Final Exam Notes

Unit 1: Matter, Transformation, and Energy

• Chemistry: The study of properties of materials and the changes the materials undergo
• Matter- Anything that occupies space and has mass. We can classify it based on:
• The State of Matter: Solid, Liquid, or Gas
• Solids and liquids are referred to as the condensed phases. 
• Gases have different characteristics than the latter. 
• What state a material is in depends largely on two factors:
• The amount of kinetic energy the particles possess
• The strength of attraction between the particles, potential energy (inter molecular forces)
• A change in state is a physical change where the physical form changes but keeps its composition.
• Reversible by changing the temperature.
• Solids: 
• High potential energy and low kinetic energy. 
• The only motion is vibration of the atoms and molecules. 
• Definite shape and volume.
• Liquids:
• Attraction between atoms is still high but it has higher kinetic energy and lower potential energy
• definite volume but not definite shape for it takes up the shape of the container. 
• Vibration, rotation, and some translation
• Gas:
• There is little attraction between the atoms
• High kinetic energy, low potential energy
• Vibration, rotation, and lots of translation
• The Composition of Matter: 
• Matter is either a mixture or substance.
• Substance: form of matter with definite composition and distinct properties. 
• Elements: A pure substance that cannot be chemically converted into a simpler substance. Consists of Atoms or Elements. 
• Atoms: Any no diatomic elements; Atoms are the building blocks of matter and they combine to make molecules.
• Molecule: Any Diatomic element (elements that only occur in pairs in nature)
• Br2,  I2, N2, Cl2, H2, O2, F2
• A chemical combination of two or more atoms
• Compounds: elements can combine with other elements to form compounds
• Mixture: physical combination of two or more substances.
• homogeneous: uniform throughout
• heterogeneous: NOT uniform throughout
• Physical Change:
• A change that occurs without changing the molecular composition. The same element is present before and after.
• Ex. Water is the same no matter if it is liquid or solid (ice).
• Examples of physical properties:
• Mass, volume, color, solid, texture, smell, taste, ect.
• Extensive Properties: Properties that depend of the amount of a substance (mass, volume)
• Intensive Properties: Does not change with the amount of a material (density, specific heat, melting and boiling points)
• Chemical Change:
• The property's potential to undergo change to form a different product. 
• Results of a bonds breaking in the reactants and rearrangements of the atoms to different substances (s). 
• Combustible, corrosive, toxic, flammable
• indicated by temp., color, bubble and precipitate formation
• The Atomic Theory of Matter
• Define Laws of Chemistry:
1. Law of Mass Conservation: In a chemical reaction, matter is neither created not destroyed. 
• The combined masses of the reactant must equal combined masses of the product
2. Law of Definite Proportions/ Law of Definite Composition:
• All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements. 
3. Laws of Multiple Proportions (John Dalton):
• The ratio of the two masses is a small whole number
• John Dalton's Atomic Theory:
1. Each element is composed of elements
2. Each element is composed of only one type of atom
3. Compounds are formed when atoms of more than one element combine in simple, whole number rations (ex. H2O)
4. Atoms of an element are not changed into atoms of different element by chemical reactions
• John Dalton:
• Modern atomic theory

The Discovery of the Electron:

• JJ Thomson: Discovered the Electon 
• Cathode Ray Tubes: A high voltage passed through a partially evacuated glass tube, produced radiation

Robert Millikan Experiment: 

• Oil drop experiment
• Find out charge of electron (1.6X10^-19)

The Discovery of Atomic Nucleus:

• Ernest Rutherford: Gold Foil Experiment; discovered the mystery of how electrons fit inside the atom
• Used beams of particles as projectiles to explore the structure matter
• Discovered nucleus and it is contents: protons and neutrons.

• Atoms are composed of three subatomic particles: protons, neutrons, and electrons that effect the chemical behavior of an atom.

Protons and neutrons reside in the nucleus of the atom (held together by strong nuclear force) while the area surrounding the nucleus is a diffused region of negative charge where the electrons reside.

