Unit 4: Inorganic NomenclaturePeriod EMr. DesjardinsDue February 10, 2012Editor: Brendan Morrissey
Inorganic nomenclature is the system of naming inorganic chemical compounds. The systematic name of each compound provides the structure and composition of the molecule. This naming system is better than using the simple common ame of a compound because it gives a clear and precise description of the compound. Inorganic compounds are compounds that do not deal with formation of carbohydrates, meaning they do not include carbon and hydrogen structures. In this chapter, you will learn about the different froms of inorganic nomenclature, such as iionic compounds, molecular compounds, acids, and bases. You will also learn how these different compounds are bonded together. This video will explain further how inorganic nomenclature aids chemists.



Editor: Brendan Morrissey


Pages 187-193: Ions

Group Members: Austin Burlone and Siri Devlin

Austin Burlone (Pages 187-190):
Valence Electrons
- All elements of each group of the periodic table behave similarly since they all have the exact same number of valence electrons
  • Valence electrons are electrons that are found in the highest occupied energy level of each atom of an element
  • These electrons determine the chemical property of that element
  • Only electrons used in chemical bonds
- The group number represents the number of valence electrons in the atoms of the elements that are in that group
- Electron dot structures are diagrams that have dots that represent the valence electrons of an element's atom
valence_electrons.gif
The Octet Rule
- Gilbert Lewis developed the octet rule in 1916
  • This rule says that in forming compounds, atoms achieve the electron configuration of noble gases
- Atoms in metallic elements lose valence electrons which results in a complete octet in the next lowest energy level
- Atoms of nonmetallic elements have to share electrons with other nonmetallic elements to reach the complete octet
Formation of Cations
- Atom is neutral because of equal number of protons and electrons
- To form a cation, an atom loses electrons
  • With metallic elements, the name of the cation and the name of the element are the same
  • Yet there is still differences between the two
- Most common cations are formed from metal atoms losing valence electrons
- Cations of elements that are in Group 1A have a charge of +1
- Cations of elements that are in Group 2A have a charge of +2
- Yet for transition metals, the charges of the cations are not always the same
- Some ions of transition metals are exceptions to the octet rule
- Rarely are ions with charges of three or greater found



Siri Devlin (Pages 191-193)

Formation of Anions

-An anion is an atom or a group of atoms with a negative charge

-The gain of negatively charged electrons by a neutral atom produces an anion

  • The name of the anion typically ends in –ide

-Halide ions are the ions that are produced when atoms of chlorine and other halogens gain electrons

  • All halogen atoms have seven valence electrons and need to gain only one electron to achieve the electron configuration of a noble gas.- Thus all halide ions have a charge of -1

Some common anions

1-
2-
3-
F- fluoride
Cl- chloride
Br- bromide
I- iodide
OH- hydroxide
ClO- hypochlorite
NO3- nitrate
C2H3O2- acetate
HCO3- hydrogen carbonate
O2- oxide
S2- sulfide
SO4^2- sulfate
CO3^2- carbonate
N3- nitride
P3- phosphide
PO4^3- phosphate

Pages 194-200: Ionic Bonds and Ionic Compounds

Group Members: Delia Calderon

Delia Calderon (Pages 194-200):

Formation of Ionic Compounds

  • Compounds composed of cations and anions are called ionic compounds, they are usually composed of metal cations and nonmetal anions.
  • Although they are composed of ions, ionic compounds are electrically neutral, the total positive charge of cations equal the total negative charge of anions
  • Ionic Bonds
    • The electrostatic forces that hold ions together in ionic compounds are called ionic bonds.
    • Sodium has a single valence electron that it can easily lose. Chlorine has seven valence electrons and can easily gain one.
    • Sodium and Chlorine atoms combine in one-to-one ratio and both ions have stable octets.
  • Formula Units
    • Chemists represent the composition of a substance by writing chemical formulas, which show the kinds and numbers of atoms in the smallest representative unit of a substance.
    • A formula unit is the lowest whole number ratio of ions in an ionic compound.

