Chapters+7,+8,+and+9


 * Ionic, Metallic, and Covalent Bonding and Chemical **** Names and Formulas **

Katherine Perry-Editor

Introduction to the Wiki: This page covers the topics discussed in Chapters 7, 8, and 9 in Prentice Hall's //Chemistry// textbook. The seventh chapter focuses on ionic and metallic bonds. This chapter describes ions, bonds, and compounds. Additionally, it includes information on valence electrons, the Octet Rule, cations, anions, and alloys. Chapter 8 focuses on molecular compounds, covalent bonding, bonding theories, and polar bonds and molecules. This chapter has additional information on various molecular topics, covalent bonding, resonance, molecular orbitals, the VSEPR Theory, and polarity. Finally the ninth chapter focuses on chemical names and their formulas. This chapter describes ions, formulas for their ionic and metallic compounds, and how to write the names and formulas. Additionally, it includes information about different types of ions and how to write formulas and names of acids and bases. This wiki-page outlines these chapters in their entirety, and has included all information crucial to the understanding of these chapters.

Group 1 Christos Anastos- 187-190 Heather Bowditch- 191-195 Maggie Bie- 196-201
 * (Co-Editor)** Kendyl Barron- 202-203

Group 2 Tom DeMarco- 213-216 Marion Burdick- 217-229 Eileen Corkery- 230-236
 * (Co-Editor)** Nick Brault

Group 3 Liz Sieber- 239-243 and 253 Courtney Gareau- 254-260 Evan Grandfield- 261-266
 * (Co-Editor) ** Nina DeMeo- 237-238

Group 4 Mike McShane- 271-273 Steve Denison- 274-275 Andrew Sciotti- 268-270
 * (Co-Editor)** Adam Shanahan- 276-278

__ Crystalline Structure of Metals __
The structure of metallic bonds are close-packed in set patterns known as crystalline structures. Metals containing one atom are the simplest form of crystalline structures. These compact arrangements are possible because metal atoms are of the same size. Arrangments include: --Body-centered cubic: each atom has eight neighbors. ex: sodium, potassium. iron, chromium, tungsten --Face-centered: each atom has 12 neighbors. ex: copper, silver, gold, aluminum, lead --Hexagonal close-packed: each atom has 12 neighbors, but structure has a hexagonal shape. ex: magnesium, zinc, cadmium

__ Alloys __ Most metals utilized in everyday life are not in their pure forms but //alloys.// Alloys are mixtures composed of two or more elements, one of which is a metal. --Sterling silver is an alloy composed of silver and copper. --Cast iron alloy is iron and carbon. --Stainless steel alloy is iron, chromium and nitrogen.

The properties of alloy can be superior to compounds in their pure forms. Sterling silver is known to be more durable than pure silver. Steel alloys are widely used today in construction, machinery, tools and kitchen hardware. Substutional alloys from when component atoms are the same size as metal atoms and replace them in their crystalline pattersn. Smaller atoms may be able to fit into interstices, spaces between atoms, between larger atoms. These alloys are known as interstitial alloys, such as most steels.



Properties of Ionic Compounds - Maggie Bie (pgs 196-198)
Key Concept - Most ionic compounds are crystalline solids at room temperature Key Concept - Ionic compounds generally have high melting points
 * The component ions in such crystals are arranged in repeating 3D patterns

Key Concept: Ionic compounds can conduct an electrical current when melted or dissolved in water media type="youtube" key="bzVKD2Pw1D0?fs=1" height="385" width="480"
 * __Coordination Number -__﻿** (of an ion) is the number of ions of opposite charge that surround the ion in a crystal
 * When an ion is melted, the orderly crystal structure breaks down
 * if voltage is applied across this molten mass, cations migrate freely to one electrode and anions migrate to each other
 * this allows electricity to flow between the electrodes through an external wire
 * ionic compounds also conduct electricity if they are dissolved in water

Metallic Bonds and Metallic Properties
Key Concept- the valence electrons of metal atoms can be modeled as a sea of electrons - the valence electrons are mobile and can drift freely from one part of the metal to another - forces of attraction that hold metals together Sea of Electrons Model media type="youtube" key="ap5pHBWwpu4?fs=1" height="385" width="480"
 * metals are made up of closely packed cations rather than neutral atoms
 * __Metallic Bonds-__** consist of the attraction of the free-floating valence electrons for the positively charged metal ions
 * As electrons enter one end of a bar of metal, an equal number leave the other end
 * Metals are ductile (can be drawn into wires) and malleable (can be hammered or forced into shapes)
 * a sea of drifting valence electrons insulated the metal cations from one another
 * if a metal has pressure, the cations easily slide past one another
 * in contrast, an ionic crystal pushes the ions into contace. They repel and the crystal shatters