• An atom must have the correct ratio of protons and neutrons for stability. 
• Protons: The number of protons in the nucleus in the nucleus is called the atomic number (Z) and determines the identity of the atom 
• Neutrons: neutral
• Electrons: Same number as protons

PhaseChanges,Liquids,&Solutions ‐Chapters10,12,13,14 (Unit 8)

8.1.Review physical states of matter and relate Kinetic Molecular Theory[Reading12.2]
8.2.Characterize properties of liquids.[Readings12.4,12.5,12.6]
8.3.For Physical States of Matter: Identify name and process of phase changes including the energetics of physical changes of matter.[Readings12.1,12.5,12.6]
 8.4.Determine a change in heat energy for physical change in matter: Using a Heat Curve[Reading10.4,12.7] 
8.5. Interpret phase diagrams.[Readings13.1,13.2]
8.6.Differentiate Ionic, Network Covalent, Molecular, and Metallic Solids                       [Reading13.5,13.6,13.7]

Identifying Types of Solids

• Solid particles experience a vibrational motion
• Amorphous Solids: When crystalline solids have ordered, defined arrangements of atoms, ions, or molecules with irregular arrangements. 
• Amorphous solids are disordered-there is no orderly structure. As a result, intermolecular forces are irregular, which affects the properties (such as melting point) of the substance. 
• The arrangement of the particles in a crystalline solid is called the crystal lattice.
• The smallest unit that sows the pattern of arrangement for all particles is called the unit cell. 

Summary table 12/4 pg. 516

• What are the factors that affect the strength of the ionic bond?
• Magnitude of the charge and the size of the ion.
• Covalent bonds are stronger that intermolecular forces

8.7. Identify role of entropy and intermolecular forces in solution formation.[Readings14.2]
8.8.Characterize the energetics of solution formation.[Readings14.3]
8.9.Identify equilibrium in solutions and the effect of temperature and pressure.[Readings14.4]
8.10.Characterize colligative properties of solutions.[Readings14.6,14.7]

Chapter 11: Gases

Matter is paticulate and ts behavior can be understood in terms of particles

7.1 Describe measuring gas pressures using barometers and manometers. Convert between pressure units. (11.1 and 11.2) Explain and measure gas pressure. Relate pressure units. 

• The behavior of gases can be explained and predicted by a model called the kinetic molecular theory. The core of this model is that gases are composed of particles in constant motion.
• 1.) Observations on the properties of gases
• 2.) Observations led to scientific laws (that summarized the observed behavior)
• 3.) The observations and laws led to the development of the kinetic molecular theory-the model the nature of a gas. 
• Observations---> Scientific Laws ---> Kinetic Molecular Theory
• Gases:
• Have no definite shape and no definite volume for gases will expand to fill any container. 
• Characteristics of Gases:
• Compressibility: Change in volume of a sample resulting from a pressure change acting on the sample. 
• Solids and gases are not significantly compressible but gases can be compressed under sufficient pressure. 
• Gas molecules are widely spread and therefore have low densities. 
• Pressure: Particles in a gas collide with each and with the surfaces around them. Each collisions exerts only a small force, but when the forces of many particles are summed, they quickly add up. 
• The result of the constant collision between the atoms and molecules in a gas and the surfaces around them is pressure. 
• The pressure that a gas sample exerts is the force that results from the collisions of gas particles divides by the area of the surface with which they collide.
• Pressure= (Force/ Area) 
• P= (F/A)
• Pressure is the force (F) that acts on a given area (A)
• Atmospheric Pressure: The pressure that is exerted upon the earth's surface by the atmosphere. 
• Variation in pressure in Earth's atmosphere create wind, and changes in pressure help us to predict weather. 
• The number of gas particles in a given volume decreases with increasing altitude (See Below)*
• Pressure decreases with increasing altitude
• Pressure Observations:
• The total pressure exerted by a gas depends on several factors such as:
• Concentration: Number of gas particles in a given volume. The lower the concentration (fewer gas particles), the lower the pressure. Likewise, a high density of gas partiles results in high pressure. 
• Altitude *: Because the number of gas particles in a given volume decreases with increasing altitude, pressure decreases with increasing altitude. 
• Ex. In the air, airplane cabins have to be artificially pressurized due to lack of oxygen at high altitudes. 
• Ex. When you ascend in the plane or hike up a mountain, the external pressure (the pressure that surrounds you) drops, while the pressure within the ear cavity (internal pressure) stays the same. This imbalance, creates the ear popping effect. 
• Volume of the Container: A low density of gas particles results in low pressure.
• Average Speed of the Gas Particles: A high density of gas particles results in high pressure. 
• SI Units of  Pressure:
• Force: 
• (N) newton
• 1 N= 1 kg x m/s2
• Pressure:
• (Pa) Pascal
• 1 Pa=1 N/m2
• Non SI Units of Pressure that are Commonly Used:
• atmosphere (atm)
• millimeter of mercury (mm Hg)
• torr
• Note that 1 mmHg=1 torr