Properties of Ionic Compounds

  • Most ionic compounds are crystalline solids at room temperature. The component ions are arranged in repeating Ionic_compound.jpg
three-dimensional patterns
  • Each ion is attracted strongly to each of its neighbors and repulsions are minimized, the large attractive forces result in
a very stable structure.
  • Ionic compounds generally have high melting points
  • The coordination number of an ion is the number of ions of opposite charge that surround the ion in a crystal.
  • Another characteristic of ionic compounds is that they can conduct an electric current when melted or dissolved in water.



Pages 201-205: Bonding in Metals

Group Members: Brett Chatfield

Brett Chatfield (Pages 201-205):

Metallic Bonds and Metallic Properties

- Metals are made of closely packed cations rather than neutral atoms
- Valence electrons of metal atoms can be modeled as a sea of electrons
- This means that valence electrons are free to move where they wish
- Metallic Bonds are the attraction of free floating valence electrons for the positively charged metal ions
For example, the sea-of-electrons model can explain many of the physical properties of metals
- Both ductility and malleability of metals can be explained in terms of the mobility of valence electrons
- The a metal is pressured, the metal cations easily slide past one another like two balls dipped in oil

Crystalline Structure of Metals

- Similar arrangements can be seen between crystalline structures of metals and also apples stacked at a grocery store
- As a matter of fact, metals that contain just one atom are actually arranged in a very simple form of crystalline structure
- Metal atoms are arranged in very compact and orderly patterns
- In body centered cubic structures, every atom except those on the surface, have eight neighbors.
- Sodium, potassium, iron, chromium, and tungsten all crystallize in a body centered cubic pattern
- In a hexagonal close packed arrangement, every atom has twelve neighbors

http://www.youtube.com/watch?v=32_JVR6zFcg


external image Au-bs.jpg
external image Au-bs.jpg
external image gold-bars4.jpg
external image gold-bars4.jpg
- The crystalline structure of gold compared to physical shape

Alloys

- Everyday we use metallic items, like spoons and forks, but many of these items are not really pure metals
- These items are really mixes called alloys
- Alloys are mixtures composed of two or more elements, at least one of which is a metal
- Brass, for example, is an alloy of copper and zinc
- Alloys are important because their properties are often superior to those of their component elements
- Sterling silver is actually part silver and part copper
- He most important alloys today are steels
- The principal elements in most steel are elements such as iron and carbon
- Alloys can form from their component atoms in different ways
- The components are around the same size and can replace each other in the crystal
- There are various types of steel, for example, carbon atoms occupy the spaces between the iron atoms which are why steels are interstitial alloys


external image chrysnit.gif
external image chrysnit.gif


- Here the Chrysler Building, in New York City, is made of a steel frame and has a spire that is sheathed in shiny stainless steel
decrease size
decrease size



Pages 213-216: Molecular Compounds

Group Members: Marissa Chura

Marissa Chura (Pages 213-216):
Molecules and Molecular Compounds
-Two atoms joined by sharing electrons are held together by a covalent bond.
-A group of atoms joined by covalent bonds is a molecule.
-A diatomic molecule is made up of 2 atoms.
-Molecules combine to form compounds called molecular compounds.
-All of the molecules of a molecular compound are the same.
-Molecular compounds tend to have lower boiling and melting points than ionic compounds.
-Most are composed of two or more nonmetals.

Molecular Formulas
-Molecular formula=chemical formula of a molecular compound.
-It shows the amount of atoms of each element contained in a molecule.
-For example:
*Water- composed of two hydrogens, one oxygen- molecular formula: H2O
*Ethane-composed of two carbons, six hydrogens- molecular formula: C2H6
-Not just used for compounds.
-Oxygen molecule- composed of two oxygen atoms- molecular formula: O2.
-Does NOT tell the structure/arrangement of the molecule.

