Regents Review Sheet 1 – Atomic Structure These are some important points to remember about Atomic Structure. Use this sheet when you do the Review Problems. Points of Interest

• Law of Conservation of Matter-matter can not be created or destroyed

• Law of Constant Composition-a compound always contains the same elements in the same proportion or ratio by weight

Atom - The smallest particle of an element that retains the chemical identity of the element • A) Electrons • B) Nucleons-particles in the nucleus; 2 kinds

– 1. Protons-positively charged • Have a mass of 1 amu (atomic mass unit)

– 2. Neutrons-neutral • Also have a mass of 1 amu

Experiment • Rutherford did an experiment where he shot alpha particles (very

small particles) at a piece of gold foil – He found that most of the alpha particles go right through, but

that occasionally one of the particles was deflected a whole lot, as if it had ricocheted off of something solid. It hit the nucleus.

Atomic Number-Tells the number of protons in the nucleus Ions-atom with an electrical charge

– Anions are negative ions (GAINED –e) – Cations are positive ions (LOST -e)

Isotopes-same elements have different numbers of neutrons The Bohr Model of the Atom

• The closer the electron is to the nucleus, the lower the energy

• Note: only a certain amount of –e fit into a given energy level

• These principal energy levels approximate how far an electron is from the nucleus

Ground State-When the electrons are in the lowest available energy levels, the atom is in the ground state Excited State - If the atom absorbs energy, the electrons become “excited” and may jump up to a higher energy level Orbitals - The space within an atom where an electron or pair of electrons is likely to be found Valence Electrons - Electrons in the outermost principal energy level (also called shell) of an atom

• Cl has 17 electrons. Of these, 7 are in the 3rd principal energy level. Therefore, it has seven valence electrons

• The kernel is the atom except for the valence electrons

Ionization Energy-The amount of energy required to remove the most loosely bound electron from an atom in the gaseous phase Reference Table Information for this Unit A. Periodic Table – gives Atomic Mass, Atomic Number, Electron Configuration, etc.

Regents Review Sheet 2 – Periodic Table These are some important points to remember about Periodic Table. Use this sheet when you do the Review Problems. Points of Interest

• The periodic table organizes the elements by their properties • The elements are arranged in rows called “periods” and columns

called “groups” • The groups hold elements that are similar and have related properties • The periods (7 of them) represent the principal energy levels. The 1st

period contains elements with electrons in the 1st principal energy level. The 4th principal energy level has elements with electrons in the 4th principal energy level

• The elements are arranged in increasing atomic number (in increasing numbers of protons in the elements)

• Groups (columns) are labeled 1 -18 • Group 1 is the alkali metals • Group 2 is the alkaline earth metals • Group 17 is the halogens • Group 18 is the noble gases • In the table, there is usually a “staircase” drawn toward the right-hand

side; the elements to the left are metals, and the elements to the right are non-metals

• Bordering the steps are the metalloids or semi-metals; B, Si, As, Te, Ge, Sb, At

• Most elements are solids at room temperature except: – Hg and Br liquids – H, O, N, F, Cl and the noble gases, which are all gases

Trends and Properties

• 1) Atomic Radius- half the distance between the centers of 2 atoms of the element that are just touching each other

– Atoms get larger as you go down a group; this makes sense – Atoms get smaller as you go across a period from left to right.

It occurs because as an atom gets more protons, the electrons

are more strongly attracted to the nucleus, and the atoms radius decreases.

– Also as an electron is gained or lost by an atom, the radius changes

- If you add electrons, forming a negative ion, you get a larger radius - If you lose electrons, forming a positive ion, you get a smaller radius

• 2) Ionization Energy- The amount of energy required to remove the most loosely bound electron from an atom in the gaseous phase

• Ionization Energy decreases as you move down a group • It increases as you go across a period from left to right These trends occur because larger atoms are not able to hold their electrons as tightly as smaller atom

• Electronegativity - How well an atom attracts another atom’s electrons

Regents Review Sheet 3 – Bonding These are some important points to remember about Bonding. Use this sheet when you do the Review Problems. Points of Interest

Ionic Bonds-really an attraction between a positively charged metal ion and a negatively charged non metal ion