Valence Electrons- Christos Anastos Electron dot structure
 * //__Valence Electrons-__// electrons in the highest occupied energy level
 * the amount of valence electrons determines the properties of an atom
 * to find the number of valence electorns in an atom you look at it's group number
 * Electron dot structures show valence elctrons as dots

The Octet Rule

 * //__Octet Rule-__// In forming compounds atoms tend to acheive the elctron 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 nonmetallic elements tend to gain electrons or to share electrons with another nonmetallic element to achieve a complete octet.

Formation of Cations
Formation of Anions--Heather Bowditch (pg.191-195) **anion **- atom or group of atoms with a negative charge
 * an atom's loss of valence electrons produces a cation or a positively charged ion.
 * for metallic elements the name of the cation is the same as the element. For example Na is Na+
 * The most common cations are formed by metallic elements losing valence electrons

- produced when a neutral atom gains electrons - name //usually// in //-ide// ~for example: a Fluorine atom (F) forms a fluor**ide** ion - non-metallic or metalloid atoms (top right side of periodic table) have valence shells that are mostly full with 5,6, or 7 electrons. Therefore, they form ions mainly by gaining electrons. - halide ions are formed when halogen atoms (group 7) gain electrons--since they have 7 valence electrons, they only need to gain 1 to be full -this means that all halogen ions have a charge of -1, because they gain 1 electron - elements in group 6 have a charge of -2 because they gain 2 electrons and elements in group 5 have a charge of -3  Formation of Ionic Compounds - **ionic compounds **- compounds composed of cations and anions - <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">usually metal cations and nonmetal/metalloid anions - <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">always have a neutral charge because positive charge of cation and negative charge of anion cancel out - <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">held together by **ionic bonds** <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: normal; margin: 0in 0in 0pt 0.75in; text-indent: -0.25in;">**﻿** - <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">one example is sodium chloride: <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: normal; margin: 0in 0in 0pt 0.75in; text-indent: -0.25in;">~ sodium has 1 valence electron, chlorine has 7 valence electrons <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: normal; margin: 0in 0in 0pt 0.75in; text-indent: -0.25in;">~ when the two react and form a compound, sodium gives its electron to chlorine <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: normal; margin: 0in 0in 0pt 0.75in; text-indent: -0.25in;">~ now they both are stable <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">- **<span style="font-family: 'Times New Roman',serif; font-size: 12pt;">chemical formula **<span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">- a representation of the composition of a substance using the kinds and numbers of atoms in the substance <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: normal; margin: 0in 0in 0pt 0.75in; text-indent: -0.25in;">~ for example: sodium chloride is NaCl - **<span style="font-family: 'Times New Roman',serif; font-size: 12pt;">formula unit **<span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">- lowest whole-number ration of ions in an ionic compound - <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">because Na has a charge of +1 and Cl has a charge of -1, only one of each element is needed to form a neutral compound - <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">therefore, NaCl is the chemical formula - <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">another example is magnesium chloride

· <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">Mg has a +2 charge and Cl has a -1 charge · <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">to become neutral, there must be 2 Cl (2 times -1= -2) · <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">therefore, the chemical formula is MgCl ₂

Group 2 (Pages 213-236)
Nick Brault (co-editor), Tom DeMarco, Marion Burdick, Eileen Corkery

** Section 8.1- Molecular Compounds ** Thomas DeMarco Molecules and Molecular Compounds Matter takes on many forms. It can be **monatomic** which means it is made up of just one atom. Atoms of the same elements can be combined to form salts like sodium chloride. Other compounds can have very different properties. Hydrogen chloride is a gas at room temperature, but water is a liquid at room temperature. A **covalent bond** is a bond where atoms share electrons. Many elements in nature form molecules. A **molecule** is a neutral group of atoms joined together by covalent bonds. A **diatomic molecule** is a molecule consisting of two atoms. An example of this is an oxygen molecule which consists of two oxygen atoms. A **molecular compound** is a compound composed of molecules. Molecular compounds have relatively lower melting and boiling points than ionic compounds. Most ionic compounds are formed from a non-metal and a metal combining, but most molecular compounds are formed from two or more non-metals. ** Example of an ionic compound: ** ** Salt (NaCl) **

