Nomenclature+Chp.+7,+8,+9+(period+A)

= __Inorganic Nomenclature Chapters 7, 8 and 9__ =

Editor: Abby White
Chapter 7 is about Ionic and Metallic bonding. An ionic bond is the electrostatic force that holds the ions together in ionic compounds. Metallic bonds are the attraction of the valence electrons for the positively charged metal ions that are free floating Some of the major points are Ions, Ionic Bonds and Ionic Compounds, and Bonding in metals. Chapter 8 talks about Covalent bonding. A single covalent bond is when two atoms join due to the fact there is an electron pair to hold them. Molecular compounds, the nature of covalent bonding, the bonding theories and the polar bonds and molecules are discussed throughout the chapter. Chapter 9 focusses on chemical names and formulas. There are monatomic and polyatomic ions. Monatomic is when they consist of a single atom that either has a positive or negative charge. The polyatomic ion is composed of more than one atom. Naming Ions, Naming and writing formulas for ionic compounds, molecular bases, and acids/bases and the laws governing the formulas and names are explained in further detail.These 3 chapters focus on nomenclature, which is a system of terms and names by a specific community.

Chapter 7: Ionic and Metallic Bonding

__**Valence Electrons**__
__**The Octet Rule**__ __**Formation of Cations**__
 * Cations of Group 1A elements always have a charge of 1 +
 * Cations of Group 2A elements always have a charge of 2+
 * Atoms lose enough electrons to attain the electron configuration of the noble gas
 * Example: all Group 2A elements have two valence electrons and in losing two electrons, they form 2+ cations
 * Some ions formed by transition metals do not have noble-gas electron configurations (ns2np6) and are exceptions to the octet rule
 * Example: silver
 * To achieve the structure of krypton, a silver atom would have to lose 11 electrons which is impossible or highly unlikely
 * Other atoms such as copper can ionize to form a 1+ cation (Cu+)
 * [[image:mrdschemistryhwiki/T6c42d_1.jpg width="309" height="333"]]
 * This is an example of certain common ions that form multiple cations such as copper.

**__Formation of Anions__**

 * An anion is an atom or 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
 * Atoms of nonmetallic elements attain noble gas electron configurations easier by gaining electrons than losing them because of their full valence shells
 * Example: chlorine is in Group 7A and has 7 valence electrons.
 * A gain of 1 electron gives chlorine an octet and converts chlorine into a chloride ion.
 * Chloride ion is an anion with a single negative charge and has the same electron configuration as the noble gas argon
 * Use of electron dot structures help differentiate from the formation of a chloride ion from a chlorine atom.
 * Each dot represents an electron in the valence shell
 * The ions produced when atoms of chlorine and other halogens gain electrons are called __halide ions__
 * All halogen atoms have seven valence electrons and need to gain only one electron to achieve the electron configuration of a noble gas
 * [[image:mrdschemistryhwiki/Wiki.jpg width="560" height="317" caption="This is a chart of common anions and cations and their poly atomic symbols"]]

**__Writing the Symbols and Names of Ions__**
1. Analyze the relevant concepts a. An atom that gains electrons forms a negatively charged ion(anion). The name of an anion of a nonmetallic element ends in -ide. b. An atom that loses electrons forms a positively charged ion (cation). The name of a cation of a metallic element is the same as the name of the element. 2. Solve: Apply concepts to this situation