– During a reaction between these metals and the non-metals, electrons in the valence shell of the metal are transferred to the valence shell of the non-metal

– Note: By forming ionic bonds, a metal and non-metal achieve the stable octet in the valence shell

Covalent bonds are formed by a shared pair of electrons between 2 atoms in a molecule

- Bonds-the atoms share 4 –e (2 pairs of –e)

-Triple Bond-the atoms share 6 –e (3 pairs of –e)

General Rules for Lewis Dot Diagrams • 1) On a sheet of paper write the element symbol with the number of valence

electrons around it • 2) Put all the atoms together in such a way as to satisfy the octet rule

Octet Rule • Atoms tend to gain, lose or share electrons in order to acquire a full set of valence

electrons • In other words, each atom will become involved in bonds which will result in

them becoming like noble gases

Crisscross Method • If you have two atoms that you know will react, there is an easy way to see what

the molecular formula of the compound they will form is. – 1) Look at the periodic table and find the charges in the upper right hand

corner of the box. Write these numbers to the upper right of the atomic symbol

– 2) Crisscross the numbers • Example - Find the molecular formula of the product of a reaction between Ca

and F – 1) Ca2+ F1- – 2) Ca1F2

ElectronegativityShows the strength of attraction that one atom has for an electron

• Different atoms have different electronegativities; what this means is that some atoms attract electrons better than other atoms do

• The atom with the stronger electronegativity will actually hold the shared electron closer to itself than to the other atom in the covalent bond

• Electronegativity Difference

• Shows whether a bond is Ionic, Polar Covalent or Non-Polar Covalent • If the difference is ≤ 0.4 it is Non-polar Covalent • If the difference is between 0.4 and 1.7 it is Polar Covalent • If the difference is > 1.7 the bond is Ionic

You can also determine if a bond is Ionic, Polar Covalent or Non-Polar Covalent in this way Ionic – between a metal and a non-metal Polar Covalent – between 2 different non-metals

Non-Polar Covalent – between 2 identical non-metals (diatomic molecules)

Coordinate Covalent Bond - When one atom in the covalent bond has donated both of the electrons in the shared pair – (example is Ammonia) Network Solids - Atoms are linked throughout the sample by covalent bonds; this makes the substance very hard and strong

• The substance will also have very high melting points • They are poor conductors of heat and electricity • Diamonds, silicon carbide and silicon dioxide are examples

Metallic Bonding - Occurs between atoms of metals

• Metals form + ions; these + ions are linked by an attraction to a bunch of free floating –e (this is why metals conduct electricity)

• The metal ions don’t have full valence shells, so they pick up –e and lose –e often • This makes these substances very malleable

Molecular Attraction or Interactions

• Remember polar covalent bonds-molecules that have this type of bond are called dipoles

• Dipoles- one side of the molecule is slightly positive and one side is slightly negative.

• Example is HCl - each HCl has a + side and a – side. When 2 HCl molecules come in contact with each other, the + H of one becomes attracted to the – Cl of the other and they stick to each other

Hydrogen Bonds

• Occur when H is bonded to an element with a small atomic radius and high electronegativity

• H becomes slightly + since the other element doesn’t share the electron evenly with it

• Because of this slight + charge, H is attracted to - atoms Attractions and Molecule Shape

• When you have more than 2 atoms in a molecule, the shape of the molecule determines its polarity; when polar bonds are uniformly distributed around a molecule, the slight charges cancel each other out and the molecule is actually non-polar

Van Der Waals Forces

• Even without polar attractions and H-bonds, there are still weak attractive forces between molecules

• A trick to remembering the characteristics of Van Der Waals forces is to think of it as being like gravity; large atoms have powerful VDW forces, while small atoms have weak VDW forces.

• Also, the closer the atoms the more powerful the VDW forces.