** Example of a molecular compound: ** ** Water (H2O) **

A **molecular formula** is the chemical formula of a molecular compound. It shows how many atoms of each element a compound contains. The molecular formula of water is H2O. The subscript shows how many atoms of a particular element are in the molecule. If there is one atom the subscript is not included. Molecular formulas can also describe molecules consisting of just on element like an oxygen molecule which is O2. The molecular formula does not, however, show you the structure of the molecule. It only shows you how many atoms of each element are in it. ** 8.2 The Nature of Covalent Bonding ** // Marion Burdick Pages 217-229 (All pictures and text) //

** The Octet Rule in Covalent Bonding ** **Key Concept:** In forming covalent bonds, electron sharing tends to occur so that atoms, with shared electrons included, attain the configurations of noble gases. -Combinations of nonmetallic groups (Groups 4A, 5A, 6A, and 7A) are likely to form covalent bonds in order to acquire a total of eight electrons

** Single Covalent Bonds ** **//Single covalent bond://** two atoms held together by sharing a pair of electrons are joined by this Ex. Hydrogen gas (H2)

**Key Concept:** An electrons dot formula such as H: H represents the shared pair of electrons of the covalent bond by two dots. **//Structural Formula://** a chemical formula that shows the arrangement of atoms in a molecule or polyatomic ion; each dash between a pair of atoms indicates a pair of shared electrons Ex. H-H **//Molecular Formula://** indicates number of that particular atom in the molecule Ex. H2 **//Unshared Pair://** A pair of valence electrons that is not shared between atoms (also called a //lone pair// or a //nonbonding pair//)





** Double and Triple Covalent Bonds ** **Key Concept:** Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons. //**Double covalent bond:**// a bond that involves two shared pairs of electrons //**Triple covalent bond:**// A bond formed by sharing three pairs of electrons -Oxygen (O2) does **not** obey the octet rule and the Lewis dot structure that one would predict

** Coordinate Covalent Bonds ** <span style="display: block; font-family: 'Times New Roman'; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: left;">**// Coordinate Covalent Bond: //** a covalent bond in which one atom contributes both bonding electrons <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: left;"> - Show coordinate covalent bonding in which with an arrow ( à ) pointing from the atom donating the pair of electrons to the atom receiving them <span style="display: block; font-family: 'Times New Roman'; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: left;">** Key Concept: ** In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms.



**//Polyatomic ion://** A tightly bound group of atoms that has a positive or negative charge and behaves as a unit Ex. NH4 + -The negative charge of a polyatomic ion shows the number of electrons in addition to the valence electrons of the atoms present

** Bond Dissociation Energies ** **// Bond Dissociation Energy: //** The total energy required to break the bond between two covalently bonded atoms (usually expressed in kJ/mol) ** Key Concept: ** A large bond dissociation energy corresponds to a strong covalent bond.

** Resonance ** -There is more than one way to express the electron dot structure of some molecules… WHY? Because earlier scientists believed that electron pairs rapidly flip back and forth (resonate) between different electron dot structures, they use double headed arrows to indicate that two or more structures are in resonance **Key Concept:** Actual bonding of oxygen atoms in ozone is hybrid, or mixture, of the extremes represented by the resonance forms. **//Resonance Structure://** A structure that occurs when it is possible to write two or more valid electron dot formulas that have the same number of electron pairs for a molecule or ion ** Exception to Octet Rule ** **Key Concept:** 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. -There are odd numbered electrons that disobey the octet rule Ex. NO2 (17 valence electrons- it is impossible to write a Lewis dot structure for NO2 that satisfies the octet rule for all atoms) Additional examples: (chlorine dioxide) ClO2 and nitric oxide (NO) -There are some even numbered electrons that also disobey the octet rule also Ex. (BF3)

**__ 8.3 Bonding Theories __ -pgs. 230-236 ** ** Eileen Corkery (Includes all of the following text, pictures, and videos for 8.3) **


 * __ Key Concepts: __**
 * How are atomic and molecular orbitals related?
 * How does VSEPR theory help predict the shapes of the molecules?
 * In what ways is orbital hybridization useful in describing molecules?

-//Molecular orbital//- an orbital that applies to the entire molecule
 * __ I) Molecular Orbitals __**

-Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole.

- An atomic orbital is filled once it contains two electrons; a molecular orbital that can be occupied by two electrons of a covalent bond is called a bonding orbital.

__-Sigma Bonds__
 * A sigma bond is formed when two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei.
 * The Greek symbol for "sigma" is σ
 * Covalent bonding results from an imbalance between the attractions and repulsions of the nuclei and electrons involved (The nuclei and electrons attract each other because of their opposite charges.)
 * In the above diagram, the two s atomic orbitals are squished together to make a molecular (bonding) orbital. In this bonding orbital, the electron density between the two nuclei is high.