__Formation of Ionic Compounds__
__**Ionic Bonds**__ Properties of Ionic Compounds Coordination number Conductivity
 * **Ionic compounds**are compounds composed of cations and anions
 * usually composed of metal cations and nonmetal anions
 * sodium chloride is composed of sodium cations and chloride anions
 * **Ionic bonds**are electrostatic forces that hold ions together in ionic compounds
 * Sodium chloride provides a simple example of how ionic bonds are formed
 * Sodium has a single valence electron that it can easily lose
 * Chlorine has seven valence electrons and can easily gain one
 * When sodium and chlorine react to form a compound, the sodium atom gives its one valence electron to a chlorine atom
 * The two atoms combine in a one-to-one ratio and each has a stable octet
 * __Formula Units__**
 * Chemists represent the composition of substances by writing chemical formulas
 * **Chemical formulas** show the kinds and numbers of atoms in the smallest representative unit of a substance
 * NaCl is the chemical formula for sodium chloride, however it doesn't represent a single discreet unit.
 * a **Formula unit**is the lowest whole-number ratio of ions in an ionic compound.
 * for sodium chloride, this is 1:1 (one Na+ to each Cl-)
 * the number of ions of opposite charge that surround the ion in a crystal
 * this is sodium cloride
 * because each Na+ ion (light blue) is surrounded by 6 Cl- ions (dark green) the coordination number is 6 [[image:http://upload.wikimedia.org/wikipedia/commons/e/eb/Sodium_chloride_crystal.png width="236" height="178"]]
 * ionic compounds can conduct an electric current when melted or dissolved in water
 * if voltage is added to a molten mass, cations move freely to one electrode, and anions move to the other
 * this movement allows electricity to flow between electrodes to move through an external wire

By: Ali Fortier and Nikki Steiner
>
 * Metallic Bonds and Metallic Properties**
 * Metals are made up of closely packed cations
 * Valence electrons of metal atoms can be modeled as a sea of electrons
 * Valence electrons are mobile and can drift freely to different parts of the metal
 * Metallic bonds are made up of the attraction of free-floating valence electrons
 * These bonds hold metals together
 * Sea of electrons model: [[image:http://image.tutorvista.com/content/chemical-bonding/electron-sea-model.gif width="287" height="197" align="middle" caption=" sea of electrons model"]]
 * Explains physical properties of metals (good conductors of heat; ductile; as electrons enter one end of a bar of metal, an equal number leave other end; malleable)
 * Ductility and malleability explained in terms of mobility of valence electrons


 * Crystalline Structure of Metals**
 * Metals are crystalline
 * Simplest form of crystalline solids = metals that contain one kind of atom
 * Metals are arranged in compact and orderly patterns
 * Several closely packed arrangements are possible
 * Body-centered cubic; face-centered cubic; hexagonal close-packed
 * Body-centered cubic arrangement
 * Every atom besides those on surface has 8 neighbors


 * Face-centered cubic arrangement
 * Every atom besides those on surface has 12 neighbors


 * Hexagonal close-packed arrangement
 * Every atom besides those on surface has 12 neighbors
 * Pattern is different from face-centered though

Alloys: - very few objects are pure metals, most are alloys - alloys: mixtures composed of two or more elements, at least one of which is a metal - they are important because their properties are often superior to those of their component elements - the most important ones today are steels - they have a wide range of properties (corrosion resistance, ductility, hardness, and toughness) - form from their component atoms in different ways - can replace each other in a crystal (if the atoms of the components in the alloy are about the same size) - called a substitutional alloy - the smaller atoms can fit into the interstices (spaces) between the larger atoms (if the atomic sizes are different) - called an interstitial alloy (ex: steel)



By: Kyle Gallagher
**Molecules and Molecular Compounds**
 * Monoatomic atoms consist of single atoms
 * Include noble gas elements, such as neon and helium
 * Different compounds have different properties
 * Salts like sodium chloride (NaCl) have different properties than hydrogen chloride (HCl) and Water (H20)
 * **Covalent bond** – “tug of war” for the electrons between the atoms bonds them together, sharing them (unlike ionic bonding where electrons are given up or accepted)
 * **Molecule** – a neural group of atoms joined together by covalent bonds
 * **Diatomic Molecule**– a molecule consisting of two atoms
 * Such as an oxygen molecule (O2)
 * Atoms of different elements can combine chemically to form compounds
 * **Molecular Compound** – a 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 molecular compounds are gases or liquids at room temperature and are composed of atoms of two or more nonmetals
 * Ionic compounds are usually formed from a metal combined with a non-metal

**Molecular Formulas**
 * **Molecular Formula**– chemical formula of a molecular compound
 * such as H2O for water
 * A molecular formula shows how many atoms of each element a molecule contains
 * The subscript after each element symbol shows how many atoms of that element are in the molecule
 * A molecular formula reflects the actual number of atoms in each molecule
 * The subscripts are therefore not necessarily in the lowest whole number ratios
 * Like ethane, C2H6
 * Molecular formulas also describe molecules consisting of one element, like O2
 * A molecular formula does not tell the molecule’s structure
 * Does not show the arrangement of various atoms in space or which atoms are covalently bonded to one another
 * The ways of showing the structure can be seen in the picture below which uses Ammonia as an example
 * Different molecules form different structures
 * Straight (like carbon dioxide), bent (like water), or more complex (like ethanol)



