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404 Cards in this Set

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What is the basic building block of matter, representing the smallest unit of an element?
An atom
An atom is composed of...?
An atom is composed of subatomic particles called protons, neutrons, and electrons.
Nucleus
Protons and neutrons from nucleus. It is the core of the atom.
Electrons
Carry a charge equal in magnitude but opposite in sign to that of protons.
Protons
Carry a single positive charge and have a mass approximately one atomic mass unit or amu.

The atomic number (Z) of an element is equal to the number of protons found in an atom. All atoms of a given element have the same atomic number.
Atomic Number (Z)
Is equal to the number of protons found in an atom of that element.

All atoms of a given element have same atomic number.
Neutrons
Neutrons carry no charge and have a mass slightly larger than that of protons.

Different isotopes of one element have different numbers of neutrons, but same number of protons.

Mass number = total number of protons and neutrons.
Mass Number (A)
Equal to the number of protons and neutrons.

Convention = A/Z (X)
Valence Electrons
Electrons farthest from the nucleus. Farther they are from nucleus = the weaker the attractive force the positively charged nucleus and more likely these electrons are influenced by other atoms.

Generally, these electrons determine the reactivity of an atom.
Determine the number of protons, neutrons, and electrons in a Nickel-58 atom and in a Nickel-60 2+ cation.
Ni-58 has an atomic number of 28 and a mass number of 58. Therefore, Ni-58 will have 28 protons, 28 electrons, and 58-28, or 30, neutrons.

In Ni-60 2+ species, the number of protons is same as in neutral Ni-58. However, Ni-60 2+ has a positive charge because it has lost two electrons and thus, Ni 2+ will have 26 electrons. Therefore, mass number is higher by 2, which means neutrons = 32.
Avogadro's Number
A mole unit = 6.022 x 10^23 particles.

e.g. The atomic weight of carbon is 12.0 g/mol, which means that 6.022 x 10^23 carbon atoms weigh 12.0 g.
Isotopes
Same number of protons, but DIFFERENT number of NEUTRONS. (i.e. different mass numbers)
Element Q consists of three different isotopes: A, B, and C. Isotope A has an atomic mass of 40.00 amu and accounts for 60%, B is 44.00 amu and accounts for 25%, and C is 41 amu and accounts for 15.00%. What is the atomic weight of element Q?
Answer: 41.15 g/mol
Solution: 0.60(40 amu) + 0.25(44 amu) + 0.15(41 amu) = 41.15 amu
Quantum Theory
Developed by Max Planck, proposing that energy emitted as electromagnetic radiation from matter comes in discrete bundles called quanta.
Energy Value of a Quantum
E = hf

h = proportionality constant known as Planck's constant = 6.626 x 10^-34 J-s.

f = (sometimes designated v) = frequency of radiation.
Rydberg constant
Experimentally determined constant = 2.18 x 10^-18 J/electron.
Electron at Ground State
electron is in its lowest energy state.
Atomic Emission Spectra
Used as a fingerprint for the element.

Because each element can have its electrons excited to different distinct energy levels. When these electrons return to their ground states, each will emit a photon with a wavelength characteristic of the specific transition it undergoes.
Calculation of Electromagnetic energy of photons
E = hc/wavelength

h = Planck's constant
c = velocity of light (3.00 x 10^8 m/s)
wavelength = wavelength of radiation.
Atomic Absorption Spectra
When an electron is excited to a higher energy level, it must absorb energy. The energy absorbed as an electron jumps from an orbital of low energy to one of higher. Every element possesses a characteristic absorption spectra. (identify elements in a gas phase)
Orbital
Regions of space around the nucleus.