- __Pi Bonds__
 * In the above diagram (the first row only!), the two p atomic orbitals combine to form a sigma-bonding molecular orbital.
 * Pi bonds are covalent bonds in which the bonding electrons are most likely to be found in sausage shaped regions above and below the bonds axis of bonded atoms
 * The symbol for a pi bond is π
 * Unlike sigma bonds, pi bonds line up side to side, not end to end.
 * In the above diagram, the side-by-side overlap of the two p atomic orbitals produces a pi-bonding molecular orbital.

media type="youtube" key="ree49ge4VA4?fs=1" height="385" width="480" align="center"

- The above video has a really good summary of both sigma and pi bonds (I think the visual really helps). In the second half of the video, she starts to explain hybridization, which you will read about right below this section.

-Just as a photograph does not show a structure's 3D composition, an electron dot structure does not show a molecule's true 3D formation
 * __ II) VSEPR Theory __**

-VSEPR stands for "valance-shell electron-pair repulsion theory". According to the VSEPR theory, the repulsion between electron pairs cause molecular shapes to adjust so that the valance electron pairs stay as far apart as possible.

-Above is the molecular formula, electron dot structure, and structural formula for methane. Methane molecules are really 3 dimensional. The 4 hydrogens in a methane structural formula or electron dot structure really represent the 4 corners of a geometric solid called a tetrahedron. It's called a tetrahedron because each angle is 109.5 degrees, the measure of a tetrahedral angle.

-Bonding angles are affected when one or more of the valance electron pairs is unshared. Because they are unshared, no bonding atom is pulling them. They are held closer together than the bonding pairs, affecting the bonding angle so it is NOT 109.5 degrees.

-Look in your book on the top of pg. 233 for examples of how water and carbon dioxide each bond, even though they each have pairs of unshared electrons. -The above chart shows some common molecular shapes, and the measures of their bond angles. Also, in the far right column is listed several compounds that take the shape of those particular molecule forms. (This is like the lab we did on 1/7/11)

-Orbital hybridization provides information about both molecular bonding and molecular shape.
 * __ III) Hybrid Orbitals __**

-//Hybridization// is the mixing of several atomic orbitals to form the same number of equivalent hybrid orbitals.

-Some examples would be the combination of 2s and 2p orbitals, or the overlap of 2sp^2 hybrid orbitals and a 1s orbital to form a sigma bond. It helps me to think of the compounds in 3D form in order to understand the concept.

-Overview of Hybridization: media type="youtube" key="g1fGXDRxS6k?fs=1" height="385" width="480" align="center"

-Look to pg. 234 in text book for diagram of single hybrid bonds in methane and pg 236 for an example of a triple bond in ethyne.

Group 3 (Pages 237-266)
Nina DeMeo (co-editor), Liz Sieber, Courtney Gareau, Evan Grandfield

<!--[if gte mso 10]> Key Concept: The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. - With covalent bonds, electrons are shared between atoms - covalent bonds differ in terms of how the bonded atoms share the electrons which depends on the kind and number of atoms joined together which thus determine its molecular properties - Nonpolar covalent bond – a covalent bond in which the electrons are shared equally by two atoms - Polar covalent bond – a covalent bond between atoms in which the electrons are shared unequally - The higher the electronegativity value, the greater the ability of an atom to attract electrons to itself - The electronegativity difference between two atoms tells what bond is likely to form
 * __Bond Polarity__** – Nina DeMeo
 * as electronegativity difference increases, the polarity increases



__**Polar Molecules**__ - Elizabeth Sieber - the presence of a polar ond in a molecule often makes the entire molecule polar. - polar molecule: a molecule in which one side of the molecule is slightly negative and the opposite side is slightly positive - dipole: (also known as a dipole molecule) a molecule that has two poles, or regions with opposite charges - <span class="key_concept">When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates. <span class="key_concept"> <span class="key_concept">Dipoles Moment