By: Brynna Harum and Andrea Vale
__The Octet Rule in Covalent Bonding__ __Single Covalent Bonds__ <-- METHANE
 * In forming covalent bonds, electron sharing usually occurs so that atoms attain the electron configuration of noble gases
 * atoms usually acquire a total of eight electrons, or an octet, by sharing electrons, so that the octet rule applies
 * a single covalent bond joins two atoms that are held together by sharing a pair of electrons
 * hydrogen gas consists of diatomic molecules whose atoms share only one pair of electrons, forming a single covalent bond
 * An **electron dot structure** such as **H:H** represents the shared pair of electrons of he covalent bond by two dots
 * a **Structural formula** represents the covalent bonds by dashes and shows the arragment of covalently bonding atoms (H-H)
 * Unshared pair is a piar of valence electrons that is not shared between atoms and are also known as a lone pair or a nonbonding pair
 * when carbon forms bonds wih other atoms it ussually forms four bonds
 * one of carbons 2s electrons is promoted to the vacant 2p orbital to form the following electron configuration- only requires a small amount of energy

__Double and Triple Covalent Bonds__
 * a **double bond** is a bond that involves two shared pairs of electrons
 * a **triple bond** is formed by sharing three pairs of electrons
 * **atoms form double or triple covalent bonds if they can attain a noble gas structure**
 * atoms can also form triatomic molecules which is a molecule consisting of three atoms. An example of this is the Carbon dioxide molecule

__Coordinate Covalent Bonds__


 * Coordinate Covalent Bond: A covalent bond in which one atom contributes both bonding electrons

-- 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

-- Compounds containing polyatomic ions include both ionic and covalent bonding

-- A negatively charged polyatomic ion is part of an ionic compound, so the positive charge of the cation of the compound balances these additional electrons

__Bond Dissociation Energies__
 * Bond Dissociation Energy: The energy required to break the bond between two covalently bonded atoms
 * A large bond dissociation energy corresponds to a strong covalent bond

__Resonance__ -- 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
 * Double covalent bonds are usually shorter than single bonds – however, the two bonds in ozone are the same length
 * The actual bonding of oxygen atoms in ozone is a hybrid, or mixture, of the extremes represented by the resonance forms
 * Resonance Structure: A structure that occurs when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion

__Exceptions to 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
 * In molecules that have an odd number of electrons, complete pairing of electrons is not possible (Chlorine dioxide and nitric oxide)
 * Some molecules that have an even number of electrons may not follow the octet rule because an atom requires less than an octet of eight electrons (Some compounds of Boron)
 * Some atoms expand the octet to include ten or twelve electrons (Phosphorus and sulfur)