Manometer: The instrument used to measure the pressure of a gas sample in the laboratory. They are U shaped tubes partially filled with a liquid that are connected to the gas sample on one side and open to the air on the other.

7.2 Apply the ideal gas law to relate and calculate values for pressure, volume, temperature, and amount of a gas (11.3-11.5) Apply the ideal gas law to relate and calculate values for pressure, temperature, and amount of gas. 

Pressure/Gas Laws:

• Measurable properties of gases include: 
• Pressure (P) in atm
• Volume (V) in L
• Temperature (T) in Kelvin (C+273.15)
• Amount in moles (n)
• The properties are interrelated-when one changes, it affects the others. 
• The Gas Laws: mathematical relationships that describe the quantitative relationships and behavior of gases as they are mixed, or subjected to pressure or temperature changes. 

These simple gas laws describe the relationship between pairs of these properties.

• Boyle's Law: describes how volume (V) varies with pressure (P) at constant temperature and amount of gas. 
• Boyle and Hook used a J shaped tube to measure the volume of a gas sample of gas at different pressures. They trapped a sample of air in the J tube and added mercury to increase the pressure on the gas. 
• Boyle's law states that volume and pressure have an inverse relationships where an increase in one causes a decrease in the other. 
• If the volume is increased, the pressure will decrease
• We can use Boyle's ;aw to calculate the volume of a gas following a pressure change or the pressure of a gas following a volume change. 
• As long as the temperature change and the amount of gas remain constant. 

p1v1=v1v2; Where p1 and v1 are the initial pressure and volume of the gas and p2 and v1 are the final volume and pressure. 

Charles's Law: The volume of a fixed quantity of gas maintained at constant pressure is directly proportional to its absolute temperature * (T in Kelvin) assuming that the pressure and the number of moles are constant. 

• AJ Charles was interested in gases and was among the first people to ascend in a hydrogen filled balloon. 
• When the temperature of a gas sample increases, the gas particles move faster; collisions with the walls are more frequent 
• The volume of a gas increases with increasing temperature. 
• Volume and temperature are linearly related. If two variables are linearly related, plotting one against the other produces a straight line. 
• Absolute zero * refers to the temperature at 0 K-colder temperatures do no exist. Anything less would have a negative volume which is physically impossible. 
• If a balloon is moved from an ice water bath to a boiling water bath, its volume expands as the gas particles within the balloon move faster (due to the increased temperature) and collectively occupy more space. 
• Ice water (low kinetic energy)----> Boiling water (high kinetic energy)

(v1/t1)=v2/t2 at constant pressure (P) and number of moles (n)

The pressure exerted on a sample of a fixed amount of gas is doubled at constant temperature, and then the temperature of the gas in kelvins is doubled at constant pressure. What is the final volume of the gas?

(a) The final volume of the gas is twice the initial volume.

(b) The final volume of the gas is four times the initial volume.

(c) The final volume of the gas is one-half the initial volume.

(d) The final volume of the gas is one-fourth the initial volume.

(e) The final volume of the gas is the same as the initial volume.

Avogradro's Law: The volume (v) of a gas maintained at constant temperature and pressure is directly proportional to the number of moles of the gas (n)

• When the amount of gas in a sample increases at constant temperature and pressure, its volume increases in direct proportion because the greater number of gas particles fills more space. 
• More gas molecules=larger volume
• We can use Avogrado's law to calculate the volume of a gas following a change in the amount of the gas as long as the pressure and temperature are constant. 
• The volume of a gas sample increases linearly with the number of moles in the gas sample. 
• When the amount of gas in a sample increases at constant temperature and pressure, its volume increases in direct proportion because the greater number of gas particles fill more space. 
• Equal volumes of gases contain equal numbers of molecules (gas doesn't matter)