Pages 217-224: The Nature of Covalent Bonding

Group Members: Nicholas Cunha and Benjamin Ross

Nicholas Cunha (Pages 217-223):

8.2 The Nature of Covalent Bonding

Connecting to Your World

You know that without oxygen, we all wouldn’t be living. However, oxygen has another crucial role to life. High in the atmosphere, oxygen in another form called ozone forms a layer that filters out the sun’s harmful radiation. In this section you will learn how oxygen atoms can join in pairs to make the oxygen we breath, and how it also joins in groups of 3 to form the ozone high in the atmosphere. In this diagram, different ozone concentrations in the atmosphere are shown.




globeozone.jpg
globeozone.jpg


The Octet Rule in Covalent Bonding

Just like the way electrons tend to be transferred to allow each ion to reach noble gas configuration when ionic compounds are formed, a similar rule applies for covalent bonds. In forming covalent bonds, sharing usually occurs so that atoms attain the electron configurations of noble gases. For example, when a hydrogen atom (that usually has one electron) forms a covalent bond with another hydrogen atom, they then share each of their one electron which makes a hydrogen molecule. Each hydrogen atom therefore reaches the configuration of a noble gas, helium, which has 2 electrons. Combinations of the nonmetallic elements in groups 4A, 5A, 6A, and 7A of the periodic table are likely to form covalent bonds. In such a case, the atoms acquire a total of 8 electrons (octet) by sharing electrons so that the octet rule applies.

Single Covalent Bonds
A single covalent bond is when two atoms are held together by sharing a pair of electrons. The atoms in a hydrogen atom are mainly held together by the attraction of the shared electrons to the positive nuclei. A hydrogen molecule is made up of two atoms that share only one pair of electrons, forming a covalent bond.
hydrogen.gif
hydrogen.gif

An electron dot formula such as H:H represents the shared pair of electrons of the covalent bond by 2 dots. (can also be represented as a dash H-H). A structural formula represents a covalent bond by dashes and shows the arrangement of covalently bonded atoms. In contrast, the hydrogen molecular formulas is H2 indicating only the number of hydrogen atoms in each molecule.
Halogens also form covalent bonds in their diatomic molecules: Fluorine has 7 valence electrons and thus needs one to reach a noble gas figuration. When two Fluorine atoms covalently bond, each achieves the noble gas configuration of neon, each atom giving one electron to complete the octet. * Fluorine atoms share only one valence electron. A pair of valence electrons that are not shared between atoms is an unshared pair.
Electron dot structures for molecules of compounds can be drawn the same way as for diatomic elements. H2O is a molecule containing 3 atoms with 2 covalent bonds. 2 Hydrogen atoms share electrons with 1 Oxygen atom. All atoms attain noble gas configuration through the sharing of electrons. Same applies for ammonia.

bonding_types-water.gif
bonding_types-water.gif


Methane has four covalent bonds with 4 different hydrogen atoms, each contributing one valence shell electron to pair with one of the four valence electrons that the carbon has, attaining noble gas configuration. Although it would not be assumed by its electron configuration, carbon usually makes four bonds. One of the carbon’s 2s electrons is promoted (moved/jumped) to the 2p orbital. This requires only a small amount of energy, providing four electrons of carbon enabling them to form covalent bonds with hydrogen electrons. Because Methane is more stable than CH2, it is more energetically favored even though it needs energy for the movement of electrons that switch orbitals.

methane.gif
methane.gif



Double and Triple Covalent Bonds
Sometimes atoms bond sharing more than one pair of electrons. Atoms share double or triple covalent bonds if they can attain a noble gas structure by sharing 2 or 3 pairs of electrons . Bond with 2 shared pairs- double covalent bond; 3 pairs- triple covalent bond .
Oxygen forms a double, but does not obey the octet rule, an electron dot structure cannot be drawn that adequately describes the bonding in an oxygen molecule.
Nitrogen contains a tripe covalent bond. Each nitrogen atom must gain 3 electrons to attain the noble gas configuration- of neon.
o2.jpg
o2.jpg
nitrogen.jpg
nitrogen.jpg

Diatomic elements table
This table illustrates the properties and uses of the elements that exist as diatomic molecules.
diatomic_chart.jpg
diatomic_chart.jpg

Single, double, and triple covalent bonds can also exist between unlike atoms (Co2). This is a carbon dioxide molecule which contains 2 oxygens each of which forms a double bond with the carbon atom to form a molecule with 2 carbon-oxygen double bonds. The bonds are identical. Carbon dioxide is an example of a triatomic molecule which is a molecule consisting of 3 atoms.




co2.jpg
co2.jpg

Benjamin Ross (Pages 224-229):


