Chapters 7 & 8
Bonding and Nomenclature
Editor: Neal McGovern

external image Hydroxyl_Real_3.jpg<--example of polarity in a water molecule
example of nitrite, a polyatomic ion------>external image nitrite3.jpg
difference between ionic and covalent bonds
difference between ionic and covalent bonds

Diference between ionic and covalent bonds


Group 1
Coeditor: Kayla Anghinetti
Group Members: Elena Conroy, Bryan Dextradeur, Lindsey Chou

Ions
Pgs.187-193
By: Elena Conroy

Valence electrons
  • Electrons in the highest occupied energy level of an element’s atom
  • The number of valence electrons in a atom of a representative element is its group number of the periodic table
  • Valence are usually the only electrons used in chemical bonds
  • Electron dot structure are diagrams that show the valence electrons as dots


http://butane.chem.uiuc.edu/cyerkes/chem102ae_fa08/homepage/chem102aefa07/lecture_notes_102/lecture%2012%20.htm
http://butane.chem.uiuc.edu/cyerkes/chem102ae_fa08/homepage/chem102aefa07/lecture_notes_102/lecture%2012%20.htm

Octet Rule
· Noble gases are un-reactive in chemical reactions
· Gilbert Lewis used this fact to explain the octet rule in forming compounds, atoms tend to achieve the electron configuration of a noble gas
· Atoms of the metallic elements tend to lose their valence electrons leaving a complete octet in the next-lowest energy level
· Atoms of some non-metallic elements tend to gain electrons or to share electrons with another nonmetallic element to achieve a complete octet

http://www.hibbing.edu/chem/V.12/page_id_14649.html
http://www.hibbing.edu/chem/V.12/page_id_14649.html















Formation of Cations
  • An atom’s loss of valence electrons produces a cation, or a positively charged ion
  • When an element is metallic, it’s cation is the same as its element name, although there is many different chemical differences between them
  • The most common cations are ones produce from loss of valence electrons from metals
  • Transition metals have cations that vary

Formation of Anions
  • The gain of negatively charged electrons by neutral atom produces an anion
  • The name of an anion usually ends in -ide
  • Metals of nonmetallic elements attain noble-gas electron configurations more easily by gaining electrons than losing them because their valence shells are fuller
  • Halide ions are produced when atoms of chlorine and other halogens gain electrons

http://crescentok.com/staff/jaskew/ISR/chemistry/anions.gif
http://crescentok.com/staff/jaskew/ISR/chemistry/anions.gif
















Elena Conroy with above images

Bryan Dextradeur

Pages 194-195

Formation of Ionic Compounds


  • Ionic Compounds: Compounds composed of cations and anions.
    • Ionic Compounds are usually composed of metal cations and nonmetal anions.
    • EXAMPLE: Sodium Chloride (Table Salt) is composed of Sodium cations and Chloride anions.
  • Although they are composed of ions, ionic compounds are electrically neutral.
    • The total positive charge of the cations equals the total negative charge of the anions.
  • Anions and cations have opposite charges, so they attract each other through electrostatic forces.
    • Ionic Bonds: The electrostatic forces that hold ions together in ionic compounds.
    • EXAMPLE: In Sodium Chloride, sodium has one valence electron that it can easily lose to become a stable cation and chlorine has seven valence electrons, so it can easily gain one to become a stable anion, so the two elements can easily combine by sodium giving one electron to chlorine, causing both ions to have stable octets.
  • Chemical Formula: Shows the kinds and numbers of atoms in the smallest representative unit of a substance.
    • Chemists write chemical formulas to represent the composition of substances.
    • EXAMPLE: NaCl is the chemical formula for Sodium Chloride. It does not represent a single discrete unit.
  • Formula Unit: The lowest whole-number ratio of ions in an ionic compound.
    • EXAMPLE: In Sodium Chloride, the lowest whole-number ratio of ions is 1:1 because there is one Na+ for each Cl-, so the formula unit for Sodium Chloride is NaCl. However, in the ionic compound Magnesium Chloride , the ratio of ions is 1:2 because there is one Mg2+ cation for every two Cl- anions (there are twice as many Cl ions, each with a 1- charge as Magnesium ions, each with a 2+ charge), so the ions balance out making the ionic compound electrically neutral. Because of this, the formula unit for Magnesium Chloride is MgCl2 .
ionic_compounds_fig1.gifnacl.jpg
Above Images: Bryan Dextradeur