An orbital is a representation of the probability of finding an electron within a given region.
Heisenberg uncertainty principle
States that it is impossible to determine, with perfect accuracy, the momentum and position of an electron simultaneously.
Quantum numbers
An electron in an atom can be completely described by four quantum numbers: n, l, m(l), and m(s).
Pauli Exclusion Principle
No two electrons in a given atom can possess the same set of four quantum numbers.
Energy State
Is the position and energy of an electron described by its quantum numbers.
Principal Quantum Number
- First quantum number
- Denoted by letter n
- Information about the SIZE
- Larger the value of n = higher energy level.
- Maximum number of electrons in energy level n (electron shell n) is 2n^2.
Azimuthal Quantum Number
- second quantum number or angular momentum
- designated by letter l
- Information about the SHAPE
- Refers to the subshells or sublevels that occur within each principal energy level.
- For any given n, the value of l can be any integer in the range of 0 to n - 1.
- Four subshells corresponding to l = 0, 1, 2, and 3 are known as the s, p, d, and f subshells.
- Maximum number of electrons that can exist within a subshell is given by equation 4l + 2.
Magnetic Quantum Number
- Third quantum number.
- Denoted by m(l)
- Info about orientation of the orbital.
- An orbital is a specific region within a subshell that may contain no more than two electrons.
- Specifies the particular orbital within a subshell where an electron is highly likely to be found at a given point in time.
- Possible values of m(l) = integers from l to -l, including 0.
- s (0) = 1 orbital
- p (1) = 3 orbitals (-1, 0, 1)
- d (2) = 6 orbitals
- f (3) = 7 orbitals
Spin Quantum Number
- Fourth quantum number.
- Denoted by m(s)
- Two spin orientations are designated +(1/2) and -(1/2).
- Whenever two electrons are in the same orbital, they must have opposite spins.
- Electrons in different orbitals with same m(s) values have parallel spins.
Electron configuration
The pattern by which subshells are filled and the number of electrons within each principal level and subshell.

In electron configuration notatn, the first number denotes the principal energy level, the letter designates the subshell, and the superscript gives the number of electrons in that subshell.

e.g. 2p^4 = four electrons in the second (p) subshell of the second principal energy level.
Subshell fill chart
1s (start here and zig zag)

2s
2p
3s
3p
4s 3d
4p
5s 4d
5p 4f
6s 5d
6p 5f
7s 6d
7p (ends here)
(n + l) rule
Used to rank the subshells by increasing energy. Rule states that the lower values of the first and second quantum numbers, the lower the energy of the subshell.

Subshell with lower value will fill first.
Hund's Rule
States that within a given subshell, orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins.

electrons prefer empty orbitals to half-filled ones, because a pairing energy must be overcome.
What are the written electron configurations for Nitrogen (N) according to Hund's rule?
Nitrogen = 1s^2, 2s^2, 2p^3
It has an atomic number of 7. According to hund's rule, the two s orbitals will fill completely, while the p orbitals will each contain one electron, all with parallel spins.
Paramagnetic
material that has unpaired electrons.

A magnetic field will align the spins of these electrons and weakly attract the atom.
Diamagnetic
Materials that have NO unpaired electrons and are slightly repelled by a magnetic field.
Valence Electrons
Groups IA and IIA = outermost s electrons

Groups IIIA and VIIIA = outermost s and p electrons in highest energy shell

Transition elements = Outermost s and d subshell of next-to-outermost energy shell.

Inner transition elements = s subshell of the outermost energy shell, d subshell of next-to-outermost energy shell, and f subshell of energy shell two levels below outermost shell.
Periods
Rows

- 7 periods, representing the principal quantum numbers n = 1 to n =7, and each is filled sequentially.
Groups
Columns
- Represent elements that have the same electronic configuration in their valence, or outermost shell, and share similar chemical properties.
Roman Number above group
Represents the number of valence electrons.
Representative elements
- A elements, which have either s or p sublevels as outermost orbitals.
Nonrepresentative elements
B elements, including the transition elements, which have partly filled d sublevels.
Atomic Radius
Equal to one-half the distance between the centers of two atoms of that element.

Atomic radius decreases across a period from left to right and up a given group.

Atoms with largest atomic radii will be located at bottom of groups, and in Group 1.
Ionization Energy (IE)
or Ionization potential, is energy required to completely remove an electron from a gaseous atom or ion.

Removing an electron always requires input of energy.

Ionization energy increases from left to right across a period and up a group.
First and Second Ionization Energies
First ionization energy is energy required to remove one valence electron, the second is energy needed to remove a second valence electron.