Regents Review Sheet 4 – Naming Chemical Compounds, Shapes of Compounds and Balancing Chemical Equations These are some important points to remember about Naming Compounds. Use this sheet when you do the Review Problems. Rules for Naming Chemical Compounds

• A) Binary Compounds composed of a Metal and Non-Metal – Metal is written first – Non-metal is written second, but the ending –ide is added to it – Ex. NaCl sodium chloride

• B) Binary Compounds with 2 non-metals – Less electronegative element is written first – The other element is given the ending-ide – Ex. NBr nitrogen bromide

• C) Other Binary Compounds

– In some binary compounds in which there can be more than one atom of the same element, we use prefixes to state the number of that atom

– Ex. CO carbon monoxide CO2 carbon dioxide

• D) Tertiary Compounds composed of Polyatomic Ions – You can find a list of polyatomic ions in the reference tables – Metal is written first – Polyatomic ion is written last – Ex. Na(OH) sodium hydroxide

Mg(SO4) magnesium sulfate There is an exception to the rule

– Ammonium is written first – Ex. NH4Cl ammonium chloride

• E) Binary Acids

– Hydrogen bound to another element is often a binary acid – Write the other element with the prefix hydro-, the suffix –ic,

then add the word acid – Ex. HCL hydrochloric acid

H2S hydrosulfuric acid HF hydroflouric acid

• F) Ternary Acids-usually are written with H’s in front

– No hydro- prefix – Take the polyatomic ion name, change the ending to –ic, and

add the word acid – Ex. H2SO4 sulfuric acid

H2CO3 carbonic acid

• G) Compounds With Metals That Make Multiple Ions • Some metals can make several different ions, such as iron

• Example Fe+2 0-2 makes Fe0 this is called iron (II) oxide – the (II) refers to the ion or oxidation state of the metal iron

• I could also have Fe+3 0-2 come together to make Fe2O3 – this is called iron (III) oxide

• Note: you do the criss-cross method to find these molecular formulas

Molecular Shape Or Geometry Linear Molecules

• All molecules that have just two atoms are linear or in a straight line. • Linear molecules with three atoms occur when the center atom is

bonded to two other atoms and there are no additional unpaired electrons

Trigonal Planar

• To get a trigonal planar molecular shape, the central atom must be bonded to 3 other atoms and must not have any extra unshared electrons

Tetrahedral

• Central atom is bonded to 4 other outside atoms Pyramidal

• It looks like a tetrahedral molecule except you have only 3 atoms bound to the central atom and 1 pair of unshared electrons on that central atom.

Balancing Equations • Each side of the equation must have the same number of atoms of

each element; this is because of the Law of Conservation of Matter. If an element is in a reactant prior to a chemical reaction, then it has to be somewhere in the products after the reaction takes place.

• By Balancing Equations we insure that this is true. • A way to do this is to list your Reactants and Product Al2(SO4) 3 + Ca(OH) 2 ó Al(OH) 3 + Ca(SO4) Reactants Products 2 Al 1 Al 3 SO4 1 SO4 1 Ca 1 Ca 2 OH 3 OH

Regents Review Sheet 5 – Matter and Energy These are some important points to remember about Matter and Energy. Use these sheets when you do the Review Problems. Matter and Energy

• Chemistry-the study of the composition, structure and properties of matter, the changes that matter undergoes and the energy accompanying those changes

• A. Substances

• homogeneous matter having identical properties and composition

• All samples of a particular substance will have the same heat of vaporization, melting point, etc.

• Substance include: • 1. elements-composed of atoms of the same atomic number;

cannot be decomposed by chemical means

– 2. compounds-two or more different elements chemically combined in a definite ratio by weight; can only be decomposed by a chemical means, not by a physical means

• Binary compounds-composed of just 2 elements • Ternary compounds-composed of 3 elements

– B. Mixtures-combinations of varying amounts of two or more substances that differ in properties and composition

– Homogeneous mixture-uniform intermixture of particles; also called solutions

– One substance dissolves in the other – Heterogeneous mixtures-have uniformly dispersed ingredients,

but they don’t dissolve into one another Energy- The Ability to Do Work

• Forms of energy- – A. mechanical energy – B. heat energy – C. Radiant energy-light, radio waves, etc. – D. chemical energy – E. nuclear energy

Types of energy 1. Potential energy- the energy of position 2. Kinetic energy- the energy of motion Heat energy