Properties of Ionic Compounds
p. 196-200, Kayla Anghinetti
I. Most ionic compounds are crystalline solids at room temperature.
A. component ions in such crystals are arranged in repeating 3-dimensional patterns.
B. ex. in solid NaCl, each sodium ion is surrounded by six chloride ions and each chloride ion is surrounded by six sodium ions.
C. large attractive forces result in a very stable structure
II. Ionic compounds generally have high melting points.
A. The coordination number of an ion is the number of ions of opposite charge that surround the ion in a crystal.
B. Another characteristic of ionic compounds has to do with conductivity.
1. Ionic compounds can conduct an electric current when melted or dissolved in water.
2. If a voltage is applied across this molten mass, cations move freely to one electrode and anions migrate to the other.
3. Ionic compounds also conduct electricity if they are dissolved in water. When dissolved, the ions are free to move about in the aqueous solution.

external image images?q=tbn:ANd9GcS4eFeGLymkKFDE4aZbBYvlDEizQanKYiI8FLsw4FIDIv_jQnvjGAexternal image images?q=tbn:ANd9GcTnWn2VZ6VHrXSRwk5g1S_8T3w07Tf2AirZQMAM__NqwYnkZ76r
Fluorite (CaF2) & Fluorite Crystal Structure

external image images?q=tbn:ANd9GcQ6XWUJQ986IXNBhKadg9AYwA1gVnA6n3U15fQWBUFSqJaG0-bwvAexternal image images?q=tbn:ANd9GcR23y6thSCYGqq1-8_DUuO-apHfGeihufcVSv__U4Zx_jQ7PGDa
Calcite (CaCO4) & Calcite Crystal Structure
Kayla Anghinetti with above images

Lindsey Chou
7.3 “Bonding in Metals” (pg. 201-205)

Metallic Bonds and Metallic Properties
· Metals = composed of closely packed cations (ions with a positive charge)
· The attraction between the free-floating electrons and positive metal ions is responsible for metallic bonds
· Valence electrons of metal atoms can be modeled as a “sea of electrons”, which means they are mobile and “drift” from one area of the metal to another
Sea_of_Electrons.jpg
(Lindsey Chou)

· This gives metals their properties of conductivity, ductility and malleability
Ø Conductivity is the ability of metals to conduct an electrical current. They can do this because electrons flow freely in them.
Ø Ductility is how metals can be made into wires. Malleability is how metals can be forced into shapes. Both of these are possible because the drifting electrons act as “insulation” and allow the metal cations to slide past each other easily when put under pressure.
Ø In other substances, like ionic crystals, a force causes ions with the same charge to repel from each other, and the substance shatters.