<span class="key_concept">__**Attractions Between Molecules**__ - Elizabeth Sieber <span class="key_concept">- Intermolecular attractions are weaker than either ionic or covalent bonds <span class="key_concept">- van der Waas forces: <span class="key_concept"> the two types of weakest attraction between molecules <span class="key_concept">- hydrogen bonds: <span class="key_concept"> attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom __**Intermolecular Attractions and Molecular Properties**__ - Elizabeth Sieber - the state (liquid, solid, or gas) that a compound is at room temperature depends on the type of bonding (ionic or covalent) - melting and boiling points of covalently bonded molecules are usually lower than those of ionic bonds - network solids (a.k.a. network crystals) are harder to break apart to make into another state; they are the exception to a low melting point for covalent bonds <-- Network Solid - the Chart below shows the differences in properties of ionic and covalent compounds
 * <span class="key_concept">named after French Chemist, Johannes van der Waals
 * <span class="key_concept">two types are dipole interaction and dispersion forces
 * dipole interactions : occur when polar molecules are attracted to one another
 * the positive side is attracted to another's negative and vise versa, making them stick together
 * dispersion forces: <span class="key_concept"> the weakest of all molecular interaction; are caused by the motion of electrons
 * <span class="key_concept">can occur between polar or non-polar molecules
 * <span class="key_concept">electrons puss against each other
 * <span class="key_concept">the more electrons, the stronger the dispersion force (generally)
 * <span class="key_concept">always involves hydrogen and an already bonded atom
 * <span class="key_concept">seen in drops of water
 * network solid: solid in which all of the atoms are covalently bonded to each other
 * the melting of a network solid would require the breaking of all these covalent bonds
 * ex. diamonds, silicon carbide (they are like single molecules)


 * Characteristic || Ionic Compound || Covalent Compound ||
 * representative unit || formula unit || molecule ||
 * bond formation || transfer of one of 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 C) || Low (usually below 300 C) ||
 * Solubility in water || usually high || high to low ||
 * electrical conductivity of aqueous solution || good conductor || poor to nonconducting ||

=9.1 Naming 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 - they can be either cation (+) or anions (-) __**Cations**__ - Elizabeth Sieber - metallic elements tend to lose valence electrons - group 1A has a +1 charge (ex. Lithium and Sodium), group 2A has a +2 charge (ex. Magnesium and Calcium), group 3A had a +3 charge (ex. Aluminum) - When the metals in Groups 1A, 2A,and 3A lose electrons, they form cations with positive charges equal to their group number - when naming them, follow the element name with the word "ion" (ex. Al+ is aluminum ion)
 * __Monatomic Ions__** - Elizabeth Sieber

<span class="key_concept">** __ Anions – Courtney Gareau __ ** <span class="key_concept"> Nonmetals gain electrons to form anions, so the charge of a nonmetallic ion is negative. <span class="key_concept"> **Key Concept –** The charge of any ion of a Group A nonmetals is determined by subtracting 8 from the group number. <span class="key_concept"> Example – Elements in Group 7A form anions with a 1- charge (7 -8 = -1) <span class="key_concept"> Anion names start with the stem of the element name and end in //–ide//.
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Ionic Charges of Representative Elements ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">1A || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">2A || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">3A || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">4A || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">5A || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">6A || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">7A || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">8A ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Li+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Be2+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">N3- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">O2- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">F- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Na+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Mg2+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Al3+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">P3- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">S2- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Cl- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">K+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Ca2+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">As3- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Se2- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Br- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Rb+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Sr2+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">I- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- ||
 * <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Cs+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">Ba2+ || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- || <span style="display: block; margin-bottom: 0pt; margin-left: 0in; margin-right: 0in; margin-top: 0in; text-align: center;">- ||

Ionic Charges of Representative Elements Table – Courtney Gareau

**Key Concept** – The charges of the cations of many transition metal ions must be determined from the number of electrons lost. The metals of Groups 1A, 2A, and 3A consistently form cations with charges of 1+, 2+, and 3+ respectively. Many of the transition metals form more than one cation with different ionic charges. Formulas and Names of Common Metal Ions with More than One Ionic Charge – Courtney Gareau The preferred method that is used to name these ions is the Stock System. This uses a Roman numeral in parentheses, which is placed after the name of the element to indicate the numerical value of the charge. No space is left between the element name and the Roman numeral in parentheses. Example: Fe2+ is named iron(II). An older less useful method for naming the cations uses a Latin root word for the element and different suffixes at the end of the name. A few transition metals have only one ionic charge. The names of these do not have Roman numerals. The exceptions include cations that have a 1+ charge (Ag+), and cations with a 2+ charge (Cd+ and Zn+). Many transition metals form brightly colored compounds that are used in the making or paint. Chromium is used to make yellow, orange, red, and green paints. Cadmium compounds are used to make a range of colors from yellow to red and maroon. Prussian blue is composed of iron combined with carbon, hydrogen, and nitrogen. Cadmium Orange Paint Tube – Courtney Gareau
 * __ Ions in Transition Metals – Courtney Gareau __ **
 * __ Polyatomic Ions – Courtney Gareau __ **
 * Polyatomic ions – ** a tightly bound group of atoms that behaves as a unit and has a positive or negative charge
 * Key Concept – ** The names of most polyatomic anions end in //–ite// or //–ate//.

