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Honors Chemistry Study Guide. Units 11-17

Leo Taffe, Vijay Ravindran, Jimmy Wiaduck, and Ben Ludtke

UNIT 11Modern Atomic Theory

Key Terms:Wavelength - The distance between two waves

Frequency - The number of waves that pass a given point per unit of time

Photon - Tiny packet of energy

Excited State - A state that contains an electron with excess energy

Ground State - The lowest possible energy state for an atom

Orbital - The region where an electron resides in an atom

Pauli Exclusion Principle - An atomic orbital can hold no more than two electrons, and they must be of opposite spin

Valence Electrons - Electrons in outermost occupied principal energy level

Core Electrons - The inner electrons of an atom

Ionization Energy - The energy required to remove an electron from an atom in the gas phase

Chapter Summary:

Rutherford's model of the atom

Atoms are so small that they are difficult to study directly Atomic models are constructed to explain experimental data on collections

of atoms He discovered this through his golf foil experiment; there is a small,

positive, and dense region known as the nucleus. Most of the atom is empty space (electrons)

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Electromagnetic Radiation

Light has 2 natures: it can exist as a particle of energy (photon or quanta) or as a wave

Electromagnetic radiation is the whole spectrum of light, not just visible light

The different types of waves are: radio waves, microwaves, infrared, visible light, ultraviolet, x-rays, and gamma rays

The shorter the wavelength, the more powerful the wave is A wavelength is the length of one wave and is represented by "λ" Different waves travel at different speeds The frequency of a wave determines the length of the wave. Inversely

proportional Frequency is represented by "v" Speed = wavelength (in meters) x frequency per second C = λv C = 2.998 x 10^8 m/s Max Planck discovered that energy and frequency are directly proportional E = hv Planck's constant (h) = 6.626 x 10^-34 J x S Planck called the energy packets "quanta"

Atoms and Light/Energy

Atoms can give off light. They first must receive energy and become excited. The energy is released in the form of a photon. When an electron is excited and filled with energy, it goes to the excited

state When it releases its energy in the forms of photons of light, it falls back

down to the ground state Bohr’s model of the atom --> Quantized energy levels, electrons move in a

circular orbit, and electrons jump between levels by absorbing or emitting photon of a particular wavelength

Electrons don't move in a circular orbit, but they do reside in orbitals

Orbitals and Energy Levels

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Orbitals don't have sharp boundaries All atoms contain energy levels. There are 7 different energy levels The more electrons an atom has, the more energy levels it has The energy levels are divided into sublevels (s, p, d, and f) 2 electrons with opposite spins can fit in 1 orbital The s sublevel only contains 1 orbital, the p sublevel contains 3 orbitals, the

d sublevel contains 5, and the f sublevel contains 7 Each row on the periodic table has sublevels

Electron configuration

The electron configuration determines elemental properties. It is the arrangement of electrons (for example, for Ne - 1s2, 2s2, 2p6) Noble gas shorthand - using a noble gas before the next set of levels to

decrease the amount of writing needed (Ex: At = [Xe] 6s2 4f14 5d10 6p5) ** always use the diagonal rule when determining the electron configuration

1s first principal energy level2s 2p second principal energy level3s 3p 3d third principal energy level4s 4p 4d 4f fourth principal energy level5s 5p 5d 5f fifth principal energy level6s 6p 6d sixth principal energy level7s seventh principal energy level

Orbital diagrams - use the amount of electrons and put them over the sublevels in order. Use arrows to show the spins

Core electrons are the inside electrons, valence electrons are the outside ones and are typically the ones that react. Atoms with the same valence electron arrangement often times behave in a similar manner

Metals lose electrons and form cations, nonmetals gain e-s and form anions The size of an atom decrease across a row and increase down a column The ionization energy, however, increases across a row and decreases down

a column

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Important names and equations:

Ernest Rutherford

Max Planck

Niels Bohr

Albert Einstein

Schrodenger

c = λv

E = hv

*** NOTE: You can manipulate these equations to make a different equation that will help you solve the problem

Ex: If c = λv and E = hv, then λ = ch/E

Review Questions:

1. As the wavelength increases, the frequency _____________

a) increases b) decreases c) doesn't change d) cannot be determined

2. When an electron in an atom absorbs energy, it _____________

a) moves to a higher energy level b) emits light c) moves to the ground state

3. Which of the following has the highest ionization energy?

a) K b) Ge c) Si d) O

4. Which neutral atom has seven 3d electrons? _____________

5. Write the electron configurations for: Silicon, Selenium, and Oxygen. You may use noble gas shorthand.

UNIT 12

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Chemical Bonding

Key Terms:

Bond - A force that holds atoms together

Bond Energy - The energy required to break a bond

Covalent Bond - Electrons are shared by 2 atoms

Ionic Compound - Formed by a metal and a nonmetal

Polar Covalent Bond - Electrons are shared unequally between 2 atoms

Electronegativity - The ability of an atom to attract shared electrons to itself

Resonance - More than one valid Lewis dot structure may be drawn for a molecule

Linear Structure - 2 electron pairs (2 bonds) on a central atom

Trigonal Planar Structure - 3 electron pairs (3 bonds) on a central atom

Tetrahedral Structure - 4 electron pairs (4 bonds) on a central atom

Trigonal Pyramid Structure - 3 single bonds and a lone pair on a central atom

Bent or V-Shaped Structure - 2 single bonds and a lone pair on a central atom

Chapter Summary:

Bonds

The strong electrostatic forces of attraction holding atoms together in a unit are called chemical bonds

Ionic compound results when a metal reacts with a nonmetal A covalent bond results when electrons are shared by nuclei A polar covalent bond results when electrons are shared unequally by nuclei This is when one atom attracts the electrons more than the other atom Electronegativity plays a large role in this. The more electronegative, the

more the pull on the electrons. It increases up and right If the difference in electronegativity between the 2 atoms is great, a polar

covalent bond can be formed

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If the difference is between 0.4 and 1.7, it is polar For example, HF is very polar Polar covalent molecules have dipole moments. This means that the center

for positive charge and the center for negative charge are separate Water has a dipole moment. This contributes to certain properties of water

Ions and Ionic Bonding

This is always helpful

Ions are formed after metals or nonmetals transfer valence electrons Atoms in stable compounds usually have a noble gas electron configuration Metals lose electrons to get to this state, nonmetals gain electrons Chemical compounds are always electrically neutral Ions of opposite charges are packed together in a tight bond Cations are smaller than the original atom, anions are larger Polyatomic ions work the same as regular ions, it's just the atoms in a

polyatomic ion are bonded together covalently to act as a unit

Lewis Dot Structures

ONLY include the valence electrons All atoms should be bonded so that they reach the octet rule (unless the

problem specifies otherwise) Bonding pairs are shared, lone pairs are not

. Single bond – covalent bond in which 1 pair of electrons is shared by 2

atoms

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Double bond – covalent bond in which 2 pairs of electrons are shared by 2 atoms

Triple bond – covalent bond in which 3 pairs of electrons are shared by 2 atoms

A molecule shows resonance when more than one Lewis structure can be drawn for the molecule (see below)

There are very few exceptions to the octet rule (BF3)

Molecular Structure

Three dimensional arrangement of the atoms in a molecule Linear - Atoms in a line (ex: all molecules with only 2 atoms, CO2, etc) Trigonal planar - Atoms in a triangle (BF3) Tetrahedral - CH4 These last three are nonpolar molecules, meaning they are symmetrical The VSEPR model (valence shell electron pair repulsion)

The molecular structure is determined by minimizing repulsions between electron pairs

Bent - Such as H20. It is NOT symmetrical. This is because there are lone pairs present, and the lone pairs force the shared electrons down, causing them not to be linear