Successive ionization energies grow increasingly large.
Electron Affinity
Energy change that occurs when an electron is added to a gaseous atom, and represents ease with which an atom can accept an electron.

noble gases have affinity of zero, because already possess an octet.

Increases left to right and up.
Electronegativity
Measure of attraction of an atom has for electrons in a chemical bond.

Greater the electronegativity, the greater its attraction for bonding electrons.

Increases from left to right across periods and up a group.
Three categories of Periodic Table
1. Metals (left side and middle)
2. Nonmetals (right side)
3. Metalloids or semimetals (along diagonal line between the two).
Characteristic of Metals
Shiny solids (except mercury) at room temp., and generally have high melting points and densities.

Ability to be deformed without breaking (malleability and ductility)

large atomic radius, low ionization energy, and low electronegativity.

Good conductors of heat and electricity.

Groups IA and IIA (most reactive metals)

transition (metals with partially filled d orbitals)
Characteristic of Nonmetals
Generally brittle in solid state and show little or no metallic luster.

High ionization energies and electronegativities, and are usually poor conductors of heat and electricity.

Gain electrons easily.

Located on upper right side of periodic table, partially filled p orbitals.
Characteristic of Metalloids
Semimetals found along line btw the metals and nonmetals.

Properties vary: densities, bp, mp fluctuate widely.

electronegativities and IE lie btw metals and nonmetals.

e.g. Silicon has metallic luster, yet is brittle and not efficient conductor (semiconductor).

Reactivity dependent upon element with which they are reacting.
Alkali Metals
Elements of Group IA

Possess properties of metals, yet densities lower than other metals.

Have one loosely bound electron in outermost, giving largest radius.
Alkaline Earths
Elements of Group IIA

Possess metallic properties, but dependent upon ease with which they lose electrons.

Have two electrons in valence.
Halogens
Elements of Group VIIA

Highly reactive nonmetals with seven valence electrons.

Highly variable in properties.
Noble gases
Also called inert gases.

Found in Group VIII (Also group O)

Fairly nonreactive because complete octet.

Low bp and all are gases at RT.
Transition elements
Elements of Groups IB to VIIIB

Considered metals.

Hard and have high melting points and boiling points.

Moving across period, the give d orbitals become progressively more filled.
Intermolecular forces
Forces between molecules. Although weaker than the intramolecular chemical bonds, these are of considerable importance in understanding physical properties of substances.
Octet Rule
States that an atom tends to bond with other atoms until it has eight electrons in its outermost shell, thereby forming stable electron configuration.

Exceptions are hydrogen, lithium and beryllium.
Types of chemical bonds
Ionic bonds and Covalent bonds
Ionic bonding
Electrons from an atom with smaller ionization energy is transferred to an atom with a greater electron affinity, and the resulting ions are held together by electrostatic forces. (Transfer of electrons from less to more) - difference needs to be greater than 1.7
Covalent bonding
Electron pair is shared between two atoms.

Atoms can share more than one pair of electrons. One, two, and three electron pairs are said to be joined by single, double, and triple bonds, respectively.

Can be characterized by bond length and bond energy.
Polar Covalent Bonds
Bond is partially covalent and partially ionic.
Cation
Atom that loses electrons become a positively charged ion.
Anion
Atom that gains electrons becomes a negatively charged ion.
Ionic Compounds
Have high melting and boiling points due to strong electrostatic forces.
Bond length
In reference to covalent bonds.

Average distance between two nuclei of atoms involved in the bond.

As number of electrons shared increases, the two atoms are pulled closer together, which decreases bond length.

Triple < Double < Single
Bond Energy
In reference to Covalent bonds.

Required to separate two bonded atoms.

Strength of bond increases as the number of shared electrons pairs increases.

Triple > Double > Single
Nonbonding Electrons
Valence electrons not involved in the covalent bond. (or lone electron pairs)
Bonding electrons
Shared valence electrons of a covalent bond.
Lewis Structure
Used to represent the bonding and nonbonding electrons in a molecule, facilitating "bookkeeping".
Formal Charges
Difference between the valence electrons and the number of electrons officially assigned to an atom in a Lewis Structure.

Formal charge = V - (1/2)Nbonding - Nnonbonding

V = valence electrons
Resonance Structures
Two or more nonidentical Lewis Structures. Molecule exists as composite of the resonance structures.

