new molecules.
B. Bonds can affect the electrical charge of an atom or molecule.
C. Bonding of atoms promotes atomic or molecular stability in nature. (It mimics the stability of Noble
gas elements.)
1. Potential Energy (PE) per atom decreases in a bond. More stable atoms mean less potential
Energy per molecule.
2. Molecules release energy when a bond is created; and to break apart molecules require energy
input.
D. There exists three types of chemical bonds:
1. Covalent Bonds
a. The atoms valence shells “share” electrons to make each atom more stable. Each behaves
like a Noble gas. (“co” means together and “valent” means “of the valence”)
b. Two types of molecules can result from this type of bonding.
i.
Non-polar molecule
α. In this type of molecule, there is an equal distribution of electrons
b. These molecules typically have electronegativity (desire for electrons) differences
ranging from 0 (no desire) to 0.3 (minimal desire).
ii.
Polar molecules
α. These molecules have unequal distribution of electrons within the molecule.
* This makes one end (pole) positive (less electrons).
* This makes one end (pole) negative (more electrons).
b. These molecules typically have electronegativity differences ranging from
0.3 –1.7.
c. The negative end is represent by (δ-); the positive end is (δ+)
d. Water is the most important polar molecule on the planet. This polar quality helps
support life in all organisms.
c. Covalent bonds tend to be stronger than all other bonds, but not always.
2. Ionic Bonds
a. This bond forms from the electrical charge attraction (positive to a negative) between
oppositely charged atoms.
i.
The charge on the atom is the result of an “electron swap” between the two atoms.
a. This swap allows both atoms to behave “like” a Noble gas.
b. The atom that lost an electron (called oxidation), is now positively charged and
called a cation. Think “Cats have a positive happy effect on people.”
 It is called oxidation because an atom usually loses the electron to an
Oxygen atom, just like at the end of the Electron Transport Chain in
Cellular Respiration.
c. The atom that gained an electron (called reduction), is now negatively charged
and is called an anion. Think “Angry people are negative.”
 Remember the Law of Conservation of Matter? Matter is neither
created, nor destroyed; just transferred or transformed. Well the
electrons have been transferred from one atom to another atom.
b. These molecules typically have electronegativity differences greater than 1.7.
c. These bonds are usually weaker than covalent bonds, but can be stronger sometimes.
3. Metallic Bonds
a. These are bonds involving electrically charged metals. One positive and one negative.
(These are like Ionic, but there is no electron swap resulting in ions.)
b. Metals tend to have very few valence shell electrons. Look at the Periodic Table to see.
c. These electrons tend to be highly mobile and therefore can travel from one atom to another
adjacent atom. This moving field of electrons is often referred to as “a sea of electrons” and
is the reason why metals are great for generating electricity.
i.
The electrons are said to be delocalized (as the can roam from atom to atom).
d. Metals tend to have luster (Shine).
i.
Remember the Photoelectric effect and electrons coming of the surface and emitting
light (energy loss).
e. Metals are also malleable. Think mallet (hammer).
i.
This means they can be hammered into thin sheets. (Such as aluminum foil.)
f. Metals are also ductile.
i.
This means they can be “pulled” into long thin wires.
g. The strength of metallic bonds varies with nuclear charge and the number of valence
electrons present in the atoms of those metals.
h. Enthalpy (the total amount of energy in a system) of Vaporization is used to calculate the
relative strength.
i.
This is the total amount of energy require to turn a liquid metal into a gas (vapor) at
a constant pressure.
ii.
More valence electrons in the “sea” the higher the amount of energy.
Unit 3: Bonding and Nomenclature Content Outline: Intermolecular Forces (3.2)
I.
Intermolecular Forces (Between molecules; like an “Interstate” goes between states.)
A. These are forces or attractions between molecules of solids or liquids mainly, but can be gases too.
B. They vary in strength; but are generally weaker than chemical bonds, as there is no real physical
interaction like there are in bonds…hence the term forces.