Trigonal Pyramidal - Such as NH3. There is one lone pair present **NOTE - The bond type does not matter when you are predicting the

overall shape of the molecule. A triple bond is counted the same as a double or single bond

Review Questions:

1. A bond formed between two elements that have a very large difference in electronegativity is called _______________a) A covalent bond b) A polar covalent bond c) A double bond d) An ionic bond

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2. Which of the following molecules exhibits resonance?

a) NH3 b) SO3 c) CCl4 d) H2S

3. The shape of the NF3 molecule is ____________

a) Linear b) Bent c) Tetrahedral d) Trigonal Planar

4. Draw Lewis dot structures for each of the following molecules/ions and predict the structure for each

a) NI3

b) CS2

c) COCl2 (C is the central atom)

d) NH4+

e) OF2

5. Draw Lewis structures for the carbonate ion, CO3-2, showing all possible resonance structures:

UNIT 13Gases

Key Terms:

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Absolute Zero - The temperature at which a gas theoretically occupies zero volume

Universal Gas Constant - 0.08206 Liters x Atmosphere / (mol x temperature)

Boyle's Law - The volume of a given amount of gas is inversely proportional to its pressure

Charle's Law - The volume of a given gas is directly proportional to its temperature in Kelvin

Avogadro's Law - Volume of a gas is directly proportional to the number of moles of gas

Combined Gas Law - P1V1/T1 = P2V2/T2 OR P1 V1 T2 = P2 V2 T1

Dalton's Law of Partial Pressures - The total pressure of a mixture of gases is equal to the sum of each gas

Molar Volume of Ideal Gas Law - 22.4 L

Chapter Summary:

Gases

Gases, unlike liquids and solids, are easily compressed because of the space between the particles

Under pressure, the particles in a gas are forced closer together. The amount of gas, volume, and temperature are factors that affect gas

pressure. Pressure, Volume, Temperature, and the number of moles are 4 very

important things in describing a gas The more moles, higher pressure. Higher temperature, higher pressure.

Lower volume, higher pressure. Barometer – device that measures atmospheric pressure. It uses mm Hg Invented by Evangelista Torricelli in 1643 1 standard atmosphere

= 1.000 atm = 760.0 mm Hg = 760.0 torr

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= 101,325 Pa A manometer measures the pressure of a gas in a container. Boyle figured out that pressure and volume were inversely proportional Another way of stating Boyle’s Law is P1V1 = P2V2

Charles found that volume and temperature were directly proportional His equation is V1/T1 = V2/T2 OR V1T2 = V2T1 Avogadro's law --> V1/N1 = V2/N2

Ideal Gas Law

This is the most crucial thing to know. It is the hardest to remember and also the most important.

Combining all of the other equations gives us this equation: PV = nRT In this equation, R is the universal gas constant (R = 0.08206) Volume in liters at STP = 22.4L

Partial Pressures

the total pressure exerted is the sum of the partial pressures of the gases present.

Ptotal = P1 + P2 + P3

The volume of the individual particles is not very important The forces among the particles must not be very important Total pressure is the pressure of the gas + the vapor pressure of the water. Diffusion is the tendency of molecules to move toward areas of lower

concentration until the concentration is uniform throughout. During effusion, a gas escapes through a tiny hole in its container. Gases of lower mass move out more quickly than gases with higher masses Graham’s law of effusion states that the rate of effusion of a gas is

inversely proportional to the square root of the gas’s molar mass.

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Ex: You have hydrogen gas (H2) and oxygen gas (O2). That's root of 32/4

Study the Kinetic Molecular Theory for Ideal Gases

Reminders: ALWAYS convert any temperature given to you to degrees Kelvin. Otherwise you'll fail. Use pressure in ATM and Volume in Liters or else certain equations won't work. Concentration in moles, not grams or anything else.

Review Questions:

1. There are 11.2 Liters of an ideal gas at STP. The number of molecules of gas in this sample is:

a) 3.01 x 10^23 b) 6.02 x 10^23 c) 11.2 d) 22.4

2. A sample of gas has a volume of 50.0 L at a temperature of 300K. What temperature would be needed for this sample to have a volume of 60 L if its pressure remains constant?