- Lewis structure with small or no formal charges is preferred over large formal charges.

- Lewis structure in which negative formal charges are placed on more electronegative is more stable than when placed on less E.
Types of Covalent Bonding
1. Polar Covalent Bond
2. Nonpolar Covalent Bond
3. Coordinate Covalent Bond
Polar Covalent Bond
Occurs between atoms with small differences in electronegativity, generally in the range of 0.4 to 1.7.

Bonding electron pair is NOT shared equally, but pulled more towards the element with higher electronegativity.
Nonpolar Covalent Bond
Occurs between atoms that have the same electronegativities.

Bonding electron IS shared equally, with no separation of charge across the bond.

Occur in diatomic molecules such as H2, Cl2, O2, and N2.
Coordinate Covalent Bond
Shared electron pair comes from the lone pair of one of the atoms in the molecule.

Typically found in Lewis acid-base compounds.
Lewis Acid
Compound that accepts an electron pair to form a covalent bond.
Lewis Base
Compound that can donate an electron pair to form a covalent bond.
Valence Shell Electron-Pair Repulsion Theory (VSEPR)
Uses Lewis structures to predict the molecular geometry of covalently bonded molecules.

States that the 3-D arrangement of atoms surrounding a central atom is determined by the repulsions between the bonding and nonbonding electron pairs in the valence shell of central atom. These electron pairs arrange themselves as far apart as possible to minimize repulsion.
Valence Electron Arrangements:

2, 3, 4, 5, 6 (atoms around central atom), what are electron arrangements?
2: linear, 180 degree
3: trigonal planar, 120 degree
4: tetrahedral, 109.5 degree or pyramidal
5: trigonal bipyramidal, 90-120-180 degrees
6: octahedral, 90-180 degrees
Polar molecule
Molecule with a net dipole moment.

Depends on the polarity of the constituent bonds and on the shape of the molecule.
S orbitals
Spherically symmetrical
P orbitals
Dumbbell shape
Bonding orbital
If the signs of the two atomic orbitals are the same.
Antibonding orbital
If the signs are different.
Sigma Bond
When orbitals overlap head to head.
Pi Bond
When the orbitals are parallel.
Types of Intermolecular Forces
Are attractive forces between molecules.

1. dipole-dipole interactions
2. hydrogen bonding
3. dispersion forces

dipole-dipole and dispersion are often referred to as van der waals forces.
Dipole-Dipole Interactions
Polar molecules tend to orient themselves such that the positive region of one molecule is close to the negative region of another molecule.

Are present in solid and liquid phases but become negligible in gas phase.

Tend to have higher boiling points than nonpolar species of comparable weiht.
Hydrogen Bonding
Specific, unusually strong form of dipole-dipole interaction, which may be either intra- or intermolecular.

Substances with hydrogen bonding tend to have high boiling points.

Important in behavior of water, alcohols, amines, and carboxylic acids.
Dispersion Forces
Short lived dipoles. Also called London forces.

Generally weaker than other intramolecular forces.

Do not extend over long distances and therefore most important when molecules are close together.
Compound
Pure substance that is composed of two or more elements in a fixed proportion.
Molecule
Combination of two or more atoms held together by covalent bonds.
Molecular Weight
Sum of the atomic weights (in amu) of the atoms in the molecule.
Formula Weight
Of an ionic compound is found by adding up atomic weights according to the empirical formula of the substance.
What is molecular weight of SOCL2?
To find molecular wieght of SOCl2, add together the atomic weights of each of the atoms:

1S = 1 x 32 amu = 32 amu
1O = 1 x 16 amu = 16 amu
2Cl = 2 x 35.5 amu = 71 amu

molecular weight = 119 amu
Mole
Defined as the amount of a substance that contains the same number of particles are found in a 12.000 g sample of carbon-12.

1 mole = 6.022 x 10^23 particles

mol = (weight of sample (g)) / (molar weight (g/mol))
Molar weight or molar mass
Expressed as g/mol

e.g. molar mass of H2CO3 = 62 g/mol
How many moles are in 9.52 g of MgCl2?
Answer = 0.10 mol of MgCl2

Solution:

First, find the molar mass of MgCl2.