Ammonium ion – Courtney Gareau Sometimes the same two or three elements combine in different ratios to form different polyatomic ions. The charge on each polyatomic ion in a given pair is the same. The //–ite// ending indicates one less oxygen atom than the //–ate// ending. However, the ending does not tell the actual number of oxygen atoms in the ion. All anions with names ending in //–ite// or //–ate// contains oxygen. Hydrogen-containing polyatomic ions are part of many compounds that affect our daily lives.

Many compounds from the days of early chemistry were named whatever the discoverer wished, but it was not uncommon for the name to describe some property of the substance of its source. The French chemist Antoine-Laurent Lavoisier determined the composition of many compounds in his experiments to show how chemical compounds form. He realized that it was impossible to memorize all of the unrelated names of compounds, so he worked with other chemists to develop a systematic method for naming chemical compounds. Antoine-Laurent Lavoisier – Courtney Gareau
 * __ Binary Ionic Compounds – Courtney Gareau __ **

Pages 261-266 Evan Grandfield I. **Naming Binary Ionic Compounds** A. Binary compound-a compound composed of 2 elements; either molecular or ionic 1. In a binary ionic compound, you can write the name if you know the formula. a. The first step to doing this is verifying that the compound is composed of one monatomic metallic cation and one monatomic nonmetallic anion. b. To name any binary ionic compound, place the cation name first, followed by the anion name (ex. sodium bromide—NaBr). c. However in metals that form more than one monatomic ion, you need to be more specific. 2. For example, in CuO, you CANNOT call it simply “copper oxide”. a. Copper forms the monatomic ions Cu+ and Cu2+ (copper (I) ion and copper (II) ion). b. With this is mind, you must realize that in the formula there is only one copper per oxygen. The oxygen has a 2- charge, so you need a copper ion with a 2+ charge to balance the sum of the charges at 0. c. In a few moments, you can come up with the name for CuO as cooper (II) oxide. c opper (II) oxide-Evan Grandfield []

B. You can determine the charges of monatomic ions from the periodic table, but you must memorize the charges of polyatomic ions. 1. When writing the names of formulas, you must bear in mind the ratio of cation to anion (in CO2, 1:2) and realize that you must balance the charges that they add up to 0. 2. Ex. CuS. There is a formula unit of 1:1. We know sulfur has a -2 charge and copper (I) has a 1+ charge and copper (II) has a 2+ charge. 2+ + -2 = 0, so we can see that the name for this compound is copper (II) sulfide. II. **Writing Formulas for Binary Ionic Compounds** A. You can write the formula for any binary ionic compound so long as you know its name. 1. Write the symbol of the cation and then the anion. You then add whatever subscripts are needed to balance the charges. 2. In order to form a compound, the positive charge and the negative charge have to add up to 0. a. Ex. Potassium chloride. A Potassium ion has a 1+ change and a chloride ion has a 1- charge. Therefore, one of each in the compound would balance the compound’s charge at 0. The ratio of potassium cations to chloride anions is then 1:1. We write the formula as KCl. b. Ex. Calcium Bromide. Calcium ions have 2+ charges and bromide ions have 1- charges. The ratio of calcium cations to bromide anions must be 1:2 to balance the charge at 0. We write the formula as CaBr2. c. Ex. Iron (III) oxide. Iron (III) ions have 3+ charges and oxide ions have 2- charges. You have 2 ways to find out how to balance the charge. 1. Find the least common multiple of the charges (6). Multiply iron (III)’s 3 charge by 2 (two 3 charges) and you get 6, which balances with oxide’s 2 charge multiplied by 3(3 2 charges) (6). This way, both get 6 (temporarily ignore +/-). Use the multipliers as the subscripts and you have the formula: Fe2O3. Iron’s now 6+ sum charge balance oxide’s now 6- sum charge. 2. Use the crisscross method. Write each atomic symbol of both ions with the charges above them. Then take the numerical value of the charges and put them as the subscripts, switching them to the opposing ions.

-Evan Grandfield [] a. While this results in a balanced charge, you must make sure the ratio is the lowest whole number ratio. If it is not, reduce it like you would a fraction.