• R (gas constant) is the proportionality constant in the ideal gas equation
• R relates pressure, volume, temperature, and the number of moles of gas in the ideal gas equation.
• Molar Volume: Volume of one mole of a gas determines at STP (standard temperature and pressure)
• An ideal gas is a gas whose volume, pressure, and temperature may be described by the ideal gas equation

7.3 Apply Dalton's Law of Partial Pressure to calculate the pressure of combined gases and to calculate the partial pressures of gases in mixtures. (11.6)

7.4 Describe gases in terms of the Kinetic Molecular Theory (11.7, 11.8)

• Kinetic Model Theory: The simplest model for the behavior of gases. In this theory, a gas is modeled as a collection of particles (either molecules or atoms, depending on the gas) in constant motion. The basic postulates (or assumption) of kinetic molecular theory are listed below:
• 1.) The size of a particle is negligibly small. KMT assumes that the particles themselves occupy no volume, even though they have mass. 
• Smaller molecules move faster then larger molecules. 
• 2.) The average kinetic energy of a particle is proportional to the temperature in kelvins. The motions of atoms ad molecules in a gas is due to thermal energy, which distributes itself among the particles in the gas.  
• The average kinetic energy of the gas particles is directly proportional to the Kelvin temperature. Nothing else. 
• The temperature of the gas increases, the average speed of the particle increases. 
• However, not all the gas particles are moving at the same speed. 

• 3.) The collision of one particle with another (or with the walls of its container) is completely elastic. This means that when two particles collide, they may exchange energy, but there is no overall loss of energy. 

7.5 Differentiate Effusion and Diffusion (11.9)

Diffusion: The process by which gas molecules spread out in response to a concentration gradient, and even though the particles undergo many collisions, the root square velocity still influence the rate of diffusion.

Effusion: The process by which a gas escapes from a container into a vacuum through a small hole.

7.6 Distinguish between Ideal and Real Gases (11.11)

• Ideal gases are hypothetical-there is no gas that will exactly follow the ideal gas behavior. 
• However, may gases will behave very closely to ideal gases under certain conditions. 
• According to KMT, ideal gas laws assume:
• No attractions between gas molecules
• gas molecules do no take up space
• At low temperatures and high pressures these assumptions are not valid. 
• Real gases will behave more like ideal gases under conditions of:
• Low pressure (at higher pressures, the volume of the gas molecules become significant
• High temperatures (at lower temperatures, the inter molecular attractions can lower the expected pressures)

Chapter 9: Introduction to Solutions and Aqueous Reactions

9.3 Solution Stoichiometry (307)

• In aqueous reactions, quantities of reactants and products are often specified in terms of volumes and concentration. We can use the volume and concentration of a reactant of a reactant or product to calculate its amount in moles. 
• Volume A==> Amount A (in moles) ==> Amount B (in moles) ==> Volume B

9.4 Types of Aqueous Solutions and Solubility

• When a solid is put into a liquid solvent, the attractive forces that hold the solid together (the solute-solute interactions) compete with the attractive forces between the solvent molecules and the particles that compose the solid (solvent-solute interactions)
• Electrolytes and Non electrolyte Solutions
• A salt solution (Electrolyte) conducts electricity while a sugar solution (Non electrolyte) does not. 
• Salt (ionic) 
• Electrolyte: Substances that dissolve in water to form solutions that conduct electricity.
• Strong Electrolyte: Substances that completely dissociate into ions when they dissolve in water. Strong acids are strong electrolytes. 
• Weak Electrolyte: Does not completely ionize in water. Weak acids are weak electrolytes that conduct electricity weakly. 
• Sugar (molecular)
• Most molecular compounds-with the exception of acids- dissolve in water as intact molecules. Sugar dissolves because the attraction  between sugar molecules and water molecules, overcome the attraction of sugar molecules to each other. 
• Compounds such as sugar that do not dissociate into ions when dissolved in water are nonelectrolytes that do not conduct electricity. 
• Acids (Molecular)
• Ionize to form H+ ions when they dissolve in water. 
• Ex. HCI ionizes into H+ and Cl- when it dissolves in water.  HCI (aq) ==> H+(aq) + Cl- (aq)
• Strong acids completely ionizes in solution. 
• In general, a compound is soluble if it dissolves in water and insoluble if it does not.