C. These forces usually affect the boiling points of chemicals.
1. Higher boiling points = more interactions/attractions between molecules.
2. Lower boiling points = less interactions/attractions between molecules.
D. There are 4 basic types of Intermolecular Forces (from strongest  weakest):
1. Network Solid
A. This is a solid molecule that is a 3D -network of covalently bonded atoms
For example, a diamond and graphite made of Carbon or elemental silicon such as sand and
quartz.
Diamond is the hardest substance, best heat conductor; but very expensive.
Graphite and silicon can be lubricants.
B. They are tough and have very high melting points and boiling points as they have numerous
intermolecular bonds.
2. Hydrogen Bonds
a. These are forces of attraction that usually are associated with a polar molecule with a positively
charged Hydrogen molecule and a negatively charged second molecule, in close proximity,
containing Oxygen, Fluorine, or Nitrogen.
b. These are relatively weak attractions, but probably the most important in terms of life.
i. They hold the two halves of a DNA molecule together.
ii. They give water the ability to move other charged molecules (food) through our blood
vessels.
iii. They allow trees to move water in massive quantities from the roots to the leaves.
c. They usually affect the boiling point of substances, including water.
i. The more Hydrogen bonds present within a liquid, the higher the boiling point or Specific
heat.
α. This is the amount of heat energy needed to raise 1 gram of a substance by 1OC.
b. Water has a specific heat of 4.18 J… so it takes a whole lot of energy to break all for
Hydrogen bonds, at one time, and turn it into a gas.
c. Iron has a specific heat of 0.449J… so it heats up very quickly in that sea of electrons.
This is why we use it to cook food on a stove or fire. Remember, we also use
metals for conducting electricity.
Which would you rather touch after 1 hour in the summer sun, a glass of water or an
iron pan?
ii. The fewer Hydrogen bonds, the lower the boiling point usually.
iii. A single water molecule has 4 simultaneous Hydrogen bonds, so it requires a high
temperature to boil, which is good since water covers 2/3rds of Earth and makes up
90% of your body.
d. Hydrogen bonds are represented by dotted lines, for example H…O.
e. As Hydrogen bonds are fairly weak attractions, they are easy affected/broken by changes in
temperature (high heat breaks them), changes in pH concentrations, or changes in ionic
solutions, such as saline (salt water).
3. Dipole-Dipole Forces (Polar molecules)
a. These attractions occur between two opposite charged parts of two different molecules.
b. “Dipole” means “two poles” on one single molecule. Like the Earth has two poles – north and
south poles.
c. These forces are represented by a single crossed arrow, such as +-->.
i. The crossed part (+) represents the positive part of the first molecule.
ii. The arrow part --> represents the attraction to the negative part of the second molecule.
d. For linear molecules, orientation of molecules is not an issue; just attraction by opposite
charges.
e. For lattice or 3D folded chains, the molecule orientation is important. Just like in the folding of
proteins – string of folded chain of amino acids.
f. Induced Dipole Moment
i.
This is a temporary moment where a non-polar molecule behaves “like” a polar molecule
temporarily because of electron attraction with another molecule. (“Induce” means “to
make into”)
ii. The most important example of this temporary attraction is Oxygen gas (O2) dissolved in
water.
α. Oxygen gas is non-polar, but when it interacts with water, it can become induced to
act polar and be relatively attracted to the highly positive portion of water (the
Hydrogen atoms).
 Cold water holds onto the Oxygen gas tighter because the molecules are
moving slower.
 Warmer water holds onto the Oxygen gas weaker because the molecules are
moving faster.
 Regardless, the dissolved Oxygen gas in water allows for life forms to live in
the water. They mainly use gills or diffusion to get the Oxygen gas out of the
water.
4. London Dispersion Forces (Non-polar molecules) (Please help students “see” the term.)
a. These were discovered by Fritz London in 1930.
b. They are also known as Van der Waals Interactions associated with enzymes.
c. They are the temporary interactions between molecules due to temporary
“clumping/dispersion” of electrons on one side of an atoms nucleus. This temporary “clumping”
creates a temporary polar “like” molecule. Now it can quickly and briefly interact with another
charged particle/molecule.