** Compounds with Polyatomic Ions ** I. **With Compounds with polyatomic ions, write the symbol for the cation followed by the formula for the polyatomic ion and balance the charges.** A. An –ate or an –ite ending on a compound shows that the compound contains a polyatomic ion that includes oxygen. 1. The polyatomic ion acts like one ion, so you follow the same procedure as usual to balance the equation and write the formula. 2. If more than one polyatomic ion is needed, you put it in parentheses and put the subscript outside of the parentheses, like so: Ca(NO3)2. The subscript usually associated with the polyatomic ion by itself is kept in the parentheses. 3. If there is not more than one polyatomic ion, do not use parentheses, like in the formula SrSO4.

II. **Naming Compounds with Polyatomic Ions** A. Guidelines for naming compounds with polyatomic ions when given the formulas: 1. Recognize the formula contains a polyatomic ion. 2. Look up unfamiliar polyatomic ions. 3. If it contains a metal cation, then the first shown symbol is the cation, so find the charge by looking on a periodic table. 4. Name the compound by stating the cation first and the anion second like the way you do in naming binary compounds. B. Some ionic compounds containing polyatomic ions do not contain a metal cation. 1. In the example (NH4)2C2O4, polyatomic ammonium acts as the cation. Ammonium has a charge of 1+ so the oxalate (C2O4) must have a 2- charge to balance it (remember, no parentheses so there is only one of this polyatomic ion). a. This compound is named ammonium oxalate.

media type="youtube" key="gUldch1n0J4?fs=1" height="385" width="480"-Evan Grandfield [|http://www.youtube.com/watch?v=gUldch1n0J4&feature=player_embedded#!]

** Group 4 (Pages 268-278) **

 * Adam Shanahan (co-editor), Mike McShane, Steve Denison, Andrew Sciotti **


 * Pgs 268-270 – Andrew Sciotti**

__Naming Binary Molecular Compounds__


 * Key concept: The prefix in the name of a binary molecular compound tells us how many atoms of each element are present in each molecule of the compound.**

Binary molecular compounds are composed of two nonmetals and are not ions (opposite of binary ionic compounds).

Because elements can combine in many different ways therefore you need ways of distinguishing between different compounds.

Ex: Co2 is non-poisonous and is not harmful, but CO is poisonous that can kill a person. So we need a way of distinguishing between these.

Prefixes are used to tell between the different amounts of atoms of each element are present in each molecule of a compound.


 * Picture by Andrew Sciotti - prefixes for naming nomenclature chemicals*

Ex: Mono- means there is one oxygen atom in CO, di- means there is two oxygen atoms in Co2. Because of this they are named carbon monoxide and carbon dioxide.


 * All binary molecular compounds en in -ide***

Steps to naming binary molecular compounds


 * 1) Confirm it is a binary molecular compound (compound of two nonmetals)
 * 2) Name the elements listed in the formula in order (use prefixes the amount of each element)
 * 3) Remove mono- from the name if all the first element only has one atom
 * 4) Add –ide to the end of the second element

__Writing Formulas for binary molecular compounds__ Key concept: Use the prefixes in the name to tell you the subscript of each element in the formula. Then write the correct symbols for the two elements with the appropriate subscripts.

Ex: Silicon carbide has no prefixes so silicon only has one atom and carbon only has one atom so it is (SiC). Di**nitrogen** tetr**oxide has the prefix di- before nitrogen and tetr- before oxide. So nitrogen has 2 atoms and oxygen has 4, therefore it is written N2O4.**


 * Pgs 271-273 – Mike McShane**
 * __ Naming and Writing Formulas for Acids and Bases __**

Naming Acids:
 * Acids are groups of ionic compounds with different, unique properties. An acid by definition is a compound that contains one or more hydrogen atoms and produces hydrogen ions when dissolved in water. **
 * Acids can be used for many different things. In the picture below, an acid was used to make this design by eating away at the glass. **