Crystalline Structure of Metals
· Metals are crystalline
· The arrangement of atoms in metals is tightly packed in certain patterns
· These patterns include body-centered cubic, face-centered cubic and hexagonal close-packed
Ø Body-centered cubic arrangement: Every atom has eight neighbors. Some examples of metals with this arrangement are sodium, potassium and chromium.
Body_Centered_Cubic.jpg
(Lindsey Chou)
Ø Face-centered cubic arrangement: Every atom has twelve neighbors. Some examples are silver and gold.
face_centered_cubic.jpg
(Lindsey Chou)
Ø Hexagonal close-packed arrangement: Every atom has twelve neighbors, but is different from the face-centered cubic arrangement because it has a hexagonal shape. Some examples are magnesium and cadmium.
hexagonal_close_packed_structure.gif
(Lindsey Chou)
Alloys
· Alloys are mixtures of at least two elements when at least one component is a metal
· Properties of alloys are often superior to the properties of the separate elements that make them up
· Steels are the most important alloys today
Ø Iron and carbon, along with boron, chromium, manganese, molybdenum, nickel, tungsten, and vanadium are the main elements used in most steels
Ø The properties of steel are very useful, and can include resistance to corrosion, ductility, hardness and toughness
· There are several ways alloys can form:
Ø Substitutional Alloy: Atoms of the components can replace each other in the crystal if they are about the same size
Ø Interstitial Alloy: Smaller atoms can fit into the spaces (also called interstices) between larger atoms if the atoms of the components are different sizes – steels are interstitial alloys
Building with Alloys
· Steel is an important material in modern architecture
· Used to make buildings taller, stronger, and lighter
· Steel-frame and reinforced-concrete construction are important in construction of skyscrapers (high-rise buildings)

Group 2
Co-editor: Jessen Foster
Ellie Kawa
Hannah Kumlin
Chris Kelly

Ellie Kawa
pgs 213-216
Molecular Compounds

Molecules and Molecular Compounds

  • noble gas elements exist as atoms
    • (Ellie Kawa)
      (Ellie Kawa)
      monatomic- consist of single atoms
  • atoms of some elements combine to form salts- high melting points
    • ex: sodium chloride
  • other compounds have different properties
    • exs: hydrogen chloride, water
  • covalent bond- joins atoms that are held together by sharing electrons
  • molecule- a neutral group of atoms joined together by covalent bonds
  • diatomic molecule- molecule consisting of two atoms
    • ex: oxygen moleculeexternal image Covalent.png
  • atoms of different elements can combine chemically to form compounds
    • atoms are bonded together to form molecules
  • molecular compound- compound composed of molecules
    • molecules of a given molecular compound are all the same
  • Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.
    • Many are gases or liquids at room temperature
    • ionic compounds- formed from a metal combined with a non-metal
    • molecular compounds- composed of atoms of two or more nonmetals (Ellie Kawa)
Molecular Formulas
  • molecular formula- the chemical formula of a molecular compound
  • A molecular formula shows how many atoms of each element a molecule contains.
    • Subscript written after the symbol indicates the number of atoms of each eleexternal image co2molecule_0.jpgment in the molecule
      • if there is only one atom, the subscript 1 is omitted
        • exs: carbon dioxide, ethane
    • molecular formula reflects the actual number of atoms in each molecule
    • describes molecules consisting of one element
    • does not tell you about a molecule's structure
      • does not show either the arrangement of the various atoms in space or which atoms are covalently bonded to one another
  • carbon dioxide- gas produced by the complete burning of carbon, occurs in our atmosphere and dissolves in sea water


Jessen Foster
Pages 217-222

Octet Rule in Covalent Bonds
-in forming covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases
-atoms usually acquire eight electrons by sharing electrons→ forming an octet
Single Covalent Bonds
-two atoms held together by sharing electrons are joined by a single covalent bond
-structural formula: represents covalent bonds by dashes and shows the arrangement of covalently bonded molecules
-unshared pair: a pair of valence electrons that is not shared between atoms
Double and Triple Covalent Bonds
-atoms form double or triple covalent bonds if they can attain noble gas structure by sharing two or three pairs or electrons
(common sense: double shares two pairs, triple shares three)
external image bonds.gifJessen Foster
Diatomic Elements
-elements that, when in nature, automatically link up with another atom of the same element
→examples: O2, H2, Cl2
-in other words, you wont find just one atom of oxygen; it comes in pairs so that it is more stable

Hannah Kumlin

Pages 223-229


Coordinate Covalent Bonds

-A coordinate covalent bond is a covalent bond in which one atom contributes both bonding electrons. In a structural formula, they are shown as arrows that point from the atom donating the pair of electrons to the atom receiving them.
For example, the structural formula of carbon monoxide, with two covalent bonds and one coordinate covalent bond, is 195014.jpg
*In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms.