• 1. Exothermic reactions-when heat energy is given off

• 2. Endothermic reactions-when heat energy is absorbed

Specific Heat • The amount of heat energy required to raise the temperature of 1 g of

a material by 1 C° – Q = m c (change in T)

Temperature Average kinetic energy of the particles in a system Converting from Kelvin to Celsius: Kelvin = 273 + Celsius

Phases of Matter The term phase refers to the gas (g), liquid (l) or solid (s) forms of

matter Solid Phase

• Atoms or molecules that make up the substance are in a fixed position • While these atoms or molecules can’t move around freely, they are

free to vibrate – This vibration increases as the temperature increases

Gaseous Phase

• The molecules are undergoing Translation ( that is, the attractions between the molecules are broken and the molecules are able to move about one another very freely)

• The molecules also undergo Rotation and Vibration Liquid Phase

• The atoms or molecules are vibrating, translating, and rotating, but there is still a stronger attraction between these particles than there is in a gas

• The liquid in a container stays in the container even if its top is open because of this stronger attraction

Heating Curve

Boyle’s Law-if the temperature is held constant, the volume and pressure will have and inverse relationship

– P1 V1 = P2 V2 Charles Law - At constant pressure the volume of a given mass of a gas varies directly with the Kelvin temperature

• V1 = V2 T1 T2

• Gay Lussac Law -When the volume is constant, there is a direct relationship between temperature and pressure

• T1 = P1

T2 P2 The Combined Gas Law

• Also called “The Easy Way To Remember the Gas Law Formulae” because instead of memorizing the previous 3 formulae, you only need to memorize this 1

• It allows you to also see the effect if two variables of a gas change

• P1 V1 = P2 V2 T1 T2 Standard Temperature and Pressure (STP)

– Temperature gas 0 °C or 273 K

– Pressure 760 mmHg or 760 torr or 1atm

Partial Pressures • The pressure exerted by each of the gases in a mixture of gases • P Total = P gas 1 + P gas 2+ P gas 3 +…P gas n

Ideal Gas Model

• A) A gas is composed of individual particles which are in continuous, random, straight-line motion

• B) Not all particles of a gas have the same kinetic energy, but the K temperature of the gas gives the average kinetic energy of the gas

• C) Collisions between gas particles are totally elastic. This also holds true for collisions between gas particles and the sides of the container. This means that no kinetic energy is lost when things collide.

• D) The volume of the gas particles is small compared to the volume of the container

• E) Gas particles are considered to have no attraction for one another.

• Hydrogen and Helium are the 2 most ideal gases. Do not forget this! Characteristics of Liquids

A) Vapor Pressure-since the molecules in a liquid are freely moving, they collide, transferring energy from one to another. This transfer of energy allows some to break free of the liquid and become gaseous

– This tends to occur near the surface of the liquid – This vapor near the surface actually imparts a pressure on the

underlying liquid – As temperature increases, so does vapor pressure

B) Boiling Point-a liquid will boil at the temperature at which the vapor pressure equals the pressure on the liquid

C) Heat of Vaporization-energy required to vaporize (evaporate) a unit mass of a liquid to a gas at constant temperature

Characteristics of Solids

• Density = mass/volume

• A) Crystals-particles in a crystal vibrate but they keep their regular pattern

• B) Melting Point-temperature at which a solid will change to a liquid

at 1 atm of pressure • C) heat of Fusion-energy required to change a unit mass of a solid to a

liquid at a constant temperature Sublimation

• Occurs when certain substances change from solids directly into gases, without passing into a liquid state

– Examples are dry ice, iodine crystals and naphthalene

Review Sheet 6 – Math and Chemistry The following are some points of interest in Math and Chemistry. Use this sheet when answering these questions.

• Molecular Mass- to find the molecular mass, you must add up the masses of each atom in the molecule

• Ex. Find the molecular mass of H2SO4 Atomic Element # of atoms Mass Product H 2 1amu 2amu S 1 32amu 32amu O 4 16amu 64amu 98amu Mole/Mass/Volume Interconversions

Stoichiometry

• Branch of chemistry concerned with mass and volume relationships of reacting substances

• Remember, the coefficients in a balanced equation can be thought of in 2 ways: it is the # of molecules of the chemical in the reaction or it is the # of moles of the chemical involved in the reaction. In these problems, it is best to think of them as moles.