Remember, enzymes “speed up” a chemical reaction by quickly and briefly grabbing molecules
and either puts them together with a chemical bond or grabs and breaks the bond. The
temporary polar “like” quality allows for the “grabbing”.
d. Any atom or molecule can have these forces, because every atom has moving electrons.
i. These are the only type of interaction that Noble gases can have with each other or other
atoms… brief and temporary.
e. These can also affect boiling points of substances.
i. High amount of interactions = higher boiling points.
ii. Fewer interactions = lower boiling points.
iii. The number of interactions (and boiling points) increases as atomic or molecular masses
(AMUs) increase (molecule gets larger).
iv. The larger the molecule becomes the stronger the dispersion force due to increases in
polarizability.
f.
Polarizability
i. These are the distortions to the electrical cloud because of the electrical charge associated
with the clumping.
Content Outline: Molecules (3.3)
I.
Molecule
A. The term refers to any electrically neutral group of atoms that are bound together.
1. It can be 2 or more atoms of the same element or composed of different elements.
For example: N2 (nitrogen gas) or NH3 (Ammonia)
A. Molecular Compound This term is used to refer to a chemical compound (2 or more atoms) whose
simplest units are molecules.
For example: H20 (water) or O2 (Oxygen gas) or C6H12O6 (Glucose)
B. This term is usually used with molecules that are bound together using covalent bonds.
C. These molecules can possess single bonds (-), double bonds (=), or even triple bonds (Ξ).
1. The purpose of “creating” the bonds is to achieve the lowest possible Potential Energy state by
filling the valence shell with 8 electrons (except for H, He, Li, and Be) so that a molecule behaves
“like” a Noble gas.
a. 8 electrons in the outer valence shell is referred to as the Octet Rule.
b. It is a combination of the “s” orbital (2 electrons) and the “p” orbital (6 electrons).
2. These “shared electrons” are from overlapping orbitals in the valence shells.
II. Ionic Compound
A. These molecules are composed of positive and negative ions that are combined in a lattice (3D cube) like
structure that looks “like” a crystal (crystalline).
1. The ions alternate (positive- negative) so as to maintain neutrality and reduce repulsive forces
between like charged ions.
2. The attractive strengths vary between the atoms is due to the size and charges associated with the
ions in the compound.
a. These ionic compound strengths are measured using Lattice Energies
b. This is the amount of energy released when one mole of an ionic crystalline compound is formed
from gaseous ions.
B. The electrical charges are balanced, so as to be neutral.
C. Ionic compounds are hard but brittle (crumbly).
D. In the solid state, they cannot conduct electricity (as the atoms can’t move); but when dissolved in
water, they are great conductors of electricity (the ions cans move freely).
III. Chemical Formula
A. This shows the relative number of atoms of each element in a chemical compound using symbols and
subscripts.
For example: C12H22O11 (Sucrose)
B. Common prefixes for numbers:
Di – 2
Penta – 5
Octa – 8
Mono-1
Tri – 3
Hexa – 6
Nona – 9
Tetra – 4
Hepta – 7
Deca – 10
IV. Molecular Formula
A. This type of chemical formula shows the types (Chemical symbol) and numbers (subscripts) of atoms
combined in a single molecule of a molecular compound.
B. For example: H2O or O3
V. Formula Unit
A. This type of chemical formula is used with ionic compounds to establish the simplest collection of atoms
from which an ionic compound. It is like Molecular weight. Calculated using Atomic Masses and
measured in moles.
For example: NaCl or KCl or CaF2 (There could be thousands of atoms, but this is the simplest ratio.)
VI. Empirical Formula
A. This type of chemical formula shows the kinds of atoms and their relative numbers in a substance in the
smallest whole number ratios.
For example: H2O2  HO OR C6H12O6  CH2O (Take the lowest subscript number and divide it into
the other subscripts.)
VII. Energy and Molecules
A. The natural tendency is to achieve the lowest possible Potential Energy state and thus behave “like” a
Noble gas element.