- Molecules are held together by tight bonds of several type
- The energy needed to break one of these bonds is called dissociation energy
- Bond dissociation is typically written in the form of energy needed to break one mole of bonds
- A large dissociation energy is common in covalent bonds

- Molecules can be formed from several covalent bonds
- One example is the ozone molecule, composed of three oxygen atoms in a covalent relation
- Double bonds are shorter than single bonds in most cases (the atoms lie closer together)
- A resonance structure that will be present when multiple molecules have the same number of dots in electron dot structure

- It is impossible to perfectly satisfy the octet rule in some cases
- This occurs in atoms whose valence electrons is an odd number
- In this case, one could accurately draw arrows pointing up or down
- Some molecules break the octet rule and have ten or even twelve valence electrons


Pages 230-236: Bonding Theories

Group Members: Kerry Desmond and Phillip Royal

Kerry Desmond (Pages 230-233:

Bonding Theories

I) Molecular Orbitals

-some orbitals exist only for groups of more than one atom
-when two atoms combine, this model assumes that their atomic orbitals overlap to make molecular orbitals.
-molecular orbits belong to a molecule as a whole
-each level of a molecular orbital is filled with two electrons.
Bonding Orbitals-an orbital that can be occupied by two electrons of a covalent bond

Sigma Bonds
-when two atomic orbitals combine to form a molecular orbital that is symetrical around the axis that connects the two atomic nuclei, a sigma bond is formed.
-covalent bonds happen when there is an imbalance between the attractions and repulsions of the nuclei and electrons involved.
-because their charges have opposite signs, the nuclei and electrons attract each other; nuclei repel other nuclei.
-atomic p orbitals can overlap to form molecular orbitals.
Pi Bonds
-in pi bonds, the electrons are likely to b found in sausage-shaped regions below and above the bond axs.
-tend to be weaker than sigma bonds.
VSEPR Theory
-because electrons repel each other, the shapes of molecules adjust so the valence electron pairs stay as far apart as possible.




Phillip Royal (Pages 234-236)

Hybrid Orbitals

- Orbital Hybridization provides information about both molecular bonding and molecular shape
- Hybridization: The process in which several atomic orbitals mix to form the same total number of equivalent hybrid orbitals.
Hybrdization Involving Single Bonds:
- All bonds are identical
- Example: The one 2s orbital and the three 2p orbitals of carbon atoms mix to form four sp^3 orbitals
- These orbitals extend further into space than s or p orbitals, allowing overlap
- Overlap results in unusually strong covalent bonds
carb1.gif

carb2.gif

carbo3.gif

Hybridization Involving Double Bonds
Example:
Ethene

- One Carbon double bond,four hydrogen single bonds
- Each Hybrid orbital is separated by 120 degrees
- Sigma bonds formed with Hydrogen
- Twelve Electrons fill the six bonding molecule orbitals


ethylene.jpg
image1E3.jpg
Hybridization Involving Triple Bonds
- Third type of Covalent bonds, triple bonds
- Found within Ethyne, linear molecule
- A 2s atomic orbital of carbon sometimes mixes with only one of the three 2p orbitals
Ethyne-2D-flat.png
imgres.jpg

Pages 237-244: Polar Bonds and Molecules

Group Members: Kendall Lavin-Parsons and Abbey Salvas

Abbey Salvas (Pages 237-240

Bond Polarity

-covalent bonds involve electron sharing between atoms

-differ in terms of how the bonded atoms share electrons

-character of the bonds in a given molecule depends on kind and number of atoms in bond

-bonding pairs of electrons are pulled between nuclei of atoms sharing the electrons

-when atoms pull equally, bonding electrons are shared equally and the bond is a nonpolar covalent bond

-molecules of hydrogen, oxygen and nitrogen have nonpolar covalent bonds

-diatomic halogen molecules

-polar covalent bond or polar bond: covalent bond in which electrons are shared unequally

-more electronegative atom attracts electrons more strongly and gains negative charge

-less electronegative atom has positive charge

-higher the electronegativity value, the greater the ability of an atom to attract electrons to itself

-electronegativity difference between two atoms tells what kind of bond is likely to form

-as electronegativity difference between two atoms increases the polarity of the bond increases