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a) 400 K b) 87 C c) 360 C d) 250 K

3. Under the same conditions of temperature and pressure, how many more times faster will hydrogen effuse compared to ammonia?

4. 3.50g of CO2 and 11.6g of O2 are placed in a 3.00 L container at 273 K. What will the total pressure in the container be?

5. CaCO3 decomposes when heated to form CaO and CO2. What volume of CO2 can be produced from 37.0g CaCO3 at a temperature of 25 C and a pressure of 1.15 atm?

6. A 54.0 mL sample of oxygen is collected over water at 23 C and 770 torr pressure. What is the volume of the dry gas at STP? (Vapor pressure of Water at 23 C = 21.1 torr)

Unit 14Liquids and Solids

Key Terms:

Intermolecular Forces - Forces between molecules

Intramolecular Forces - Chemical bond between the atoms of molecules (bonds)

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Dipole-Dipole Attraction - Attraction between polar molecules

Hydrogen Bonding - Strong dipole-dipole attraction including a hydrogen atom

London Dispersion Forces - Forces between nonpolar molecules

Molar Heat of Fusion - Energy required to melt 1 mole of a substance

Molar Heat of Vaporization - energy required to change 1 mole of a liquid to vapor

Evaporation - Transition from liquid to vapor phase

Condensation - Transition from vapor to liquid phase

Vapor Pressure - Pressure of a vapor in equilibrium with a sample of liquid

Crystalline Solid - Solid with a regular arrangement of components

Alloy - Matter that contains a mixture of elements and has metallic properties

Chapter Summary:

Intermolecular Forces

Intermolecular forces are forces of attraction between particles Intermolecular forces are important in determining many macroscopic

properties of a substance Whereas intramolecular forces are inside of molecules, such as ionic

and covalent bonds, intermolecular forces occur on the outside There are four types of intermolecular forces. They are: ionic,

hydrogen, dipole-dipole, and London dispersion forces

Ionic Bonding

The strongest type of intermolecular forces Very high boiling point. NaCl is an example of this. Usually solid at room

temperature

Hydrogen Bonding

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Occurs between an H and a very electronegative atom (N, O, or F)

Above is an example of a hydrogen bond between two water molecules Hydrogen bonding affects the boiling point, usually making it higher The boiling point increases because the forces are stronger, and therefore

take more energy to break That is why water is liquid at room temperature, and not a gas

Dipole-Dipole Attractions

The force of attraction between opposite partial charge of neighboring dipoles

Hydrogen bonds are example of extremely strong dipole dipole attractions

London Dispersion Forces

The force of attraction between induced dipoles As the size of the atoms increase, the forces become stronger For these forces, an electrically neutral atom can create its own dipole Since I can't really explain it that well, here is a diagram

Phase Changes

Water and its phase changes; heating and cooling curve of water

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The density of water is 1g/mL. Ice happens to be less dense than water. Its boiling point is 100C and its freezing point is 0C It takes a certain amount of energy to vaporize a certain amount of water. It

also takes energy to melt it

Liquids

Vaporization and Evaporation are not the same thing. Vaporization is a process that occurs at the boiling point of a liquid, while evaporation can occur at any given temperature.

When the maximum amount of a liquid has evaporated at a given temperature, the liquid has reached equilibrium and will remain constant

Vapor pressure is the pressure of the vapor present at equilibrium with its liquid

At the boiling point, the atmospheric pressure is equal to the vapor pressure

Solids

There are ionic solids, molecular solids, and atomic solids Ionic = Sodium Chloride (NaCl - Na+ + Cl-). These have high melting

points and are held together by strong forces between ions Molecular = Water (H2O). These melt at generally low temperatures. Held

together by weak intermolecular forces Atomic = diamond (composed only of C atoms). Properties vary greatly Metals are held together by closely packed atoms, also known as the electron

sea model. Alloys are metals that have a number of different atoms present

Formulas and notes

*The molar heat of fusion and the molar heat of vaporization have formulas

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*The molar heat of fusion for water is 6.02 kJ/mol; the molar heat of vaporization is 40.6 kJ/mol (although you probably aren't expected to memorize either of these)

*Know the equation Q = sm t

*When you're using the equation, make sure to add all of the answers from the equations. For example, if the questions asks you to find the amount of heat needed to go from 50C to 150C, you must add the Q=sm t numerical answer to the numerical answer for the molar heat of vaporization.