-A polyatomic ion, such as NH4+, is a tightly bound group of atoms that has a positive or negative charge and behaves as a unit. Here is an example of an ammonium ion which is a polyatomic ion.
ammonium_ion.GIF<--Hannah Kumlin

-Most polyatomic cations and anions contain both covalent and coordinate covalent bonds. Therefore, compounds containing polyatomic ions include both ionic and covalent bonding.
*See page 224 in your textbook for a table of some common molecular compounds*

- Remember, the electron dot structure for a neutral molecule contains the same number of electrons as the total number of valence electrons in the combining atoms. The negative charge of a polyatomic ion shows the number of electrons in addition to the valence electrons of the atoms present. Because a negatively charge polyatomic ion is part of an ionic compound, the positive charge of the cation of the compound balances these additional electrons.

Here's some practice problems (or see page 225) :
1. Draw the electron dot structure of the hydroxide ion (OH-).

2. Draw the electron dot structure of the polyatomic boron tetrafluoride anion (BF4-).

3. Draw the electron dot structures for sulfate (SO4-2) and carbonate (CO3-2). Sulfur and carbon are the central atoms, respectively.

4. Draw the electron dot structure fore the hydrogen carbonate ion (HCO3-). Carbon is the central atom, and hydrogen is attached to oxygen in this polyatomic anion.

Bond Dissociation Energies

- The energy required to break the bond between two covalently bonded atoms is known as the bond dissociation energy. This is usually expressed as the energy needed to break one mole of bonds.
- A large bond dissociation energy corresponds to a strong covalent bond. A typical carbon-carbon single bond has a bond dissociation energy of about 347 kJ/ mol.

Resonance


Ozone_molecule.png<--Hannah Kumlin
^this is an ozone molecule. The ozone molecule has two possible electron dot structures, both shown above. Notice that the structure on the left can be converted to the one on the right by shifting electron pairs without changing the positions of the oxygen atoms. As shown, the electron dot structures suggest that the bonding in ozone consists of one single coordinate covalent bond and one double covalent bond. Early scientists used the double headed arrows to show that two or more structures are in reasonance.

-Double covalent bonds are usually shorter than single bonds, so it was believed that the bond lengths of ozone were unequal. However, experimental measures showed that the two bonds in ozone are the same length. This result can be explained if you assume that the actual bonding in the ozone molecule is the average of the two electron dot structures. The electron pairs do not actually resonate back and forth.

-The actual bonding of oxygen atoms in ozone is a hybrid, or mixture, of the extremes represented by the resonance forms.
-The resonance structure is a structure that occurs when it is possible to write tow or more valid electron dot formulas that have the same number of electron pairs for a molecule or ion. Resonance structures are simply a way to envision the bonding in certain molecules. Double-headed arrows are used to connect resonance structures.

Exceptions to the Octet Rule

-The octet rule provides guidance for drawing electron dot structures. For some molecules or ions, however, it is impossible to draw structures that satisfy the octet rule.
-The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. There are also molecules in which an atom has fewer, or more, than a complete octet of valence electrons.

-Several molecules with an even number of electrons, such as some compounds of boron, also fail to follow the octet rule. This may happen because an atom acquires less than an octet of eight electrons.

-A few atoms, especially phosphorus and sulfur, sometimes expand the octet to include ten or twelve electrons.