-if electronegativity difference is greater than 2.0, electrons will be pulled away completely by one of the atoms; ionic bond will form




Polar Molecules

-presence of a polar bond in a molecule makes the entire molecule polar

-one end of the molecule is slightly negative and the other end is slightly positive

-molecule that has two poles is called a dipolar molecule or dipole

-when polar molecules are placed between opposite charged plates, they tend to become oriented with respect to the positive and negative plates

-effect of polar bonds on polarity of molecule depends on shape of molecule and orientation of polar bonds


Polarity_boron_trifluoride.png


Attractions Between Molecules

-intermolecular attractions are weaker than either ionic or covalent bonds

-responsible for determining whether a molecular compound is a gas, liquid, or solid

-Van der Waals Forces

-two weakest attractions between molecules

-Dipole Interactions

-polar molecules attracted to one another

-electrical attraction occurs between oppositely charged regions of polar molecules

-similar to ionic bonds

-Dispersion Bonds

-weakest of all molecular interactions

-caused by motion of electrons

-moving electrons happen to more on side of molecule closest to neighboring molecule, electric force influences neighboring molecule’s electrons to be momentarily more on the opposite side

-causes attraction between two molecules similar to that of permanently polar molecules


Kendall Lavin-Parsons (Pages 241-244):

HYDROGEN BONDS

  • Dipole interactions are what cause the hydrogen bonds

  • Every bond is polar

  • The Oxygen recieves the more negative charge (because of electronegativity)

  • The hydrogen has a slightly positive charge

  • Dipole interactions are only thos strong with Hydrogen containing molecules hense Hydrogen bond

  • HYDROGEN BOND= attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another elctronegative atom

  • Other atom can only be in the same molecule or a close by one

  • Hydrogen is only chemucally reactive element with non underlying electron shielded valence electrons

  • the bonds between hydrogen must be strongly polar

  • includes oygen, nitrogen, and fluorine

  • strongest intermolecule force

INTERMOLECULAR ATTRACTIONS AND MOLECULAR PROPERTIES
  • pysical prperties depend on types of bonding (ionic vs covalent)

  • melting point is much lower with covalent bonds because they are easier to break

  • NETWORK SOLIDS= solids in which all the atoms are covalently bonded to each other EX diamonds and Silicon carbide

  • Should think of network solids as single molecules

  • Ionic compounds also are soluble in water

ex of hydrogen bond in water
ex of hydrogen bond in water

ex of hydrogen bond in water

diamond molecular structure
diamond molecular structure

diamond molecular structure

notice the similarity to Diamond?
notice the similarity to Diamond?

notice the similarity to Diamond?


http://youtu.be/oNBzyM6TcK8 (great video to help tell the difference between covalent and ionic bonds!)

Pages 253-259: Naming Ions

Group Members: Chris Hart and Daniel Lynch

Chris Hart (Pages 253-256):

Monatomic Ions

Ionic compounds consist of a positive metal ion and a negative nonmetal ion combined in a proportion such that their changes add up to a net charge of zero. Some ions, called monatomic ions, consist of a single atom with a positive or negative charge resulting from the loss or gain of one or more valence electrons, respectively.


Cations

Metallic elements tend to lose valence electrons. Lithium, sodium, and potassium in Group 1A lose one electron to form cations. When the metals in Group 1A, 2A, and 3A lose electrons, they form cations with positive charges equal to their group number. The names of the cations of the Group 1A, Group 2A, and Group 3A metals are the same as the name of the metal, followed by the word ion or cation.


Anions

Nonmetals tend to gain electrons to form anions, so the charge of a nonmetallic ion is negative. The charge of any ion of a Group A nonmetal is determined by subtracting 8 from the group number. Anion names start with the stem of the element name and end in -ide. For example, two elements in Group 7A are flourine and chlorine. The anions for these nonmetals are the flouride ion (F-) and chloride ion (Cl-). Anions of nonmetals in Group 6A have a 2- charge. The three nonmetals in Group 5A can form anions with a 3- charge. The majority of the elements in the two remaining representative groups, 4A and 8A, usually do not form ions.