*As always, be able to convert from grams to moles and moles to grams for these equations

Review Questions:

1. Which of the following substance most likely has the highest melting point?

a) NaCl b) H2O c) CH3OH d) Ne

2. Which of the following most likely has the lowest boiling point?

a) MgCl2 b) C12H22O11 c) H20 d) CS2

3. If the pressure above a liquid is increased, its boiling point will:

a) increase b) decrease c) stay the same d) depends on its properties

4. Which substance has the highest vapor pressure?

a) H2O b) H2S c) CH4 d) KCl

5. How much heat is necessary to completely vaporize 75 g of ice at 0C? The molar heat of fusion for water is 6.02 kJ/mol. The specific heat capacity of water is 4.18 J/g C. The molar heat of vaporization is 40.6 kJ/mol

UNIT 15

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Solutions

Key Terms:

Solution - Homogenous mixture in which the solute is dissolved in the solvent

Solvent - The major part of the homogenous mixture

Solute - The minor part of the homogenous mixture

Aqueous Solution - Homogenous mixture in which the solute is dissolved in water acting as the solvent

Saturated - A type of solution that contains as much solute that will dissolve at a given temperature

Unsaturated - A type of solution that has not reached the maximum amount of solute able to dissolve at a given temperature

Supersaturated - A type of solution that became saturated at an elevated temperature, then allowed to cool in a way that keeps all of the solid remained dissolved

Concentrated - A situation where there is more solute than solvent able to dissolve it in a 1:1 ratio

Dilute - A situation where there is more solvent than solute able to be dissolved in a 1:1 ratio

Mass Percent - The percentage of mass of solute compared to mass of whole solution

Molarity( M) - The amount of moles of a substance per liter(volume) of the substance

Standard Solution - A solution whose concentration is accurately known in a lab setting

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Dilution - A process in a lab setting to make the concentration of the solution lower by the addition of water acting as a solvent

Neutralization Reaction - A reaction involving an acid and a base

Equivalent of an Acid - The amount of acid that furnishes(creates) 1 mole of H+ ions

Equivalent of a Base - The amount of base that furnishes(creates) 1 mole of OH- ions

Equivalent Weight - The mass in grams of one equivalent of an acid or base respectively

Colligative Property - A property present in a solution that is based on the ratio of the number of solute particles to the number of solvent particles, rather than the type of particles present in the solution (ex: Boiling point elevation or freezing point depression)

Chapter Summary:

Solution Basics

• Solution = Solute + Solvent

• Polar substances dissolve other polar substance, while non-polar substances dissolve other non-polar substances. "Like dissolves Like"

• When a substance is dissolved, either intermolecular or intramolecular forces are broken and new bonds are formed.

Solubility of Ionic Substances

• Ionic substances best dissolve in aqueous solutions

• Ionic substances break up into individual anions and cations when dissolved due to the great difference in charges between the anions and cations, combined with the attractive pull of the polar and non-polar sides of the H2O molecule

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Ex: NaCl(s) + H20(l) --> (Na+) + (Cl-)

Solubility of Polar Substances

• Polar substances are molecules with a charge. H20 is the universal solvent due to its properties of being an extremely polar molecule.

• When a polar solid dissolves in water, it often has to breaks apart from the other intermolecular forces connecting it to the other molecules of the polar solid.