By: Brittany Morgan
__ Molecular Orbitals __ __ Sigma Bonds __ __ Pi Bonds __ __ VSPER Theory __ __ Hybrid Orbitals __ __ Hybridization Involving Single Bonds __ __ Hybridization Involving Double Bonds __ __ Hybridization Involving Triple Bonds __
 * There is a quantum mechanical model of bonding that describes the electrons in molecules using orbitals that exist only for grouping of atoms
 * Molecular orbitals are when two atoms combine and their atomic orbitals overlap.
 * Atomic and molecular orbitals are similar but there are differences
 * Atomic orbitals belong to a particular atom; this is filled if it has two electrons
 * Molecular orbitals belong to a molecule as a whole; this is filled if it has two electrons
 * Bonding orbital is a molecular orbital that can be occupied by two electrons of a covalent bond
 * 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
 * Covalent bonding is formed from an imbalance between attractions and repulsions of the nuclei and electrons involved
 * Due to their opposing charges they attract each other
 * Atomic //p// orbitals can also overlap to form molecular orbitals
 * [[image:mrdschemistryhwiki/pi.gif]]
 * The side-by-side overlapping produces pi molecular orbitals
 * A pi bond occurs when a pi molecular orbital is filled with two electrons
 * In a pi bond the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis or the bonded atoms
 * Atomic orbitals in pi bonding overlap less than sigma bonding
 * In result pi bonds tend to be weaker
 * [[image:mrdschemistryhwiki/sigma.gif]]
 * The valence-shell electron-pair repulsion theory explains three dimensional shapes of molecules.
 * According to VSPER theory the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible
 * If there is an unshared electron they are held closer to the element, verses if they are shared they are held at a further distance because they are vying for each other.
 * There are many different types of shapes including but not limited to: Linear triatomic, Trigonal planar, Bent triatomic, Pyramidal, Tetrahedral and Square planar.
 * Orbital hybridization provides information about both molecular bonding and molecular shape
 * In hybridization several atomic orbitals mix to form the same total number of equivalent hybrid orbitals
 * All bonds are identical, and orbital hybridization explains this
 * Hybridization is also useful in describing a double covalent bond
 * Pi bonds and Sigma bonds are alike in structure but pi are weaker and easier to break
 * An example of this is in ethane the 2//s// and 2//p// atomic orbitals of carbon form with the four available 1//s// orbitals to form a carbon-carbon sigma-bonding orbital.
 * This is a third type of covalent bond
 * The hybrid orbital description is guided by the properties of the molecule
 * An example of triple bonding is in an ethyne molecule and a carbon-carbon-sigma-bonding molecule is formed from the overlap of one //sp// orbital from each carbon.

By: Ryan McSweeney
__ Bond Polarity __ Covalent bonds invlove electron sharing between atoms. However, covalent bonds differ in terms o fhow th ebonded atoms share the electrons. The character of the bonds in a given molecule depends on the kind and number of atoms joined together. These features determine the molecular properties. -Consider hydrogen chloride. Hydorgen has an electronegativity of 2.1 and chlorine's is 3.0. This is a significant difference so the covalent bond is polar. The lowercase delta symbol denotes atoms involved in the covalent bond. The minus sign shows that chlorine has acquired a slightly negative charge. The plus sign shows that hydrogen has acquired a slightly positive charge. The polar nature may also be represented by an arrow pointing to the most electronegative atom. The O-H bonds in water are also polar. The highly electronegative oxygen pulls bonding electrons away from hydrogen molecules. The oxygen acquires a slightly negative charge. The hydrogen is slightly positive.As shown in this picture.
 * Polar Bonds and Molecules **
 * Nonpolar Covalent Bond ** is when bonding electrons are shared equally.
 * Polar Covalent Bond ** is a covalent bond between atoms in which the electrons are shared unequally.
 * __ Key Fact __** the more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. The higher the electronegative value, the greater the ability of an atom to attract electrons to itself.



Remember, the greater the difference of electronegativity the higher polarity. __**Key Fact:**__ When polar molecules are placed between oppositely charged plates they tend to become oriented with respect to the positive and negative plates. A carbon dioxide molecule for example has two polar bonds, therefore it is linear. __ Van Der Waals Forces __ The two oweakest attractions between molecules are collectively called van der Waals forces. They consist of dipole interactions and dispersion forces. Hydrogen Bonds The dipole interactions in water produce an attraction between water molecules. Each O-H molecule is very polar. Rember that a hydrogen bond is an attraction to hydrogen atom already bonded to a strongly electronegative atom. Hydrogen bonds are the strongest in intermolecular forces. Intermolecular Attractions and Molecular Properties __**Key Fact:**__ Melting a network solid would require breaking covalent bonds throughout the solid. Diamonds are an example.
 * Polar Molecules **
 * Polar molecule** is one in which one end is slighlty positve, and the other is slightly negative.
 * Dipole** is a molecule with two poles.
 * Attractions Between Molecules **
 * __Key Fact:__** Intermolecular attractions are weaker than either ionic or covalent bonds. Which also determines whether a molecule is a liquid, solid, or gas.
 * Dipole interactions** occur when polar molecules are attracted to one another. The slightly negative region of a polar molecule is weakly attracted to the slightly positive region of another molecule.
 * Dispersion Forces** are the weakest of all molecular interactions and are caused by the motion of electrons. They occur even between nonpolar molecules. When th eelectrons happen to be momentarily more on the side of a molecule closest to a neighboring molecule, their electric force influences the neighboring molecules elctrons to be momentarily more on the opposite side. This causes a very small attraction between the two molecules.
 * Hydrogen Bonds** are 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. Hydorgen bonding always involves hydrogen.
 * Network Solids** are stable solids in whic hall of the atoms are covalently bonded together.