1(24.31 g/mol) + 2(35.45 g/mol) = 95.21 g/mol

Now, solve: (9.52 g) / (95.21 g/mol) = 0.10 mol.
Equivalent Weight
Useful to define a measure of reactive capacity.

Equivalents = Weight of Compound / Gram Equivalent Weight

Gram Equivalent Weight = Molar Mass / n

n = either number of hydrogens used per molecule of acid in reaction, or the number of hydroxyl groups used per molecule of base in reaction.
Law of Constant Composition
States that any sample of a given compound will contain the same elements in the identical mass ratio.

e.g every sample of H2O will contain two atoms of hydrogen for every atom of oxygen, or in other words, one gram of hydrogen for every eight grams of oxygen.
Ways to express a formula for a compound
Empirical formula and molecular formula.
Empirical formula
Gives simplest whole number ratio of the elements in the compound.
Molecular formula
Gives the exact number of atoms of each element in the compound, and is usually a multiple of the empirical formula.
Percent composition
By mass of an element is the weight percent of the element in a specific compound.

To determine the percent composition of an element X in a compound, the following formula is used:

% comp = [(Mass of X in Formula) / (Formula Weight of Compound)] x 100%

May be determined with either empirical or molecular formulas.
What is the percent composition of chromium in K2Cr2O7?
Answer: 35.4%

Solution:
Formula Weight of K2Cr2O7 is;
2(39) + 2(52) + 7(16) = 294 g/mol

% comp of Cr = [2(52)]/294g/mol x 100%
= 0.354 x 100%
= 35.4%
What are the empirical and molecular formulas of a compound which contains 40.9% carbon, 4.58% hydrogen, 54.52% oxygen, and has a molecular weight of 264 g/mol?
Answer: Molecular C9H12O9
Empirical C3H4O3

Solution:

# mol C = [(0.409)(264)g] / 12 g/mol = 9 mol

# mol H = [(0.0458)(264)g] / 1 g/mol = 12 mol

# mol O = [(0.5452)(264)g] / 16 g/mol = 9 mol
Types of Chemical Reactions (Inorganic)
1. Combination reactions
2. Decomposition Reactions
3. Single Displacement Reactions
4. Double displacement reactions
Combination reactions
Reactions in which two or more reactants form one product.

e.g S + O2 --> SO2
Decomposition Reactions
One in which a compound breaks down into two or more substances, usually as a result of heating or electrolysis.

2 HgO + heat --> 2 Hg + O2
Single Displacement Reactions
Occurs when an atom (or ion) of one compound is replaced by an atom of another element.

e.g. Zn + CuSO4 --> Cu + ZnSO4
Double Displacement Reactions
Also called metathesis reactions, occurs when elements from two different compounds displace each other to form two new compounds.

CaCl2 + 2 AgNO3 --> Ca(NO3)2 + 2 AgCl
Net Ionic Equations
Because reactions such as displacements often involve ions in solution, they can be written in ionic form.

Important for demonstrating the actual reaction. Shows only the species that actually participate in reaction (minus spectator ions)
Spectator ions
Ions that do not take part in the overall reaction but simply remain in solution throughout.
Neutralization Reactions
Specific type of double displacements which occur when an acid reacts with a base to produce a solution of SALT and WATER.
Stoichiometric coefficients
Used to indicate the number of moles of a given species involved in the reaction.

Chemical equations must be balanced so that there are same number of atoms of each element in the products as there are in the reactants.
Applications of Stoichiometry
1. Limiting Reactant
2. Yields
Limiting reactant
First reactant to be consumed. It limits the amount of product that can be formed in the reaction.
Excess reactant
Reactant that remains after all the limiting reactant is used up.
If 28g of Fe react with 24 g of S to produce FeS, what wold be the limiting reagent? How many grams of excess reactant would be present in the vessel at the end of the reaction?

balanced eqn: Fe + S --> FeS
Answer: Fe is limiting reagent and 8 g of S will be in excess.

Solution:
First, the number of moles for each reactant must be determined.