Ions of Transition Metals

The metals of Groups 1A, 2A, and 3A consistently form cations with the charges of +1, +2, and +3, respectively. Many of the transitions metals form more than one cation with different ionic charges. The charges of the cations of many transition metal ions must be determined from the number of electrons lost. There are two methods used to name common ions of many of the metals that form more than one ion. The first is called the Stock System. In the Stock System, a Roman numeral in parentheses is placed after the name of the element to indicate the numerical value of the charge. The name for Fe2+ is read "iron two ion".


An older method for naming these cations uses a root word with different suffixes at the end of the word. The older, or classical, name of the element is used to form the root name for the element. For example, ferrum is Latin for iron, ferr- is the root name for iton. The suffix -ous is used to name the cation with the lower of the two ionic charges. The suffix -ic is used with the higher of the two ionic charges. Using this system, Fe2+ is the ferrous ion, and Fe3+ is the ferric ion. You can usually identify an element's symbol in the name. A major disadvantage of using classical names for ions is that they do not tell you the actual charge of the ion. A classical name tells you only that the cation has either the smaller (-ous) of the larger (-ic) charge of the pair of possible ions for that element.


A few transition metals have only one ionic charge. The names of these cations do not have a Roman numeral. These exceptions include silver, with cations that have a 1+ charge, as well as cadmium and zing, with cations that have a 2+ charge. Many transition metal compounds are colored and can be used as pigments, which are compounds having intense colors that can be used to make yellow, orange, red, or green paints. Various cadmium compounds range in color from yellow to red and maroon. Prussian blue is an important pigment composed of the transition element iron combined with carbon, hydrogen, and nitrogen.


Daniel Lynch (Pages 257-259):

Polyatomic Ions

- Some ions, called polyatomic ions, are composed of more than one atom. The sulfite anion consists of one sulfur atom and four oxygen atoms. These five atoms comprise a single anion with an overall -2 charge. The formula is written SO4^-2.
- The names of most polyatomic ions end in -ite or -ate.
-Sometimes two or three of the same elements combine in different ratios to form different polys.
-Notice that in all kinds where this is true (sulfite + sulfate) the ITE ending has one less oxygen atoms than in the ATE ending.
-When the formula for a polyatomic ion starts with H, you can think of the H as representing a hydrogen ion combined with another poly.
For Example HCO3^- is a
poly
of H+ and CO3-2. Note that the charge on the new ion is the algebraic sum of the ionic charges.
http://www.youtube.com/watch?v=TZgv21FmEzk

http://www.youtube.com/watch?v=TZgv21FmEzk
Here is a video doing a great job further explaining Polyatomic ions.
15930929-1.png


Pages 260-267: Naming and Writing Formulas for Ionic Compounds

Group Members: Kati West

Kati West (Pages 260-267):

Binary Ionic Compounds

Naming Binary Ionic Compounds

  • a binary compound is composed of two elements and can be either ionic or molecular
  • verify that the compound is composed of a monatomic metallic cation and a monatomic nonmetallic anion
  • place the cation name first, followed by the anion name
  • the cation and anion charges must balance each other to 0

Writing Formulas for Binary Ionic Compounds

  • if you know the name of a binary ionic compound, you can write its formula
  • write the symbol if the cation and then the anion and add whatever subscripts are needed to balance the charges
  • The positive and negative charges must balance to make 0. If the cation, for example, has a +2 charge and the anion has a -1 charge, there must be two anions to balance the cation.
  • If the charges cannot be balanced simply, like in an ion where the cation has a +3 charge and the anion has a -2 charge, you must find the least common denominator of them, in this case, 6.
  • You can also use the crisscross method, in which the numerical value of each ion is crossed over, becoming the subscript for the other ion.
  • If the ratio can be reduced, reduce it as much as possible. For example, if the ration is 4:4, this can be reduced to 1:1.

Compounds with Polyatomic Ions

  • To write the formula for an ionic compound with a polyatomic ion, write the symbol for the cation followed by the formula of the polyatomic ion and balance the charges.
  • Use parentheses to set off the polyatomic ion in a formula when the compound contains more than one polyatomic ion (when more than one polyatomic ion is needed to balance the cation).

Naming Compounds With Polyatomic Ions

  • Recognize that the compound contains a polyatomic ion.
  • To name a compound containing a polyatomic ion, state the cation first and then the polyatomic ion.