B. Energy is released in bond formation between atoms.
C. Energy is required in the breaking of a bond between atoms.
1. The energy to make or break a bond is referred to as bond energy.
2. It is reported as kiloJoules/mole (kJ/mol)
3. The octet rule applied to the number of bonds an element can make.
D. Endothermic – Energy is absorbed from the surrounding environment.
E. Exothermic – Energy is released to the surrounding environment.
F. The energy of a chemical bond is inversely proportional with the nucleus distance.
1. This is referred to as Bond Length Energies.
1. High Potential energy = very little distance apart. (There exists lots of repulsion between like charges
of the protons in each nucleus.)
2. Low Potential Energy = very far apart. (There is very little repulsion by the nucleus protons.)
3. Do not confuse with Coulombs Forces that deal with one atom’s own electrons and protons. These
are between two different atoms.
Content Outline: Molecular Geometry and VSPER Theory (3.5)
I.
Structure = Function
A. The 3 dimensional shape of atoms in molecules helps determine the chemical properties of that
molecule.
1. Properties such as:
a. Polarity and charge
b. Attraction or Repulsion with other molecules (Intermolecular forces).
c. Re-activity or inactivity with other molecules.
II. VSEPR Theory
A. This stands for: Valence-Shell, Electron-Pair Repulsion.
B. The theory states “Repulsion between the sets of valence-level electrons surrounding an atom
causes these sets to be oriented as far apart as possible.”
1. “Shared pairs” of electrons are as far apart as possible.
2. “Lone pairs” (unshared) electrons occupy space around the central atom.
a. Some times they are shown as “. .” but they may not be all the time, it is just understood that
they are present in the molecule.
b. They have greater repulsion strengths than “paired” electrons.
3. Sigma (σ) bonds (Head to head overlap)
a. These are the primary (first) bonds formed by the overlapped orbitals of valence shells.
i. Because of this overlap, the electron density in that region increases.
4. Pi (π) bonds (Side by side overlap)
a. These are the second (double) or third (triple) bonds, in a multiple bond set.
i.
These are from over lapping “p” orbitals
C. There are 8 geometric shapes of molecules:
1. Linear (Line in one plane)
a. Number of atoms bonded to the central atom = 2.
b. The molecule has a bond angle of 180O.
c. It has a basic molecular formula of: AB2 (“A” is one element; “B” is the other element)
For example: Be F2
2. Trigonal-Planer (3 point pyramid in one plane “flat”)
a. Number of atoms bonded to the central atom = 3.
b. The molecule has bond angles of 120O.
c. It has a basic molecular formula of: AB3.
For example: BF3
3. Bent or angular planar (Boomerang shape in one plane “flat”)
a. Number of atoms bonded to the central atom = 2.
b. The molecule has bond angles of greater than 120O.
c. It has a basic molecular formula of: AB2E (E = number of Lone pair electrons…on 1 point)
For example: ONF
4. Bent or angular pyramidal (Boomerang shape of atoms but 4 point 3 side pyramid whole)
a. Number of atoms bonded to the central atom = 2.
b. The molecule has bond angles of 104.5O.
c. It has a basic molecular formula of: AB2E2. (Lone pair on top of pyramid and 1 base point)
For example: H2O
5. Tetrahedral (3 sided 4 point pyramid with flat bottom)
a. Number of atoms bonded to the central atom = 4.
b. The molecule has bond angles of 109.5O.
c. It has a basic molecular formula of: AB4.
For example: CH4
6. Trigonal-Pyramidal (3 sided, 4 point pyramid with flat bottom)
a. Number of atoms bonded to the central atom = 3.
b. The molecule has bond angles of 107O.
c. It has a basic molecular formula of: AB3E. (Lone pair on top of pyramid)
For example: NH3
7. Trigonal-Bipyramidal (2- 3 sided, 4 point pyramids with flat bases put base to base)
a. Number of atoms bonded to the central atom = 5.
b. The molecule has internal vertical bond angles of 90O and horizontal (base) angles of 120O.
c. It has a basic molecular formula of: AB5.
For example: PCl5
8.
Octahedral (2 – 4 sided, 5 point pyramids with flat bases put base to base)
a. Number of atoms bonded to the central atom = 6.
b. The molecule has bond vertical and horizontal internal angles of 90O.
c. It has a basic molecular formula of: AB6.