Bonding Theories (p.230-236)


By: Chris Kelly

Molecular Orbitals


-Molecular Orbitals are orbitals that apply to the entire molecule
-Atomic Orbitals and Molecular Orbitals have similarities
-Atomic orbitals belong to a particular atom
-Molecular orbitals belong to a molecule as a whole
-Atomic orbitals are filled with 2 electrons
-2 electrons are required to fill a molecular orbital
-A bonding orbital is a molecular orbital that can be occupied by 2 electrons of a covalent bond

Sigma Bonds

-Sigma bonds are formed when two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei
-The symbol is the Greek letter sigma (σ)

Pi Bonds


-A pi bond is a covalent bond in which the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms

VSEPR Theory


-Electron dot structure and structural formula don’t show molecules as three dimensional
-Tetrahedral angle-a bond angle of 109.5 degrees that results when a central atom forms four bonds directed toward the center a regular tetrahedron
-Valence-shell electron-pair repulsion theory, or VSEPR theory explain the three-dimensional shape
-repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible

Chris Kelly
Chris Kelly
<--Chris Kelly
Hybrid Orbitals

-Orbital Hybridization provides information about both molecular bonding and molecular shape
-Hybridization is the mixing of several atomic orbitals to form the same total number of equivalent hybrid orbitals

Group 3

Co-Editor-Shannon Leavey

members- Shannon Leavey, Allie Chabot, and Teresa Lynch

8.4 Polar Bonds and Molecules

By: Allie Chabot (Pg. 237-244)

Bond Polarity
· Covalent bonds involve electron sharing between atoms.
· The bonding pairs of electrons in covalent bonds are pulled, between the nuclei of the atoms sharing the electrons.
· Nonpolar covalent bond- a covalent bond in which the electrons are shared equally by the two atoms.
· Polar covalent bond (Polar bond) - is a covalent bond between atoms in which the electrons are shared unequally.
· *The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge.*



Electronegativity difference range
Most probable type of bond
Example
0.0-0.4
Nonpolar covalent
H-H (0.0)
0.4-1.0
Moderately polar covalent
H-CL (0.9)
1.0-2.0
Very polar covalent
H-F (1.9)
Greater than 2.0
ionic
NaCl (2.1)


Polar Molecules
· Polar molecule- a molecule in which one side of the molecule is slightly negative and the opposite side is slightly positive.
· Dipole- a molecule that has two poles, or regions, with opposite charges.
· *When polar molecules are placed between opposite charged plates, they tend to become oriented with respect to the positive and negative plates.*
· The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds.
external image dipolh2o.gif(Allie Chabot)
(Polar bond)

Attractions between Molecules
· Molecules can attract each other in a variety of forces.
· Intermolecular attractions are weaker than either ionic or covalent bonds.
· These attractions are responsible for determining whether a molecular compound is a gas, liquid, or solid at a given temperature.

Van der Waals Forces
· Van der Waals Forces- the two weakest intermolecular attractions- dispersion interactions and dipole forces. This was named after a Dutch chemist Johannes van der Waals.
· Dipole Interactions- intermolecular forces resulting from the attraction of oppositely charged regions of polar molecules.
· Dipole interactions are similar to, but much weaker than ionic bonds.
· Dispersion forces- attractions between molecules caused by the electron motion on one molecule affecting the electron motion on the other through electrical forces; these are the weakest interactions between molecules.
· The strength of dispersion forces generally increases as the number of electrons in a molecule increase.

Hydrogen Bonds
· The hydrogen in water molecules acquires a slightly positive charge.
· Hydrogen Bonds- attractive forces in which hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom.
· Hydrogen bonds always involve hydrogen.
· Hydrogen bonds are the strongest of the intermolecular forces.
external image Image8.gif(Allie Chabot)
(Hydrogen Bond)

Intermolecular Attractions and Molecular Properties
· The physical properties of a compound depend on the type of bonding it displays- whether it is ionic or covalent.
· The melting and boiling points of most compounds composed of molecules are low compared to those of ionic compounds.
· Network Solids- a solid which all of the atoms are covalently bonded to each other.
· *Melting a network solid would require breaking covalent bonds throughout the solid. *
· Diamond is an example of a network solid.