28 g Fe x (1 mol Fe/56g) = 0.5 mol Fe

24 g S x (1 mol S/32g) = 0.75 mol

Fe is limiting.

Calculate excess:
mass of S = 0.25 mol S x (32 g/mol S) = 8 g of S
Yields
Yield of a reaction is the amount of product predicted or obtained when the reaction is carried out.

It can be determined or predicted from the balanced equation.
Three ways to report Yield
1. Theoretical yield
2. Actual Yield
3. Percent Yield
Theoretical Yield
Amount of product that can be predicted from a balanced equation, assuming all of the limiting reagent has been used.

Seldom obtained.
Actual Yield
Amount of product that is isolated from the reaction experimentally.
Percent Yield
used to express the relationship between the actual yield and the theoretical yield.

% yield = (actual yield/theoretical yield) x 100%
What is the percent yield for a reaction in which 27 g of Cu is produced by reacting 32.5 g of Zn in excess CuSO4 solution?
Answer: 84%

Solution:
The balanced equation is as follows:
Zn + CuSO4 --> Cu + ZnSO4

Calculate the theoretical yield for Cu:
32.5 g Zn x (1 mol Zn/65 g) = 0.5 mol

0.5 mol Zn x (1 mol Cu/1 mol Zn) = 0.5 mol Cu

0.5 mol Cu x (64 g /1 mol Cu) = theoretical yield (32 g)

% yield = (27g/32g) x 100% = 84%
Chemical Kinetics
Study of the rates of reactions, the effect of reaction conditions on these rates, and mechanisms implied by such observations.
Mechanism
Actual series of steps through which a chemical reaction occurs.
Rate-determining step
Slowest step in a proposed mechanism, because overall reaction cannot proceed faster than that step.
Rate Law
For general reaction:

a A + b B --> c C + d D

the rate is proportional to [A]^x[B]^y, that is:

rate = k[A]^x[B]^y

k = rate constant
x and y = orders of reaction
Reaction Order
Overall order of a reaction is defined as the sum of the exponents of rate law, equal to x + y.

Chemical reactions are often classified as zero-, first-, second-, mixed-, or higher-order reactions.
Zero-order reactions
Has a CONSTANT rate, which is INDEPENDENT of the reactant's concentrations.

Thus rate law is: rate = k

Units = Msec-1
First-order reactions
Order = 1
Has a rate proportional to the concentration of one reactant.

rate = k[A] or rate = k[B]

Units: sec-1
Second-order reactions
Order = 2
Has a rate proportional to the product of the concentration of two reactants, or to the square of the concentration of a single reactant.

e.g. rate = k[A]^2, rate = k[B]^2, or rate = k[A][B]

Units = M-1sec-1
Higher-order reactions
Has an order greater than 2.
Mixed-order reactions
Has a fractional order; e.g., rate = k[A]^1/3
Activation Energy
Or the energy barrier.

Is the minimum energy of collision necessary for a reaction to take place.
- Delta (H) = ?
Exothermic = heat given off
+ Delta (H) = ?
Endothermic = heat absorbed
Factors affecting Reaction Rate
The rate of a reaction will increase if either of the following occurs: an increase in number of effective collisions, or a stabilization of the activated complex to the reactants.

1. Reactant concentrations
2. Temperature
3. Medium
4. Catalysts
Reactant Concentrations in regards to Reaction Rate.
Greater the [reactants], the greater the number of effective collisions per unit time.

Therefore, the reaction rate will infrease for all but zero-order reactions.
Temperature and Reaction Rates
For nearly all reactions, the reaction rate will increase as the temperature of a system increases.
Medium and Reaction Rates
Rate of a reaction may also be affected by the medium in which it takes place.

Certain reaction proceed more rapidly in aqueous solution.
Catalysts and Reaction Rates
Catalysts are substances that increase reaction rate without themselves being consumed.

They do this by lowering activation energy. (e.g. enzymes)
What is the expression for the equilibrium constant for the following reaction?

2 H2 + N2 --> 2 NH3
Solution:

Kc = [NH3]^2 / ([H2]^3[N2])

Note: remember that exponents of the concentrations of teh reactants and products are equal to their stoichiometric coefficients in the equilibrium expression, but NOT kinetic rate law.
Reaction Quotient
Q, is the measure of the degree to which a reaction has gone to completion.