For example: SF6
III. Hybridization of electron orbitals during bonding
A. “Hybrid” means “two combined as one”.
B. These are electron configuration models that are helpful to explain the rearranging of electron
orbitals when covalent bonds occur.
C. It is the mixing of two or more atomic orbitals of similar energies on the same atom to produce new
hybrid atomic orbitals of equal energies.
1. Carbon is a great example with its tetravalence (4 valence shell electrons).
a. Step 1: 1 of the “s” electrons elevates out to the open “p” suborbital.
b. Step 2: Now all four orbitals “mix” so that each (the 1 “s” and 3 “p”) possesses equal
energies.
c. Step 3: Each hybridized orbital now enters into a covalent bond with another atom.
For example: CH4 (Methane Gas) sp3
D. Hybridized orbitals are represented as: sp, sp2, sp3
1. The superscript number on the “p” + 1 equals the number of hybrid orbitals.
For example: sp = 2; sp2 = 3; sp3 (above) = 4.
Content Outline: Writing and Naming Ionic Compounds (3.6)
I.
Chemical Formula
A. This type of formula indicates the relative number of atoms of each element present in a chemical
compound.
B. These are used when naming and writing molecular or ionic compounds.
II. Ionic Compounds
A. These molecules consist of a lattice network of positive and negative ions held together by mutual
attraction between ions.
B. The chemical formula for an ionic compound represents one formula unit.
1. Formula Unit, remember is the simplest ratio of the compounds positive ions (cations) and
negative ions (anions)
For example, you may have a block of salt with millions of ions but the formula unit is NaCl.
a. Polyatomic ions, remember are groups of atoms that collectively can behave “like” cations or
anions.
For example, Al2(SO4)3 Al+3 is the cation; (SO4)2- is the anion.
C. Remember, in Binary (2) ionic compounds the charges must balance to be neutral.
III. Naming Monatomic, Binary, and Polyatomic Ions
A. Naming (nomenclature) Monatomic Ions:
1. These are single charged (+ or -) atoms.
a. They may have Noble Gas electron configurations.
2. Positive ions:
a. They are named using the elements normal name.
b. They can possess charges of +1, +2, +3, or even +4.
Looking at the element’s column on the Periodic Table can be helpful for this.
3. Negative ions:
a. Step 1: Drop the elements normal ending on the name.
Step 2: Replace the ending with the ending of “ide”.
For example, Oxygen becomes Oxide; Fluorine becomes Fluoride
4. Ion names with Roman Numerals: (Mainly used with “d Block elements”)
a. These Roman Numerals are part of the Stock System of naming ions and elements.
b. These elements can form 2 or more cations only, each possessing a different charge.
i.
There are no elements can form more than 1 anion.
c. The Roman Numeral indicates the charge of the ion.
For example, Fe(II)  2+, Fe (III)  3+ OR Pb(II)  2+, Pb (IV)  4+
B. Writing Binary Ionic Compounds
1. The total compound charges must equal out to be neutral.
2. Step 1: Cation symbol is written first.
Step 2: Write the element’s ionic charge (1+, 2+) above the name.
Step 3: Anion symbol is written next to the positive ion.
 Single or Polyatomic
Step 4: Write the element’s ionic charge (-1,-2) above the name.
Step 5: Balance the charges, using the subscripts, so as to equal out to neutral.
 The “crossover” method can be used for this step.
For example, Al2O3
Aluminum has a charge of +3 so Al+3
Oxygen has a charge of -2 so O-2
Now “crossover” by cross multiplication; Al (+3 x 2 = +6); O (-2 x 3 = -6)
So the formula is written as Al2O3
C. Naming Binary Ionic Compounds
1. Step 1: Use your chemical formula.
Step 2: The cation name is written first. (Include Roman numerals if shown.)
Step 3: The anion name goes second.
For example, Al2O3  Aluminum Oxide; NaF  Sodium Fluoride.
a. The ratio of the ions is not indicated in the name because it is understood based upon the
relative charges of the compound’s ions to balance out to be neutral.