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. Metals from elements in Groups 1A, 2A and 3A lose electrons to form cations. Their names are simply the elemental name (Ex: Aluminum) followed by the word "ion" (Ex: Aluminum Cation).
 * 1) ====Cations====

2. Anions
Nonmetals from Groups 5A, 6A and 7A gain electrons to form anions. Their names always start stem with the stem of the elemental name and ends with "-ide" when they become anions. (Ex: Oxygen gains two electrons and becomes "oxide".)

Ions of Transition Metals

 * Groups 1B - 8B make up the transition metal elements. This group does not strictly follow principles we see in the representative element groups of 1A - 8A.
 * The charges of the cations created in this metal grouping are determined by the number of electrons lost when the element becomes an ion.
 * The general rule of thumb is that the elements that are transition metals are labeled as +2 elements meaning they lose two electrons. However, this rule does not apply 100% of the time.**



There are two main methods used to determine the ion:
 * 1) Stock System - a system using Roman Numerals to label the different ions. (Ex: Fe+2 or Iron (II) loses two electrons when becoming a cation whereas Fe+3 or Iron (III) loses three electrons when becoming a cation.)
 * 2) Suffix System - a method using prefixes and suffixes to only tell you which ion has the larger and smaller charge. The prefixes sometimes use latin root name for the element. (Ex: Fe, Iron, when written in this is named either the "Ferrous Ion" or the "Ferric Ion". "Ferr-" comes from its lain background. "-Ous" is used for the smaller or the ions and "-Ic" is used for the larger of the two ions. Ferrous ion would be Iron (II) while Ferric ion would b Iron (III).)

Here is a very simple explanation as to why atoms form ions: [|Simple Explanation on Ion Formation]

Here is a more in-depth explanation of ions: In-Depth Ion Explanation

Polyatomic Ions
Consists of more than one atom that are bound together, acting as a single unit and carrying a single charge. (Ex: Sulfate Ion consist of one Sulfur atom and four Oxygen atoms [SO4] and has a charge of -2.) The names of polyatomic ions generally end in either "-ite" or "-ate".

-Generally speaking, the polyatomic ion with less Oxygen atoms will end with "-ite" while the polyatomic ion ending with more Oxygen atoms will end with "-ate" (Ex: Nitrite = NO2. Nitrate = NO3.) -If the polyatomic ion has a positive charge it will end with "-ium". (Ex: Ammonium [NO4] has a +1 charge.) -The prefixes "Hypo-" or "Per-" are used when there are either three of four distinguishing ions that have the same formula with different Oxygen atom amounts. (Ex: ClO is Hypochlorite, ClO2 is Chlorite, ClO3 is Chlorate, and ClO4 is Perchlorate.) -All polyatomic ions containing Oxygen at the end the end will end with either "-ite" or "-ate". -All polyatomic ions not ending with oxygen will end with "-ide". (Ex: CN is Cyanide.)
 * Read the table on pg. 257 to see the majority of polyatomic ions we must know.**

When the formula of a polyatomic atom begins with Hydrogen, simply add the word to the beginning... Ex: (H) (Hydrogen) + (CO3) (Carbonate) = HCO3 Hydrogen Carbonate (H) (Hydrogen) + (HPO4) (Hydrogen Phosphate) = H2PO4 Dihydrogen Phosphate

CHeck out the link below to see a fun rap about polyatomic ions. Notice the rules explained above used when naming the ions. Polyatomic Ion Rap

By: Ally Luongo
French chemist Antione-Laurant Lavoisier knew it was hard to memorize the names of the compounds. He worked with other chemists on a naming system. Naming Binary Ionic Compounds binary compound- composed of 2 elements and can either be ionic or molecular - to name any binary ionic compund place the cation name first, followed by the anion name - the charges of the monatomic anions can be determined from the periodic table
 * Binary Ionic Compounds**