__Writing Formulas for Acids:__
 * [[image:http://www.pearsonsuccessnet.com/ebook/products/0-13-190443-4/che00359c05.gif width="292" height="186"]]Bluish Design On Window-Mike McShane **
 * Acids will always have the formula HnX where X is the anion. There are three rules for naming acids: **
 * 1. <span class="key_concept">If the name of the anion ends in the letters "-<span class="emphasis_italic">ide", the acid will begin with hydro. <span class="emphasis_bold"> The anion will get "-ic" added to the end of it and then will be followed by the word acid. Ex: In HCl(aq); H is there because of the acid formula, Cl is the anion (X) and the (aq) is there to show that it is an acid. Because Cl is "Chlor **__IDE__**", it will get the suffix "ic". Therefore, the acid is called <span class="emphasis_italic">hydro chlor<span class="emphasis_italic">ic acid .**
 * 2. <span class="key_concept">If the name of the anion ends in the letters "-<span class="emphasis_italic">ite", the acid will end in "-<span class="emphasis_italic">ous" , and will be followed by the word <span class="emphasis_italic">acid . Ex: In H2SO3(<span class="emphasis_italic">aq ); H2 is there because of the acid formula, SO3 is the anion (X) and the (aq) is there to show that it is an acid. Because SO3 is "Sulf**__ITE__**", it will get the suffix "ous". Therefore, the acid is called sulfur<span class="emphasis_italic">ous acid. **
 * <span class="key_concept">3. If the name of the anion ends in the letters "-<span class="emphasis_italic">ate", the acid will end in "-<span class="emphasis_italic">ic", and will be followed by the word <span class="emphasis_italic">acid . Ex: In HNO3(<span class="emphasis_italic">aq ); H is there because of the acid formula. NO3 is the anion (X) and the (aq) is there to show that it is an acid. Because NO3 is "Nitr**__ATE__**", it will get the suffix "-ic". Therefore, the acid is called Nitric Acid. ﻿ **
 * Danger Acid- Mike McShane **
 * Danger Acid- Mike McShane **
 * Danger Acid- Mike McShane **

__﻿__**If you know the rules for acid names, you know how to write the formula for acids when you only have the name. All you have to do is to do the same process in reverse. Ex: In Nitric Acid; The "-ic" tells you that the anion ended in "-ate". Using the acid formula HnX where X is the anion, the formula for the acid is H (because H is always needed in acid formulas) combined with the formula for nitrate (NO3) followed by (aq) (because aq is always needed in acid formulas).**
 * So the formula is HNO3(aq) **
 * **__Names and Formulas for Bases:__
 * **__Names and Formulas for Bases:__

__ ﻿ __** Another group of ionic compounds are bases. To write formulas for bases, all you need to do is to write the symbol for the cation, followed by the formula for the hydroxide ion. Then you need to make sure that you balance the charges and give the subscripted number to the correct letters. Ex: Aluminum hydroxide; Aluminum is Al and has a charge of +3. Hydroxide's formula is OH and has a charge of -1. The formula for aluminum hydroxide is Al(OH). But then you need to make sure that the charges are even. Therefore, you need to have three Hydroxide ions to balance off the Aluminum. So the formula is Al(OH)3. **

- The Laws of Definite and Multiple Proportions
The Law of Definite Proportions - Chemical formula tells you ratio of atoms in elements - Done using ratios of masses - The masses (of the atoms) do not change no matter how much of the element is made - Atoms combine in whole number rations and the atoms within the element are always within the same proportions

The Law of Multiple Proportions - 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 - Water and Hydrogen Peroxide are examples of this


 * SAMPLE PROBLEM**

Calculating Mass Ratios

1. **Analyze** -List the knowns and the unknown. Knowns •Compound A  2.41 g C and 3.22 g O •Compound B  6.71 g C and 17.9 g O

Unknown •Lowest whole number ratio of carbon per gram of oxygen in the two compounds  ? Apply the law of multiple proportions to the two compounds. For each compound, ﬁnd the grams of carbon that combine with 1.00 g of oxygen by dividing the mass of carbon by the mass of oxygen. Then ﬁnd the ratio of the masses of carbon in the two compounds by dividing the larger value by the smaller. Conﬁrm that the ratio is the lowest whole number ratio.

2. **Calculate** Solve for the unknown.

Compound A 2.41 g C / 3.22 g O = 0.748 g C / 1.00 g O

Compound B 6.71 g C / 17.9 g O = 0.375 g C / 1.00 g

Compare the masses of carbon per gram of oxygen in the compounds. The mass ratio of carbon per gram of oxygen in the two compounds is 2:1.

3. **Evaluate** Does the result make sense? The ratio is a low whole number ratio, as expected. For a given mass of oxygen, compound A contains twice the mass of carbon as compound B.

Pgs 276- 278 – Adam Shanahan**

__Practicing Skills: Naming Chemical Compounds__

Two basic skills will be discussed in this chapter. They are writing chemical formulas and naming chemical compounds. Follow the following chart to find the correct name for a compound when you know it's formula.

Image added by Adam Shanahan

When writing a chemical formula, use the following notes to help.
 * An //ide// ending generally indicates a binary compound.
 * An //ite// or //ate// ending means a polyatomic ion that includes oxygen in the formula.
 * Prefixes in a name generally indicate that the compound is molecular.
 * A Roman numeral after the name of a cation shows the ionic charge of the cation.