9.5 Precipitation Reactions

• Precipitation Reaction: a reaction in which a solid forms upon the mixing of two solutions. 
• However, they do not always occur when two aq solutions are mixed. 
• The key to predicting reactions is understanding that only insoluble compounds form precipitate. Two solutions containing soluble compounds combine and an insoluble compound precipitate. 
• Practice Problems (317)

9.6 Representing Aqueous Reactions: Molecular, Ionic, and Complete Ionic Equations

• Molecular Equation: A chemical equation showing the complete, neutral formulas for every compound in a reaction. 
• Complete Ionic Equation: Chemical equation showing all of the species as they are actually present in solution
• Net Ionic Equation: Equation showing only the species that actually change during the reaction.

9.7 Acid-Base Reactions

9.9 Oxidation-Reduction Reactions

Chapter 8: Chemical Reactions and Chemical Quantities

8.3 Writing and Balancing Chemical Equations

• Chemical changes occur via chemical reactions. 
• A combustion reaction is a particular type of chemical reaction in which a substance combines with oxygen to form one or more oxygen containing compounds. Combustion also emit heat. 
• We represent a chemical reaction with a chemical equation. The substances on the left are reactants and the ones on the right are products. 
• When a chemical equation is not balanced, it violates the law of conservation of mass because an atom is not formed out of nothing. 
• To correct an unbalanced equation-or write an equation that closely represents what actually happens-we must balance it by changing the coefficients to ensure that the number of each type of atom on the left is equal to the right side. New atoms do not form during a reaction, nor do atoms vanish-matter must be conserved. 
• What quantity or quantities must always be the same on both sides of a chemical equation?
• the number of atoms of each kind. 

8.4 Mole to mole and mass to mass conversions

• The coefficients in a chemical equation specify the relative amounts in moles of each substance involved in the reactions. 
• The numerical relationships between chemical amounts in a balanced chemical equation are called stoichiometry, which allows the prediction of the amount of products that will form in a chemical reaction based on the amount of reactants that react. It also allows is to determine the amount of reactants necessary to form a given amount of product. 
• This is essential in knowing the quantity of chemical reactants to obtain products in the desired quantities. 
• Think of it as a recipe. The amount of ingredients (cheese, crust, sauce) make a certain amount of pizzas (products). 

2 C8H18 (l) + 25 O2 (g) ====> 16 CO2 (g) + 18 H2O (g)

• Making Molecules: Mole to Mole Conversions
• In a balance chemical equation, we have a "recipe" for how reactants combine to form products. From our balanced equation for the combustion of octane for example is the following stoichiometric ratio:
• 2 mol C8H18: 16 mol CO2
• We can use this ratio to determine how many moles of CO2 from when a given number of moles of C8H18 burns. 
• Suppose we burn 22.0 mol C8H18, how many moles of CO2 form?
• 22.0 mol C8H18 X (16 mol CO2/2 mol C8H18)= 176 mol CO2
• The combustion of 22 moles of C8H18 adds 176 moles of CO2 to the atmosphere. 
• Making Molecules: Mass to Mass Conversions
• This calculation is similar to the Mole to Mole conversions exept that we are now given mass of octane instead of the amount of octane in moles, Consequently we must first convert mass (in grams) to the amount (in moles). (pg. 280)
• Mass A ==> Amount A (in moles) ==> Amount B (in moles) ==> Mass B
• Practice Problems 8.4 and 8.5 (pg. 281 and 282)

Video Overview

Mole to Mole Conversions

Mass to Mass Conversions

Mass to Mass Conversions Part II

8.5 Limiting Reactant, Theoretical Yield, and Percent Yield

• The three most important concepts in reaction stoichiometry: limiting reactant, theoretical yield, and percent yield.
• Limiting Reactant (Reagent):The substance that limits the amount of products that can be made
• Excess Reactant: is any reactant that occurs in a quatity greater than is required to completely react with the limiting reactant.
• Theoretical Yield: The amount of product that can (theoretically) be made based on limiting reactants
• Actual Yield: The amount of product that is (actually) made. Always equal or less than theoretical yield. 
• Percent Yield: Ratio percentage. (Actual Yield/Theoretical Yield) X 100

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.