Example 1 – Mass Mass Problem • How many grams of oxygen are required to combine with 8g of hydrogen gas

to form water? 1) 8g X 2 H2 + O2 ó 2 H2O Put grams above 2mol 1mol and moles below 2) 4mol X 2 H2 + O2 ó 2 H2O Convert grams to 2mol 1mol moles 2) 4mol X 2 H2 + O2 ó 2 H2O Convert grams to 2mol 1mol moles 3) Set up a ratio with these numbers and solve for X: 4mol = X 2mol 1mol 2mol = X So you need 2mol of O2 or 64 g. Example 2 – Mass Volume Problem

• Find the mass of aluminum required to produce 1.32 L of hydrogen gas at STP from the following reaction:

• 2Al + 3 H2SO4 ó Al2(SO4) 3 + 3 H2

• 1) Write the Liters above and the moles below each substance in the equation X 1.32L 2Al + 3 H2SO4 ó Al2(SO4) 3 + 3 H2 2mol 3mol

• 2) Now, you don’t need to convert Liters to moles. Simply set up the ratio as it is:

X = 1.32L 2mol 3mol X = 0.88L You have 0.88 L of Aluminum

• 3)Now convert these Liters to moles (0.88L)/ (22.4L/mol) = 0.04mol

• 4)Lastly they want mass so convert to mass: (0.04mol) x (27g/mol) = 1.08 g of aluminum

Example 3 – Volume Volume Problem • Real easy. No need to convert anything.

• If 0.38L of hydrogen gas reacts with chlorine gas, what volume of HCl gas

will be produced? H2 + Cl2 ó 2 HCl

• 1) Put the Liters above and the moles below the substances in the equation. 0.38L X H2 + Cl2 ó 2 HCl 1mol 2mol

• 2) Set up the ratio with these numbers: 0.38L = X 1mol 2mol

• X = 0.76 L You are done! Example 4 – Mole Mole Problem

• Once again, no conversions are necessary.

• How many moles of HCl are needed to react with 2.3 moles of Zn in the following reaction?

2 HCl + Zn ó ZnCl2 + H2 • 1)Put the number of moles you have above, and the number of moles in the

equation below. X 2.3mol 2 HCl + Zn ó ZnCl2 + H2 2mol 1mol

• 2) Set up the ratio: X = 2.3mol 2mol 1mol X = 4.6 mol

Percentage Composition is the ratio of the mass of element in a compound to the mass of the whole compound Example 5- Percentage Composition of FeS Atomic Element # atoms Mass Product Fe 1 56 56 S 1 32 32 88g/mol For 1 mol: 56 g/mol Fe has a % composition = 88 g/mol = 63.6 % 32 g/mol S has a % composition = 88 g/mol = 36.4 % Empirical Formula From Percentage Composition

• If you have an unknown chemical, and you decompose it into its constituent elements, you can find its % composition. Once you have this you can determine the empirical formula of the unknown chemical and thus identify what it is.

Example 6 - A chemist analyzes an unknown compound and finds that it is 70.9% K and 29.1% S in composition. What is the empirical formula of this compound?

• Step 1 – convert the given percentages to grams – 70.9 g of K and 29.1 g of S

• Step 2 – divide these masses by the element’s atomic mass in grams – K: 70.9 g = 1.82 mol

39 g/mol

- S: 29.1 g = 0.91 mol 32 g/mol

• Step 3 – divide the larger of these numbers by the smaller. This gives you the ratio of the elements.

