Chapter 5: Molecules and Compounds

The great diversity of substances that we find it in nature is a direct result of the ability of elements to form compounds. Life, for example, could not exist with just 91 elements. It takes compounds, in all their diversity, to make life possible.

• For example: Hydrogen is an explosive gas, oxygen is not but must be present for burning to occur. Both have low boiling points but when the two are combined, a dramatically different substance results (H20). Water is nothing like the hydrogen and oxygen from which it forms.  
• When two or more elements combine to form a compound, an entirely new substance results.
• A compound is different from a mixture of elements. In a compound, elements combine in a fixed, definite proportions; in a mixture, elements can mix in any proportions whatsoever. 
• Compound: Water molecules have a fixed ratio of hydrogen (two atoms) to oxygen (one atom). (ex. Water H2O)
• Mixture: This can have any ratio of hydrogen to oxygen. 
• Homologous Mixture: has a uniform appearance and composition throughout. Commonly referred to as a solution. (ex. flour and baking soda)
• Heterogeneous Mixture: Consists of visibly different substances or phases: gas, liquid, and solid. (ex: trail mix)

5.2 Types of Chemical Bonds

• Chemical Bond: Is the force that holds atoms together in a compound. Chemical bonds form because they lower the potential energy of the charged particles that compose atoms, 
• Because the rest of the elements do not possess the stability of the noble gases, they form chemical bonds to become more stable (to lower the potential energy of the noble gases, they form chemical bonds to become more stable. 
• When two atoms approach each other, the electrons of one atom are attracted to the nucleus of the other according to Coulomb's law and vice versa. 
• Ionic Bond: A bond between a metal (-) and a non metal (+). Transfers atoms. 
• Metals have a tendency to lose electrons, and that non metals have a tendency to gain them. 
• Therefore, when a metal interacts with a non metal, it can transfer one or more of its electrons to the non metal.
• Metals become cations (positively charged)
• Nonmetals became anions (negatively charged)
• These oppositely charged ions attract one another according to Coulomb's law and form an Ionic Compound.
• Ionic Compound: Which in the sold state is composed of a lattice-a regular three dimensional array of alternating cations and anions. 
• The basic unit of an ionic compound is the formula unit, the smallest, electrically neutral collection of ions. 
• The formation of an ionic compound from its constituent elements usually gives off a bit of energy as heat (the process is exothermic-absorbs energy).
• A reaction is exothermic because of the lattice energy.
• Lattice Energy: the energy associated with the formation of a crystalline lattice of alternating cations and anions from the gaseous ions. That energy is emitted as heat when the lattice forms. 
• Ionic Compounds: Formulas and Names (5.6)

• Nomenclatures: The process of naming compounds.
• An ionic compound always involves:
• Positive and negative ions
• The sum of cations (+) must equal the sum of the charges of the anions (-)
• Reflects the smallest whole number ratio of ions.
• Examples on pg. 155
• Binary Compound: Contains only two different elements
• Name of cation (metal) + Base name of anion (nonmetal) + -ide
• Transition metals
• Binary with a metal that forms more than one kind of cations:
• For these types of metals, the name of the cation is followed by a roman numeral (in parenthesis), which indicates the charge of the metal in that particular compound. 
• [Name of Cation/metal](charge of cation/metal in roman numerals in parenthesis)[base name of anion/non metal)+-ide]

                                           

                                            

• Polyatomic Ion: An ion composed of two or more atoms.

• Polyatomic Ions: is a charged particle which has two or more atoms held together by covalent (sharing of pairs of electrons) bonds Some rules: 1. Ions that end in ate have oxygen in them. 2. Elements in the same family make similar ions. Common Polyatomic Ions
• Oxyanions- Most common polyatomic ion; anions containing oxygen and another element. Notice that when a series of oxyanions contains different numbers of oxygen, we name them systematically according to the number of oxygen atoms in the ion. If there are only two ions in the series, the one with more oxygen atoms has ending -ate and the one with fewer oxygen atoms has the ending -ite. 
• NO2- Nitrite (less than)
• NO3- Nitrate (more than)
• Hydrates- Contain a specific number of water molecules associated with each formula unit.  