Ex: Sugar molecules bond together through hydrogen bonding. The Hydrogen bonds are broken during the dissolving process due to

the polar sides of the of the water molecule attracting different sides of the polar molecule.

• Most alcohols are soluble in water, due to their polarity caused by the OH bond.

Insolubility in Water

• Non-polar substances are insoluble in water. They are only soluble in other polar substances.

Ex: Oils are not soluble in water because they are non-polar. They are soluble in other non-polar substances such as paint

thinner.

Rate of Dissolving

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• Stirring a substance that is dissolving increases the rate of dissolving due to the increased speed and likelihood of molecules hitting each other.

• Grinding a solute before putting it in a solvent to be dissolved, increases the solids surface area, which therefore increases the speed of the dissolving. This is because the increased surface area allows for more chemical processes to occur at once.

• Heating during the process of dissolving a substance increases the rate of dissolution due to the fact that heated molecules move faster, which therefore allows for more chemical processes to occur. Vice versa is true if cooling takes place during the dissolving process.

General Facts

• Dissolving a solute in a solution will decrease the vapor pressure of the solution due to the lack polar molecules on the surface

• Adding salt to water increases the boiling point of the water

• Saturated solutions can be concentrated or dillute

Math Concepts

• Molarity = (moles/Liters)

• Mass Percent = (mass of solute/mass of solution) x 100% = [mass of solute/(mass of solute) + (mass of

solvent)] x 100%

• Moles of Solute = (M1)(V1) = (M2)(V2) (M1) = Molarity of initial solution (V1) = Volume of initial

solution (M2) = Molarity of diluted solution (V2) = Volume of diluted solution

Example Problems

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1.) 70.0 mL of .15M Ag(NO3) solution is poured into 140.0 mL of 0.10M MgCl2 solution. What mass of solid AgCl will be formed?

2.) 22.5g of MgCl2 is dissolved in enough water to make 500. mL of solution. The concentration of Cl(-) ions is what?

UNIT 16Acids and Bases

Key Terms:

Acids - A molecule with a hydrogen ion

Bases - A molecule with a hydroxide ion

Arrhenius Concept of Acids and Bases - Produces either an H+ or OH- for an acid or base respectively

Bronsted-Lowry Model - Acids and bases are proton donors and acceptors

Conjugate Acid - Molecule with one more hydrogen ion compared located on the product side of equation compared to the acid on the reactant side

Conjugate Base - Molecule with one less hydrogen ion on the product side of the equation compared to the base on the reactant side

Hydronium Ion - H30+

Completely Ionized - All parts of the products completely disassociated

Strong Acid -HCl, HBr, HI, H2(SO4), H(NO3), HClO4 (Completely ionized acids)

Strong Base - Base that completely dissociates

Dipotic Acid -Acid containing two hydrogen ions

Oxyacids -Acids with proton attached to an oxygen atom

Organic Acids -Hydrogen ion attached to a molecule with Carbon

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Carboxyl Group -Group of Carbon, Oxygen atoms, and Hydrogen ion found in organic acids

Amphoteric Substance - Substance that can be an acid or base (water)

Ion-Product Constant- [H+][OH-] = 1 x 10^-14

pH Scale - Scale used to indicate the acidity of a solution

Indicators - Things that indicate the acidity or how basic a solution based on color

Indicator Paper - Paper that changes color based on pH

Neutralization Reactant - The reactants of a reaction involving an acid and a base

Titration - Delivering a measured volume of a solution with a known concentration, into the solution being analyzed

Buret - Device used for accurate measurement of the delivery of a liquid

Stoichiometric Point (Equivalence Point) - When exactly enough standard solution (titrant) has been added to react with all the solution being analyzed

Titration Curve (pH Curve) - Data plotted on a graph for a given titration showing the pH of the solution vs volume

Buffered Solution - Resists a change in its pH when either an acid or base has been added. Presence of a weak acid and its conjugate base or weak base and conjugate acid can act as a buffer.

Neutralization Reaction

• HCl is the acid in this equation and H20 is acting as a base.