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 - the positive charge of the cation must balance the negative charge of the anion so that the net charge of the formula is zero - crisscross method: the number for the charge of the ion becomes the subscript for the other ion

- write the symbol for the cation followed by the formula for the polyatomic ion and balance the charges - put parenthesis around the polyatomic ion in the formula
 * Compounds with Polyatomic Ions**

Naming Compounds with Polyatomic Ions - first recognize that the compound contains a polyatomic ion - to name a compound containing a polyatomic ion, state the cation first and then the anion just as you did in naming a binary ionic compounds - some ionic compounds containing polyatomic ions do not finclude a metal cation - instead the cation may be the polyatomic ammonium ion (NH4+)

** By: Ian Kelly **
__Naming Binary Molecular Compounds__
 * Binary Compounds are composed of ions of 2 elements
 * Binary Molecular Compounds are also composed of 2 elements, but both elements are non-metals
 * This affects the naming of the compounds and the formulas
 * Composed of molecules not ions
 * When nonmetals combine, they can come together in many different way


 * Prefix: name of a binary compound that tells how many atoms of each element are present in each molecule of the compound
 * There are guidelines to name the binary molecular compounds
 * Confirm that it is actually a binary molecular compound
 * Name must identify the elements
 * Elements should be listed in the order they are given in the formula
 * Omit the prefixes
 * This should show the formula
 * 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

By: Catherine Murray
Naming Acids Writing Formulas for Acids Names and Formulas for bases
 * **acid:** a compound that contains one or more hydrogen atoms and produces hydrogen ions (H+) when dissolved in water
 * To name an acid, you must consider the acid to consist of an anion combined with as many hydrogen ions as are needed to make the molecule electrically nuetral
 * The chemical formula is usually HnX where X= a monatomic or polyatomic anion and n is the subscript which indicates the number of hydrogen atoms
 * There are 3 rules to name acids
 * 1) When 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.
 * 2) When the anion name ends in //-ite//, the acid name is the stem of the anion with the suffix //-ous//, followed by the word 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.
 * 1) 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.
 * Use the rules for writing acids in reverse to write the formulas for acids
 * Hydrobromic acids: the hydro prefix and ic suffix point to rule one. It must be a combination of hydrogen ion and brmide ion, so the formula is HBr.
 * Common Acids:**
 * Name || Formula ||
 * Hydrochloric Acid || HCl ||
 * Sulfuric Acid || H2SO4 ||
 * Nitric Acid || HNO3 ||
 * Acetic Acid || CH3COOH ||
 * **base**: an ionic compound that produces hydroxide ions when dissolved in water
 * named by the name of the cation followed by the names of the anion
 * To write the formula, write the symbol for the metal cation followed by the formula for the hydroxide ion then you must balance the charges.
 * Aluminum hydroxide consists of the aluminum cation (Al3+) and the hydroxide ion (OH-) and you need three hydroxide ions to balance the charge, so the formula is Al(OH)3.

By: Maddie Myers
The Law of Definite and Multiple Proportions
 * the rules for naming/writing formulas are only possible because compounds form in predictable ways
 * ways are summed up in 2 laws
 * __law of definite proportions:__in samples of any chemical compound, the masses of the elements are always in the same proportions
 * subscripts tell you the ratios of atoms of each element in a compound
 * mass ratio does not change no matter how the chemical is formed or the size of the sample
 * atoms combine in whole-number ratios, so their proportions by mass must always be the same
 * __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
 * stated by John Dalton
 * stated by John Dalton

Practicing Skills: Naming Chemical Compounds >
 * uses
 * there are 100s of chemicals in your house; drugs, cleaning products, pestacides
 * what if the chemical mixed together/ began to react/ one was ingested
 * call poison control
 * would be much easier if you could supply information about the name of formula of the substance
 * i suggest looking at the chart in the book p. 277, it is really helpful**
 * guidelines
 * -ide generally indicated a binary compound
 * -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
 * chart p.278 (the one we copied in class that starts with Name of Compound)