1.82mol/0.91mol = 2 => This means you have 2 K’s for each S

• Therefore the empirical formula is K2S Density of Gases

• Density = mass/volume or Density = GFM/ (22.4L/mol) • Because for any gas you know that the Gram Formula Mass is the

Mass of 1 mole of the gas and that the volume of 1 mole is 22.4 L/mol

Regents Review Sheet 6 – Solutions These are some important points to remember about Solutions. Use this sheet when you do the Review Problems. Points of Interest Solutions - A solution is a homogeneous (“same throughout”) mixture of two or more substances in a single physical state

• Solute – substance that is dissolved

• Solvent – substance that does the dissolving

• (something that is insoluble will not dissolve) • 1. Solid solutions – not all solutions are liquids; some are

solids • Alloys are solutions where one metal has been dissolved in

another (when both were melted) • Examples of these are gold and sterling silver; both have some

copper mixed into them to strengthen the metal • 2. Gaseous solutions – all mixtures of gases, including air

• 3. Liquid solutions – can dissolve solids, liquids or gases

into liquids – Liquids – when one liquid is dissolved in another they

are said to be miscible • Water and ethanol

– When two liquids won’t dissolve together, they are said to be immiscible

• Oil and water

• Aqueous Solutions - Solutions in which water is the solvent • When an ionic substance is dissolved in water it forms ions – this

is called an electrolyte, since the solution can then conduct electricity very well

– Example: NaCl in water Concentration of Solutions

• The amount of a solute in a given amount of a solvent (or how much stuff is dissolved in the solvent)

• 1.) Molarity (M) = moles of solute

liters of solution

• 2.) Parts per Million - Another common measure of concentration is parts per million

• PPM = grams of solute

grams of solution x 1,000,000

• 3.) Percent Solution-Another way that the concentration of solutions are often expressed is in Percent solution

• Percent solution = grams solute

100 mL solvent x 100 Saturated Solutions -When no more solute can be dissolved in a solution

• A saturated solution contains as much solute as can possibly be dissolved under the existing conditions of temperature and pressure

Unsaturated Solutions - If there is less solute than can be dissolved (for example, you could actually dissolve more in if you wanted), the solution is said to be unsaturated Supersaturated Solutions - You can sometimes make a solution that holds more solute than is normally present in a saturated solution

• You can make these by cooling a saturated solution. As they cool, most saturated solutions will precipitate solute out of solution. For some solutions this does not occur and a supersaturated solution will form.

Know the Solubility Curve (Table G) – If the amount added is below the curve, it is unsaturated, on the curve is saturated, above the curve can be either saturated or unsaturated depending on whether it fully dissolve or not.

Solubility Factors – • The things that determine how something will dissolve into something

else

• 1.) Nature of the solute and solvent – “Like dissolves like” – polar solutes dissolve in polar solvents very well

– Example: CH3OH and H2O mix easily because both are polar Non-polar solutes dissolve in non-polar solvents

– Oil dissolves in turpentine because both are non-polar Non-polar substances don’t dissolve well in polar solvents

– Oil in water – CH4 doesn’t dissolve in water

“Like dissolves like”

• 2.) Temperature – gases in liquids are more soluble at lower temperature

– Solids are more soluble in higher temperatures

• 3.) Pressure – increasing the pressure increases the solubility of gases in liquids (think of the pressure as pushing the gas into the liquid)

– Pressure has no effect on the solubility of liquids or solids in a liquid solvent

Colligative Properties –

• Property that depends on the concentration of solute, but not on the nature of the solute (how much you have makes a difference, not what it is)

• 1.) Vapor Pressure – decreases as a non-volatile solute is added to

a liquid • 2.) Boiling Point – related to 1.) above

• As a non-volatile solute is added, the boiling point of the solution increases

• 3.) Freezing Point Depression – a dissolved solute will normally lower the freezing point of the solvent

Regents Review Sheet 8 – Kinetics and Equilibrium Here are some of the most important concepts in Kinetics and Equilibrium. Use this sheet when doing the Review Problems for this topic. Equilibrium – exists when two opposing processes occur at the same rate

• Chemical Equilibrium Example: • If you put NO2 into a container, the following reaction occurs:

2 NO2 ó N2O4 browncolorless gas gas

• The reason for this is because the reaction never reaches completion. The reactants are never totally converted into products.

• This is a reversible reaction, which means that while the products are being formed from the reactants, the reactants are being reforme from the products (the reaction goes in the forward and reverse directions at the same time)

• When the forward and reverse reactions reach a point where the rates are the same, they are said to be in chemical equilibrium.