• H30(+), also known as hydronium, is acting as the conjugate acid in this equation. Cl- is acting as the conjugate base.

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• As shown above, HCl and Cl- form an Acid-Conjugate base pair, while H20 and H30+ form a Base-Conjugate acid pair

• In general, if a weak acid is on the reactant side of the equation, they will have a strong acid for a conjugate acid. If a strong acid is on the reactant side of the equation, its conjugate acid will be weak. These same rules can also be applied to bases and conjugate bases. That is why in the diagram above the strong acid HCl has a weak conjugate acid.

• The forward reaction dominates the example above because the reactant side contains a strong acid being HCl. If the reactant side contained a weak acid instead of a strong acid, the reverse reaction would dominate.

Solutions

• [H+] = Hydrogen molecule concentration, [OH-] = Hydroxide molecule concentration, Kw = 1 x 10^-14 = equilibrium constant for concentration of water

• [H+] > [OH-] is an acidic solution -[H+] < [OH-] is a basic solution - [H+] = [OH-] is a neutral solution

• [H+][OH-] = 1 x 10^-14 = Kw at 25 degrees Celsius

• pH = -log[H+] pOH = -log[OH-]

• pH + pOH = 14.00

Titrations

• M(a) = Molarity of Acid V(a) = Volume of Acid

• M(b) = Molarity of Base V(b) = Volume of Base

• [M(a)][V(a)] = [M(b)][V(b)]

Example Problems

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1.) If 60.0 mL of a HBr solution requires 35.0 mL of a 0.30 M NaOH to titrate it to the equivalence point, what is the concentration, in moles per liter, of the HBr solution?

2.) What is the pH of a 0.15 M H(PO4) solution?

UNIT 17Equilibrium

Key Terms:

Collision Model - The colliding of molecules is what is needed in order for a reaction to occur

Activation Energy - Energy required to be used in order for a reaction to start

Catalyst - A substance that speeds up a reaction without being consumed

Enzyme - Catalyst in a biological system

Homogenous Reaction - All reactants and products are in one phase

Heterogenous Reaction - Reactants are in two or more phases

Equilibrium - The exact balancing of two processes, one which is the opposite of the other

Chemical Equilibrium - A dynamic state where concentrations of all reactants and products remain constant. The rates of the reverse and forward reactions are exactly the same.

Law of Chemical Equilibrium - A dynamic reversible state in which rates of opposing processes(forward and reverse reactions) are equal

Equilibrium Expression - K = (Products) / (Reactants) = [(C)^c x (D)^d] / [(A)^a x (B)^b]

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Equilibrium Constant - The set ratio of products to constants, that stays this ratio at equilibrium

Equilibrium Position - Idea of a theoretical position in the system based off the value of the equilibrium concentration

Homogenous Equilibria - An equilibria system where the products and reactants are in the same state

Heterogeneous Equilibria - The products and reactants aren't in the same state in an equilibria system

Le Chatelier's Principle - When a change is imposed on a system at equilibrium, the position of the equilibrium shifts in a direction that tends to reduce the effect of that change

Solubility Product Constant (Ksp) - Measures the solubility of a solution using the equation Ksp = [M+][X-] with M and X as solids in a reversible reaction

Collision Model

• In order for a collision to successfully break a bond, and therefore trigger a reaction, there needs to be enough energy to break the bond, the molecule must collide, and wen the molecule collide they must be aligned properly.

• Due to the laws of the collision model, the rate of the chemical reaction is determined by the temperature of the system and the concentration of molecules within the system.

• The concentration of a system is determined by the pressure exerted on it and the volume it is allowed to encompass. Pressure and volume are inversely proportional, because as the volume gets bigger, the amount of pressure in the system goes down. The more pressure that is exerted on a system, the faster the rate of the reaction becomes.

• The temperature of a system also affects the rate of a reaction by directly affecting how fast molecules move. The higher the temperature the faster

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the molecules move, the faster the molecules move and also the more often molecules bump into each other. This results in more positive reactions between molecules to take place, and overall a faster reaction of the system. If the temperature of a system is cooler than normal, vice versa occurs in regards to the effects and rate of the overall reaction.