Pages 268-270: Naming and Writing Formulas for Molecular Compounds

Group Members: Elaney Marcotte

Elaney Marcotte (Pages 268-270):

Naming Binary Molecular Compounds
  • Binary ionic compounds are composed of ions of a metal and a nonmetal.
  • Binary molecular compounds are composed of two nonmetals that are not ions.
  • When two nonmetallic elements combine, then can do it in multiple ways.
  • Prefixes in the names of the binary molecular compounds help to distinguish the compounds that contain different amounts of the same elements.
  • In a binary molecular compound, the prefix tells how many atoms of the elements are in each molecule of the compound.
    • Mono- indicates one atom
    • Di- indicates two atoms
    • The names of all binary molecular compounds end in –ide.
    • If an element begins with a vowel, then the vowel at the end of the prefix is often dropped.
    • If the first element only has one atom, then the mono- is omitted.
    • Naming binary compounds
      • Confirm the compound is composed of two nonmetals
      • The name has to show the elements and the number of atoms for each
      • Name the elements in order in the formula
      • Use the prefixes to indicate the number of each atom of each elementBinary_Compound_Prefixes.jpg
      • Add the suffix –ide to the second element
  • Example:
    • Name the Compound: N2O
    • This consists of two nonmetals, so it is a binary compound
    • There are two atoms of nitrogen and one atom of oxygen
    • Add the prefix di- to nitrogen, mono- to oxygen, and add –ide to the end of oxygen
    • N2O is dinitrogen monoxide

Writing Formulas for Binary Molecular Compounds
  • Use the prefixes in the names of each element to determine the subscript for each element in the formula
  • Write the symbols for the elements with the correct subscripts
  • Example:
    • Give the formula for dinitrogen tetroxide
      • The prefix di- tells that there are two atoms of nitrogen
      • The prefix tetra- tells that there are four atoms of oxygen
  • The formula for dinitrogen tetroxide is N2O4

Pages 271-273: Naming and Writing Formulas for Acids and Bases

Group Members: Randy Melanson

Randy Melanson (Pages 271-273):
Acids are a group of ionic compounds that contain two unique properties:
  1. An acid contains one or more hydrogen atom in a compound
  2. An acid produces hydrogen ions (H+) when dissolved in water
When naming an acid, you can consider an acid to be an anion combined with as many hydrogen ions (H+) as needed to make the molecule electrically nuetral.
  • The general form of an acid is: H(n)X
X is a monatomic or polyatomic anion and n is a subscript indicationg the number of hydrogen ions that are combined with the anion.

Following three rules can help you name an acid in the general form H(n)X. The naming system depends on the different suffix: -ide, -ite, and -ic.
  1. If the name of the anion (X) ends in -ide, the acid name begins with the prefix hydro-. The stem of the anion has the suffix -ic and is followed by the word acid. Ex: HCl(aq). X = Chloride, hydrochloric acid. H2S(aq). X = Sulfide, hydrosulfuric acid.
  2. When the anion name ends in -ite. the acid name is the stem of the anion wht the suffix -ous, followed by the word acid. Ex. H2SO3(aq), X = sulfite is named sulfurous acid.
  3. When the anion name ends in -ate, the acid name is the stem of the anion with the suffix -ic followed by the word acid. Thus HNO3(aq) is X = nitrate and is named nitric acid.

Writing Formulas for Acids

Use the rules for writing the names of acids in reverse to write the formulas for acids.
  • Rule 1, hydrobromic acid (hydro- prefix and -ic suffix) This must be a combination of hydrogen ion (H+) and bromide ion (Br-). The formula for hydrobromic acid is HBr.
  • Rule 2, phosphorous acid must be composed of hydrogen ion (H+) and phosphite ion (PO3-3). Therefore the formula for phosphorous acid is H3PO3.
  • Rule 3, Carbonic acid must be composed of hydrogen ion (H+) and carbonate ion (CO3-2). Therefore the formula for carbonic acid must be H2CO3.

Names and Formulas for Bases

Bases are named in the same way as other ionic compounds- the name of the cation is followed by the name of the anion.
  1. Write the symbol for the metal cation followed by the formula for the hydroxied ion
  2. Balance the ionic charges just as you do for any ionic compound
  3. Aluminum hydroxide = cation (Al+3) and anion (OH-) balancing the charges gives you Al(OH)3.