Characteristics
Ionic compound
Covalent compound
Representative unit
Formula unit
Molecule
Bond formation
Transfer of one or more electrons between atoms
Sharing of electron pairs between atoms
Type of elements
Metallic and nonmetallic
Nonmetallic
Physical state
Solid
Solid, liquid, or gas
Melting point
High (usually above 300 degrees Celsius)
Low (usually below 300 degrees Celsius)
Solubility in water
Usually high
High to low
Electrical conductivity of aqueous solution
Good conductor
Poor to no conducting

Shannon Leavey
Section 9.1

Monatomic Ions
-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
-when the metals in Groups 1A, 2A, and 3A lose electrons, they form cations with positive charges to their group number
-the charge pf any ion of a group of nonmetals is determined by subtracting 8 from the group number
-the charges of the cations of many transition metal ions must be determined from the number of electrons lost

Polyatomic Ions
-these are composed of more than one atom
-the names of most polyatomic anions end in –ite or –ate
-ite -ate
SO32-, sulfite SO42- sulfate
NO2- nitrite NO3- nitrate
ClO2- chlorite ClO3- chlorate

9.2 Naming and Writing Formulas for Ionic Compounds

By: Teresa Lynch

· Binary Ionic Compounds:
· Antoine-Laurent Lavoisier and other chemists worked on creating a systematic way to name compounds. Before them chemical compounds were named whatever the person who discovered wanted to name it.

· Naming Binary Ionic Compounds: A binary compound is two or more elements (either molecular or ionic).
· To name any binary ionic compound, place the cation name first, followed by the anion name.
· Ex: SrF2 is strontium fluoride

· Writing formulas for Binary Ionic Compounds: Write the symbol of the cation and then the anion. Add whatever subscripts are needed to balance the charges.
· Ex: potassium chloride is KCl because the charges are balanced.
· You can also you the crisscross method- Ex: Iron (III) oxide is Fe3+ and oxide is O2- if you crisscross the charges (making them the subscripts of the opposite element) you have your answer: Fe2 O3 (*make sure to check for simplifying)

· Compounds with Polyatomic Ions:

· Remember that –ate and –ite means that the compound contains a polyatomic ion and oxygen.
· Write the symbol for the cation followed by the formula for the polyatomic ion and balance the charges.
· Ex: calcium carbonate contains the monatomic ion Ca2+ and the polyatomic ion CO32- together they make: CaCO3


oyster_pic.gif


Oysters produce calcium carbonate for their shells and pearls. (T. Lynch)

· Naming Compounds With Polyatomic Ions: In order to name compounds with polyatomic ions follow these steps:
  1. Recognize that there is a polyatomic ion in the compound
  2. State the cation first and then the anion just like with binary ionic compounds

Group 4

David Monti, Lauren Murphy, Emily Stewart, Nikki Sheehan

Co-editor Nikki Sheehan

David Monti
(268-270) Section 9.3

NAMING AND WRITING FORMULAS FOR MOLECULAR COMPOUNDS

Naming Binary Molecular Compounds
- Binary ionic compounds are composed of the ions of two elements, a metal and a nonmetal.
- Binary molecular compounds are also composed of two elements, but both of them are nonmetals and they are not ions
- Because they are composed of molecules instead of ions, ionic charges cannot be used to write formulas or name them
- They can combine in more than one way ex. CO and CO2
- Naming molecular compounds is based on a code of prefixes:
Mono- 1
2_6_table.gif
Di- 2
Tri- 3
Tetra- 4
Penta- 5
Hexa- 6
Hepta- 7
Octa- 8
Nona- 9
Deca- 10



- These prefixes indicate the different amounts of each element in a molecular compound
- The Prefix in the name of a binary molecular compound tells how many atoms of each element are present in each
molecule of the compound Ex. Dinitrogen monoxide (N2O
- The names of all binary molecular compounds end in ide
- The vowel at the end of a prefix is omitted when the element name begins with a vowel
- The steps for naming binary molecular compounds are fairly similar:
1. Confirm that the compound is a binary molecular compound
2. Confirm that is composed of two nonmetals
3. Name the elements in order listed in the formula
4. Use prefixes to indicate the number of each element
5. Omit the prefix mono when the formula contains only one atom of the first element in the name
images.jpeg(David Monti)
Writing Formulas for Binary Molecular Compounds
- Use the prefixes in the name to tell you the subscript of each element in the formula. They write the correct symbols for the two elements
with the appropriate subscripts.
- Ex. Silicon carbide has no prefixes so the subscripts of silicon and carbon must be one. Thus the formula for silicon carbide is SiC.