D. Naming Polyatomic Ionic Compounds
1. Most polyatomic ions with a negative charge contain Oxygen and are referred to as Oxyanions.
(“oxy” refers to “Oxygen” and Anion for negation ions.)
a. The number of Oxygen atoms affects the name of the polyatomic ion.
i.
The oxygen is not written in the name though.
b. Oxyanions with two forms:
i. Ion with the greater number of Oxygens ends with “ate”. Like Nitrate  NO3
ii. Ion with the lesser number of Oxygens ends with “ite”. Like Nitrite  NO2
c. Oxyanions with more than 2 forms:
i. Same as above applies, but the following 2 also.
ii. Ion with the least number of Oxygen begins with the prefix “Hypo” (means “least”).
Like Hypochlorite.  ClOThen Chlorite. (See above) ClO2Then Chlorate. (See above)  ClO3iii. Ion with the most number of Oxygens begins with the prefix “Per” (means “most”).
Like Perchlorate  ClO4iv. The most common for this are Chlorine, Phosphorus, and Manganese
d. The name is associated with the first element if there is no Hydrogen also present.
Like Bromate  BrO3- Exceptions: Cyanide  CN-, Acetate  CH3COOi. If Hydrogen is present, it is written first, then the next element is given the primary name.
Like Hydrogen Carbonate  HCO3- Exception: Hydroxide  OHIV. Complex Ions (Hydrated Cations)
A. They are transitional (d-block) metal ions that have coordinate-covalent bonds with water
molecules.
1. Coordinate-covalent bonds
a. These are interatomic attractions resulting from the sharing of a lone pair of electrons from
one atom with another atom.
For example, Fe2+ + 6H2O  [Fe(H2O)6]2+
b. These ions are associated with acids.
c. These solutions also tend to be very colorful.
Content Outline: Writing and Naming Molecular Compounds and Acids (3.7)
I.
Molecular Compound
A. This is a compound composed of individually covalently bonded atoms, units, or molecules.
B. These can be named (nomenclature) in one of two different ways:
1. Stock Method
a. This system uses the elements or polyatomic molecules oxidation numbers/state.
2. Prefix system
a. This system uses prefixes to indicate the number of atoms, of each kind, within a molecule.
b. Prefixes: 1 –Mono
4 – Tetra
7 – Hepta
10 - Deca
2 – Di
5 – Penta
8 – Octa
3 – Tri
6 – Hexa
9 – Nona
II. Stock System (Oxidation number)
A. Oxidation number/state
1. This term refers to the number of electrons that must be added to or removed from an atom in a
combined state to convert the atom into the elemental form.
a. Indicated the distribution of electrons among bonded atoms.
B. Assigning Oxidation numbers to atoms or polyatomic groups.
1. As a rule, shared electrons go to the more electronegative atom.
2. Additional Rules:
a. Atoms in a pure element have oxidation numbers = 0.
b. The more electronegative element in a binary (2) compound is assigned it’s anion charge; the
least electronegative element gets assigned it’s cation charge.
c. Fluorine, the most electronegative element, is always assigned an oxidation number = -1.
Remember, your Periodic Trends for electronegativity on the Periodic Table? As you go
across a period they increase; as you go down a column/group they decrease.
d. Oxygen is always assigned -2 unless with Fluorine. If Fluorine is present, the Oxygen is
assigned +2.
You must always check for Fluorine, if you see Oxygen.
e. Hydrogen is always assigned +1 unless it is paired with a metal. If it is with a metal, then it is
assigned -1.
Know your metals by being able to look at a Periodic Table quickly. Remember, there are
metals, metalloids, and non-metals.
f. Neutral compounds equal zero.
g. Polyatomic ions are assigned oxidation numbers equal to the charge of the group.
Know those common polyatomic ions to make your life in Chemistry class easier.
h. This system can also be used with ions.
C. Naming compounds (molecular or ionic) using oxidation numbers:
1. This method of nomenclature will use Roman numerals for cations
2. Step 1: Use your chemical Formula.
Step 2: Write the cation name. First letter is capitalized
Step 3: Look at the subscript of the anion.