Pages 274-278: The Laws Governing Formulas and Names

Group Members: Frank Morley

Frank Morley (Pages 274-278):

The Laws of Definite and Multiple Proportions

  • Rules for naming and writing formulas for compounds are only possible because elements form compounds in predictable ways.
  • The Law of Definite Proportions
    • Because atoms combine in simple whole number ratios the masses of each element is always in the same proportion in a compound regardless of the total mass of the compound
    • Subscripts, in chemical formulas, show the ratio of atoms of each element in a compound
    • Example:
      • The molar mass of hydrogen is 1.007 (1)
      • The molar mass of oxygen is 15.99 (16)
      • In H2O there are two hydrogen atoms (total molar mass of 2) and 1 oxygen atom (total molar mass of 16)
      • So the ratio of H:O is 2:16 which simplifies to 1:8
      • If you had 100 H2O molecules instead of just 1 the total molar mass of the compound changes to be greater, but the ratios of H:O stay the same
      • It would end up being 20:160 which again simplifies to 1:8
  • The Law of Multiple Proportions
    • Whenever the same two elements form more than 1 type of compound the different masses of one element that combine with the same mass of the other element are in a ratio of small whole numbers.

Practicing Skills: Naming Chemical Compounds

  • PAY CLOSE ATTENTION THIS CAN GET CONFUSING GOING BACK AND FORTH
  • Is there an H to start the compound?
    • Yes: it is an acid
    • No continue on
  • Are there more than 2 elements?
    • If no:
      • It is a binary compound so it ends in "ide"
      • Find if the 1st element named is a metal
        • If yes:
          • Find if it is is group A
            • If yes:
              • Name the ions regularly
              • Ex: Barium Sulfide
            • If no:
              • Name ions, but use the roman numeral
              • NOT NECESSARY FOR THE TEST
        • If no: external image law%20of%20multiple%20proportion%201.gif
          • it is binary molecular
          • Use prefixes
          • Remember 2 negatives= prefix
    • If yes:
      • The compound contains a polyatomic ion
      • name generally ends in "ite" or "ate"
      • Find if the 1st element is in group A
        • If yes:
          • Name the ions (polyatomic ion as part of it)
        • If no
          • Again with the roman numerals, DO NOT NEED THEM
    • Remember with polyatomic ions, parenthesis may be needed depending on the charges

Practicing Skills: Writing Chemical Formulas

  • Ide= binary compound
  • ite/ate= includes a polyatomic ion
  • Prefixes mean it is a molecular compound
  • Roman numeral shows charge
SEE CHART FROM CLASS

PRACTICE TIME!

8 AND 18 WERE ILLEGAL SO THEY WERE ELIMINATED, I MAY HAVE MISSED OTHER ILLEGAL ONES.

Name the following ionic compounds:

1) NH4Cl _

2) Fe(NO3)3 _

3) TiBr3 _

4) Cu3P _

5) SnSe2 _

6) GaAs _

7) Pb(SO4)2 _

9) Mn2(SO3)3 _

10) Al(CN)3 _

Write the formulas for the following compounds:

11) chromium (VI) phosphate _

12) vanadium (IV) carbonate _

13) tin (II) nitrite _

14) cobalt (III) oxide _

15) titanium (II) acetate _

16) vanadium (V) sulfide _

17) chromium (III) hydroxide _

19) lead (II) nitride _

20 silver bromide _

NUMBER 1 WAS ILLEGAL BUT AGAIN I MAY HAVE MISSED OTHERS:
Write the formulas of the following ionic compounds:
2) lead (II) sulfate _

3) lead (IV) hydroxide _

4) copper (II) acetate _

5) beryllium chloride _

6) ammonium chromate _

7) silver oxide _

8) potassium sulfide _


Write the names of the following ionic compounds:
9) KI _

10) Mn2(SO3)­7 _

11) SnBr4 _

12) Mg3P2 _

13) NaF _

14) Sr(MnO4)2 _

15) Cr(PO4)2 _

16) Al2Se3 _