1/2 Hemi

• Covalent Bond: The bond that forms between two or more nonmetals. Shares atoms.
• Nonmetals tend to have high ionization energies for their electrons are relatively difficult to remove. Therefore, when a nonmetal bonds with another nonmetals, neither atom transfers electrons to the other. 
• Nonmetals share atoms. The shared electrons interact with the nuclei of both of the bonding atoms, lowering their potential energy in accordance with Coulomb's Law. 
• Covalently bonded atoms form molecules, and the resulting compounds form molecular compounds. 
• Bonding Pair: A shared pair of electrons
• Lone Pair: A pair that is associated with only one atom and therefor not involved in bonding. Also called non-bonding electrons.  
• Double Bond: When two atoms share two electron pairs. Shorter and stronger than single bonds. 
• Triple Bond: The strongest and shortest bond. 
• [Prefix][Name of 1st Element][Prefix][Base name of 2nd element + -ide]

5.9 Formula Mass and the Mole Concept for Compounds

• The term molecular mass or molecular weight are synonymous with formula mass and are also common
• Formula Mass= (Number of atoms of 1st elements in chemical formula X Atomic mass of 1st element) +                                 (Number of atoms of 2nd element in chemical formula X Atomic mass of 2nd element)

• Ionic Bonding	Covalent Bonding
  Metal + Nonmetal	Nonmetal + Nonmetal
  Non directional and hold together an array of ions	Highly directional
  	Single, Double, Triple Bond
  	Individual molecules with interactions between them

Ionic Bonding

Covalent Bonding

Metal + Nonmetal

Nonmetal + Nonmetal

Non directional and hold together an array of ions

Highly directional

Single, Double, Triple Bond

Individual molecules with interactions between them

5.3 Representing Compounds: Chemical Formulas and Molecular Models

• The quickest and easiest way to represent a compound is with its chemical formula, which indicates the elements present in the compound and the relative number of atoms or ions of each. 
• Chemical formulas normally list the more metallic (or more positively charged) element first, followed by the less metallic (or more negatively charged). 
• The type of formula we use depends on how much we know about the compound and how much we want to communicate. 
• Empirical Formula: Gives the relative number of atoms of each element in a compound. (Ex.: Hydrogen Peroxide- HO)
• Communicated the least amount of information.
• Molecular Formula: Gives the actual number of atoms of each element in a molecule of a compound. (Ex.: Hydrogen Peroxide- H2O2)
• Is always a whole-number multiple of the empirical formula. 
• For some compounds, the empirical formula and the molecular formula are identical. (Ex.: H2O because water molecules contain two hydrogen atoms and one oxygen atoms, and no simpler whole number ratio can express the relative number of hydrogen to oxygen).
• Structural Formula: Use lines to represent covalent bonds and shows how atoms in a molecule are connected or bonded to each other. (Ex. H-O-O-H)
• Conveys the most amount of information. 
• Structural formulas may also be written to give a sense of the molecule's geometry. 
• This version represents the approximate angles between bonds, giving a sense of the molecule''s shape. 
• Structural Formulas can also depict the different types of bonds that occur within molecules. (Ex.: Carbon Dioxide O=C=O)
• Single Bond: (-Single Line), corresponds to one shared electron pair
• Double Bond: (=) Generally stronger and shorter than a single bond. Corresponds to two shared electron pairs. 
• Molecular Models: a more accurate and complete way to specify a compound. The structural formula indicates how the atoms are connected.
• Ball and Stick Molecular Model: represents atoms as balls and chemical bonds as sticks, how the two connect reflects a molecule's shape. 
• Space-filling molecular model: atoms fill the space between each other to more closely represent our best estimates for how a molecule might appear if scaled to visible size. 

                        

5.4 The Lewis Model: Representing Valence Electrons with Dots 

• Bonding theories (or models) are central to chemistry because they explain how atoms bond together to form molecules. They explain why some combinations of atoms are stable and others are not. 
• Lewis Electron Dot Structures: A model where valence electrons are represented as dots to depict molecules. These structures are fairly simple to draw, have tremendous predictive power as to if a set of atoms will form a stable molecule and what that molecule might look like. The Lewis Model remains the simplest model for making quick, everyday predictions about molecules. 
• Lewis model focuses on valence electrons because chemical bonding involves the transfer or sharing of valence electrons between two or more atoms. 
• Atoms with 8 valence electrons are stable because they have a full outer principal level-are easily identified because they have 8 dots-an octet. 
• Helium only contains 2 dots, a duet, which represents a stable electron configuration because n=1 quantum level fills with only two electrons. 
• Octet Rule-The bonding atoms obtain stable electron configurations; since the stable configuration is usually 8 electrons in the outermost shell.