Step 4: Look at the subscript of the cation.
Step 5: Make the charges of each equal out to zero net charge (neutral). It may involve a little
math… adding or dividing.
Step 6: Write the numeral used to balance the cation charge with the anion charge as a Roman
numeral after the cation name. Such as Iron (II) or Lead (IV)
Step 7: Write the anion name after the Roman numeral. First letter is lower case.
Remember, to drop the ending and add the ending “ide”.
For example, PCl3  Cl = -1 so P must be +3 to equal zero; therefore, it is
Phosphorus (III) chloride.
N2O  O = -2 (no Fluorine present) so N2 (each must be 1); therefore, it is
Nitrogen (I) oxide.
Mo2O3  O = -2 (no Fluorine present) so -2 x 3 = -6; 2 Mo so 6/2 = 3 therefore
Molybdenum (III) oxide
III. Prefix System (numbers)
A. This system of naming uses the prefix to identify the number of atoms present.
B. Step 1: The Element with the smaller group (column) goes first.
 If both elements come from the same group, then the highest energy level (row) goes
first.
Step 2: If only 1 atom is present, just use the element’s name.
If more than 1 atom is present, you must add the prefix to the elements name.
For example, 1 Hydrogen present  Hydrogen Chloride (HCl)
3 Hydrogens present  Trihydrogen Antimonide (H3Sb)
2 Sodium & 2 Oxygen  Disodium Dioxide (Na2O2)
Step 3: Drop the ending on the anion and replace with “ide”.
 DO NOT confuse the steps with your polyatomic ion rules.
C. If a prefix ends with an “o” or an “a” and the word following begins with a vowel, drop the “o” or “a”
from the prefix.
D. If both elements in the compound are non-metals the first named element is decide based upon this
sequence: (First)C P N  H  S I Br  Cl O  F (last possible)
This is an easy way to remember the sequence:
C.P. Newton & H.S. Igor brought Cauliflower to the party.
IV. Naming Acids
A. Acid
1. This is a substance that generates Hydrogen Ions (H+) when placed in a solution.
a. The Chemical Formula usually begins with H. Exception Acetic Acid  CH3OOH
b. There are basically two types of acids:
i. Binary Acids (2 elements present)
α. One element is Hydrogen.
b. The second element is a Halogen(Group 17).
ii. Oxyacids (3 elements present)
α. One element is Hydrogen.
b. The second element is Oxygen.
c. The third element is a non-metal or polyatomic ion.
B. Naming:
1. Step 1: Do you have a monatomic anion or a polyatomic anion.
Step 2: If monatomic, - write “hydro” in front of the anion name.
For example, HCl Hydrochloric
If polyatomic, - use the polyatomic ion naming rules for the anion.
For example, H3PO4  Phosphoric
* The Hydrogen is not mentioned.
Step 3: If monatomic, replace “ide” with “ic” and add the word “acid”.
If polyatomic, replace “ite” with “ous” OR “ate” with “ic”
* Don’t forget your prefix, if needed… Hypo or Per
For example, HClO  Hypochlorous Acid
HClO2  Chlorous Acid
HClO3  Chloric Acid
HClO4  Perchloric Acid
C. Salt
1. This is an ionic compound composed of a cation and an anion from an acid.
Content Outline: Naming Basic Organic Molecules (3.8)
I.
Organic Chemistry
A. This is a branch of chemistry that studies Carbon containing compounds.
1. Not all Carbon containing compounds are considered organic
For example, CO, CO2, and Na2CO3
B. Carbon to Carbon bonds
1. This type of bonding is referred to as Catenation.
2. Carbon molecules can form chains or rings.
3. Types of Carbon Bonds
a. Single bonds (-)
i.
These molecules with all single bonds are called Alkanes. (end with “ane”)
ii.
If circular, they are called cycloalkanes.
b. Double bonds (=)
i.
These molecules are called Alkenes. (end with “ene”)
c. Triple bonds (Ξ)
i.