• Catalysts lower the activation energy needed for a positive reaction to occur.

Equilibrium Reactions

• They have the ability to move in both the forward and reverse direction. When both of the forward and reverse reactions are occurring at the same rate, equilibrium occurs.

• In accordance with Le Chatlier's principle, when a change is exerted on the system, the system shifts in the direction that will most effectively alleviate the effects of the change and bring back the system to equilibrium.

• The constant(K), resulting from an equilibrium expression of reaction, must always remain constant. It can do so by the equation shifting to balance the reactants and the products. The number of moles of each of the products and reactants may change, but the concentration will always stay the same in order for the reaction to remain at equilibrium. Solids and liquids aren't included in the expression, only gases and aqueous solutions.

• Adding more of a reactant to a system will cause there to be more products and a little bit less of the other reactants. Therefore, the reaction shifts right.

• Increasing the volume of a system will cause the reaction to shift towards the side where there is more moles of gas/aqueous solutions.

• Decreasing the volume of a system will cause the reaction to shift towards where there is less moles of gas/aqueous solutions.

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• Decreasing the pressure causes the reaction to shift towards where there is more moles of gas/aqueous solutions.

• Increasing the pressure causes the reaction to shift to where there is less moles of gas/aqueous solutions.

• In an equilibrium system, it always shifts away from where heat is added. In an exothermic equation, the temperature is decreased or created so heat shows up on the product(right) side. In an endothermic reaction, heat is used up in the reaction, so it shows up on the reactant(left) side of the reaction.

• If Heat = - H then the reaction is exothermic and shifts left.🔼

• If Heat = + H then the reaction is endothermic and shifts right.🔼

• If the temperature is "increased", then the reaction shifts left and is exothermic.

• If the temperature is "decreased", then the reaction shifts right and is endothermic.

Equilibrium Expressions

• K = Equilibrium Constant

• K never changes

• K > 1 means equilibrium position is far to the right

• K < 1 means equilibrium position is far to the left

• K = (Products) / (Reactants) = [(C)^c x (D)^d] / [(A)^a x (B)^b]

• Ksp = Solubility Product Constant

• Ksp = [M+][X-] with M and X as solids in an equilibrium expression

• If a coefficient is in front of one of the solid ions, it is incorporated in to the solubility product constant as an exponent.

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Example Problems

1.) Using the equation 5A(s) + 7B(g) <--> C(aq) + 4D(g) calculate equilibrium constant if [A] = .0025 M, [B] = .450 M, [C] = 1.2 M, and [D] = 0.30 M.

2.) Ksp for MgF2 = (3.7)10^-8 at 25 degrees Celsius. Calculate the solubility of MgF2 at this temperature.

EXAMPLE PROBLEMS: THE WRITTEN RESPONSE QUESTIONS

**Note - This portion is nearly 40% of the test. Each of these 5 question is 10x more important than each multiple choice. Here is an example of each.

1. (Electromagnetic Radiation Calculation) A wave of visible light has a wavelength of 5.93 x10^-7m. What is the energy of a single photon of this wave?

2. (Gas Law Calculation) Find the number of grams of CH4 that exert a pressure of 1.80 atm at a volume of 32.5L and a temperature of 50.0C?

3. (Titration Calculation) If 60 mL of a HCl solution requires 22.0 mL of 0.30 M NaOH to titrate it to the equivalence point, what is the concentration, in molarity, of the HCl solution?

4. (Solution Stoichiometry) 40.0 mL of 0.20 M AgNO3 solution is poured into 200.0 mL of 0.08 M MgCl2 solution. What mass of solid AgCl will be formed?

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5. (Balance Redox Reaction) Using the half reaction method, balance the following reaction. CeO2 + Sn+2 + H+ -----> Ce+3 + Sn+4 + 2H20