• At this point the concentration of the products and reactants remain constant over time. (You are making as much NO2 as N2O4 )

• Concentration is symbolized by placing square brackets around a substance’s symbol

• [NO2] = concentration of NO2 The Law of Mass Action – expresses the relative concentration of reactants and products at equilibrium in terms of an equilibrium constant called Keq .

• Consider this fake reaction: • aM + bN ó cP + dQ • The equilibrium expression is :

• Keq = [P]c [Q]d [M]a [N]b

• It measures the extent to which a reaction goes to completion

• If it is a very large number it means you made a lot of product, so the reaction proceeded to completion (or close to it)

• If it is a small number, you didn’t make much product, so it did not proceed to completion. The reaction barely got going.

LeChatelier’s Principle - If a change in conditions is imposed on a system at equilibrium, the equilibrium position will shift in a direction that reduces this change (makes it go back to the way it was before the change

• 1). Change in Concentration – – 2 NO2 (g) ó N2O4 (g)

– If the above reaction is at equilibrium and I add some more

NO2 (g), this extra NO2 (g) will be used up to produce N2O4 (g). The reaction will then return to equilibrium.

• 2). Change in Pressure- a reaction at a certain volume will try to counteract an increase in pressure by producing fewer molecules (since they take up less space)

2 NO2 (g) ó N2O4 (g)

2 NO2 (g) = 2moles = 44.8L N2O4 (g) = 1 mole = 22.4L

So if I increase the pressure, more N2O4 (g) will be made since it takes up less volume

• 3). Change in Temperature – depends on whether the reaction is

endothermic or exothermic You do this the same way that you do the concentration problems

H2 + I2 ó2 HI + heat

Here if I increase the heat it will force the reaction to go to the left and make more reactants

• Haber Process - N2 (g) + 3H2 (g) ó 2NH2 (g) + 91.8kJ

• In order to make more ammonia in the above reaction: • You can:

– 1. Remove the product continuously – 2. Increase the pressure

– 3. Increase the temperature – this sounds wrong, but it isn’t. Think about it in terms of the motions of particles

– Ksp Tells us the same thing that Keq does:

– Small Ksp = not much substance dissolves

– Large Ksp = a lot of substance dissolves – AgCl (s) ó Ag+ (aq) + Cl- (aq)

– Ksp = [Ag+ ] [Cl- ]

[AgCl] ó since it is solid, ignore

– Ksp = [Ag+ ] [Cl- ] Kinetics

• Chemical kinetics - the area of chemistry dealing with the speed at which reactions occur

• Rate = change / time

• Reaction rate – rate at which the reactants disappear and the products appear

• You can also think of it as the change in concentration of reactants and products in a certain amount of time

Collision Theory – molecules must collide in order to react

• Usually involves a collision between 2 molecules (occasionally 3 molecules, but the chances of 3 things colliding in the correct way at the same time are pretty slim)

• Effective collisions ó leads to product • Needs sufficient energy and needs to be aligned correctly during the

collision

Energy Diagram - Endothermic

Energy Diagram - Exothermic

• Activation energy – energy needed to start a reaction

• Activated complex – at this point the chemicals are neither product nor reactant; they are in an intermediate state between both (like a mixture of each)

• Factors Affecting Reaction Rate

• 1.) Nature of reactants – what the reactants are have a lot to do with how fast the reaction takes place

• 2.) Temperature – the higher the temperature, the faster the

rate. Why?

• 3.) Concentration – the higher the concentration, the faster the rate. Why?

• 4.) Surface Area – the higher the surface area, the faster the rate.

How do you increase surface area? Why does this increase the rate?

Energy Diagram – Addition of Catalyst

• 5.) Catalysts – these molecules orient the reactants so that they react better (basically hold the reactants in the correct position so the reaction takes place faster)

Spontaneous reaction – one that will proceed on its own without outside help § For a reaction to be spontaneous the Enthalpy must be decreasing (

= a negative number) § Entropy – a measure of the disorder of a system; Disorder is greatest

in gases and least in solids Summary

• A reaction is spontaneous if the Enthalpy decreases (is negative; see Table I in the reference tables) and the system has an increase in disorder (positive Entropy; this is easy to figure, since Entropy increases from solid to liquid to gas)