These molecules are called Alkynes. (end with “yne”)
4. Carbon loves Hydrogen. The two together are referred to as Hydrocarbons.
a. These are energy sources, such as oil and gasoline and also oils (plants) and fats (animals).
b. Types of Hydrocarbons
i.
Saturated Hydrocarbons
α. These molecules have all single bonds in the hydrocarbon area.
* These chains tend to be straight.
b. These are solids at room temperature.
For example, animal fat
ii.
Unsaturated Hydrocarbons
α. These molecules can have a double or triple bond in the hydrocarbon area.
* These cause the chains to “kink” or bend” at the double or triple bond.
b. These are liquids at room temperature.
For example, plant oils such as vegetable oil or corn oil.
iii.
Polyunsaturated Hydrocarbons
α. These molecules can have multiple double or triple bonds in the hydrocarbon area.
* These have many “kinks” or “bends”.
b. These are also liquids at room temperature.
II. Naming Basic Hydrocarbons
A. Uses a Prefix naming system
1 – Meth
4 – But
7 – Hept
10 - Dec
2 – Eth
5 – Pent
8 - Oct
3 –Prop
6 – Hex
9 – Non
B. Step 1: Count the carbons present in the chain or circle.
Step 2: Use the number of carbons to get the appropriate prefix.
Step 3: Look at the number of Hydrogens in the compound.
 “ane” – 2n + 2 (n = number of Carbons)
 “ene” – 2n
 “yne” – 2n-2
Step 4: Write the appropriate ending after the prefix.
Step 5: If circular, add “cyclo” in front of the prefix.
Content Outline: Calculating with Empirical Formulas (3.10)
I.
Empirical Formula
A. This is a type of formula that shows the element symbols combined in a compound, with subscripts
showing the smallest whole-number mole ratio (not the actual number present) of the different
elements.
1. For ionic compounds they are usually the same.
For example, C6H12O6 has an empirical formula of CH2O… hence the term Carbohydrate.
NaCl (ionic) has an empirical formula of NaCl.
II. Calculating Empirical Formula:
A. Step 1: Assume you have a 100.0 gram sample of your compound.
Step 2: Convert your percentages to gram measurements.
For example, 32.38 % Na  32.38 grams of Na
22.65% S  22.65 grams of S
44.99% O  44.99 grams of O
Step 3: Take the gram amount and convert to moles by using the conversion moles over grams
32.38 g of Na x 1mole of Na = 1.408 mol Na
22.99g Na
22.65 g of S x 1 mole of S = 0.7063 mol of S
32.07 g of S
44.99 g of O x 1 mole of O = 2.812 mol of O
16.0 g of O
Step 4: Divide the mole value for each element by the smallest mole value to convert to a ratio.
Na = 1.408/ 0.7063 = 1.993
S = 0.7063/0.7063 = 1.0
O = 2.812/0.7063 = 3.981
Step 5: Round each ratio value up or down to the nearest whole number.
Na = 1.993 rounds to 2; S = 1; O = 3.981 rounds to 4
Step 6: Combine the ratio values into an empirical number (values are subscripts).
Na2SO4
III. Calculating Molecular Formulas from Empirical formulas
A. Remember, the empirical formula is the lowest possible ratio of elements. The molecular formula is
the actual formula for a molecular compound.
1. An empirical formula may or may not be a correct molecular formula.
B. The relationship between an empirical formula and a molecular formula can be expressed as:
X (empirical formula) = molecular formula
X represents a whole number multiplier. (The number needs to convert empirical to actual).
For example C6H12O6 is expressed as 6(CH2O)
IV. Calculating Molecular Formula Mass from Empirical Formula Mass
A. The relationship can be expressed as:
X(Empirical Formula Mass) = Molecular Formula Mass
X represents a whole number multiplier
B. Step 1: You must be given a known formula mass of a compound.
Step 2: Divide the given formula mass by the Empirical Formula Mass.
 This gives you the multiplier value.
Step 3: Multiply the empirical formula using the multiplier value to get the molecular Formula
Step 4: Use the Molecular Formula to calculate the Molecular Formula Mass