Naming and Writing Formulas for Acids and Bases
Lauren Murphy

Naming Acids
· Acids are a group of ionic compounds with unique properties
· Contains one or more hydrogen atoms and produces hydrogen ions
· HnX, X is monatomic or polyatomic anion and n is a subscript indicating the number of hydrogen ions combined with the anion
· When name of anion ends in ide acid name begins with hydro/ stem has suffix ic and is followed by acid
· When anion ends in ite, the name is the stem of the anion with the suffix ous followed by the word acid
· When anion ends in ate the acid name is stem of anion with suffix ic followed by acid
Writing Formulas for Acids
· Use the rules above, but in reverse
Names and Formulas for Bases
· Base is an ionic compound that produces hydroxide when dissolved in water
· They’re named the same way as other ionic compounds, cation name is followed by anion name
FORMULAS.png(Lauren Murphy)

Emily Stewart
Practicing Skills: Naming Chemical Compounds
p276-277
external image chem_n7.jpg
(Emily Stewart)
Use this flowchart to help determine what prexifes, Roman numerals, and endings (-ate, -ide, -ite) to use when naming chemical compounds.
Apply the flowchart's general formula to name compounds.
P and Q can be atoms, monatomic ions, or polyatomic ions.
To name a chemical compound from its chemical formula, answer the questions in the flowchart.

Emily Stewart
Practicing Skills: Writing Chemical Formulas
p278


Some guidelines for writing a chemical formula:
1. An -ide ending generally indicates a binary compound.
2. An -ite ending = polyatomic ion that includes oxygen in the formula.
3. Prefixes generally indicate that a compound is molecular.
4. A Roman numeral after the name of a cation shows the cation's ionic charge.


Use this flowchart and follow the arrows to help you write chemical formulas.

A periodic table, as well as the tables on pages 254 and 257 in the textbook, might help you.external image chem_n6.jpg
(Emily Stewart)

For practice, check out this worksheet: http://cfbstaff.cfbisd.edu/louiep/images/2009-2010/formula%20writing%20flowchart%20001.jpg


The Laws of Definite and Multiple Proportions

(247-245) Nikki Sheehan


  • Compounds form from elements in predictable ways
Law of Definite Proportions
  • Subscripts tell the ratio of atoms of each element in the compound
  • Ratios of atoms can also be expressed as ratios of masses
  • For example, if you take 100.00 g of magnesium sulfide and break it down into its elements, you will get 43.13 g of magnesium and 56.87 g of sulfur
  • The ratio of these masses is then 43.13/56.87 (0.758:1)
  • This mass ratio does not change, no matter the size of the sample
  • The law of definite proportions states: in samples of any chemical compound, the masses of the elements are always in the same proportions
Law of Multiple Proportions
  • Compounds formed by the same elements can have different chemical and physical properties
  • This is because the mass ratios of the elements are different
  • For example, hydrogen peroxide and water are both formed from hydrogen and oxygen
  • In hydrogen peroxide, the mass ratio of oxygen to hydrogen is always 16:1
  • In water, the mass ratio of oxygen to hydrogen is always 8:1
  • The law of multiple proportions states: whenever the same two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers

external image 20070505150103!Hydrogen-peroxide-3D-vdW.pngexternal image 500px-Water_molecule_2.svg.png
Hydrogen Peroxide (H2O2) vs Water (H2O) - Nikki Sheehan