Qc = ([C]^c[D]^d) / ([A]^a[B]^b)

Is constant only at equilibrium, when it is equal to Kc.
Properties of equilibrium constant
1. Pure solids and liquids do NOT appear in the equilibrium constant expression.
2. Keq is characteristic of a given system at a given temperature.
3. If the value of Keq is very large compared to 1, an equilibrium mixture of reactants and products will contain very little of the reactants compared to products.
4. If it is very small compared to 1, it will contain very little product compared to reactant.
5. If Keq is close to 1, it will contain same amount of products and reactants.
Le Chatelier's Principle
Used to determine the direction in which a reaction at equilibrium will proceed when subjected to stress, such as change in concentration, pressure, temperature, or volume.
Le Chatelier's Principle: Change in Concentration
Increasing [] of a species will tend to shift the equilibrium away from the species that is added, in order to reestablish its equilibrium [].
Le Chatelier's Principle: Changes in Pressure or Volume
In a system at constant temperature, a change in pressure causes a change in volume, and vice versa.

Change in pressure or volume in liquids/solids will have little or NO effect.

Reactions involving gases, however, may be greatly affected by changes in P or V, since they are highly compressible.

Pressure and volume are inversely related. Increase in pressure will decrease moles of gas present (decreasing volume).
Le Chatelier's Principle: Changes in Temperature
Heat can be considered as a product (exothermic rxn) or as a reactant (endothermic rxn).

Depending on reaction, if temp was increased or decreased, it would alter the equilibrium.
Types of Systems
1. Isolated
2. Closed
3. Open
Isolated system
When it cannot exchange energy or matter with the surroundings, as with an insulated bomb reactor.
Closed system
When it CAN exchange ENERGY, but NOT matter with the surroundings, as with a steam radiator
Open System
When it can exchange BOTH matter and energy with the surroundings, as with a boiling pot of water.
Isothermal
Temperature of the system reamins constant.
Temperature and Reaction Rates
For nearly all reactions, the reaction rate will increase as the temperature of a system increases.
Medium and Reaction Rates
Rate of a reaction may also be affected by the medium in which it takes place.

Certain reaction proceed more rapidly in aqueous solution.
Catalysts and Reaction Rates
Catalysts are substances that increase reaction rate without themselves being consumed.

They do this by lowering activation energy. (e.g. enzymes)
What is the expression for the equilibrium constant for the following reaction?

2 H2 + N2 --> 2 NH3
Solution:

Kc = [NH3]^2 / ([H2]^3[N2])

Note: remember that exponents of the concentrations of teh reactants and products are equal to their stoichiometric coefficients in the equilibrium expression, but NOT kinetic rate law.
Reaction Quotient
Q, is the measure of the degree to which a reaction has gone to completion.

Qc = ([C]^c[D]^d) / ([A]^a[B]^b)

Is constant only at equilibrium, when it is equal to Kc.
Properties of equilibrium constant
1. Pure solids and liquids do NOT appear in the equilibrium constant expression.
2. Keq is characteristic of a given system at a given temperature.
3. If the value of Keq is very large compared to 1, an equilibrium mixture of reactants and products will contain very little of the reactants compared to products.
4. If it is very small compared to 1, it will contain very little product compared to reactant.
5. If Keq is close to 1, it will contain same amount of products and reactants.
Le Chatelier's Principle
Used to determine the direction in which a reaction at equilibrium will proceed when subjected to stress, such as change in concentration, pressure, temperature, or volume.
Le Chatelier's Principle: Change in Concentration
Increasing [] of a species will tend to shift the equilibrium away from the species that is added, in order to reestablish its equilibrium [].
Le Chatelier's Principle: Changes in Pressure or Volume
In a system at constant temperature, a change in pressure causes a change in volume, and vice versa.

Change in pressure or volume in liquids/solids will have little or NO effect.

Reactions involving gases, however, may be greatly affected by changes in P or V, since they are highly compressible.

Pressure and volume are inversely related. Increase in pressure will decrease moles of gas present (decreasing volume).
Le Chatelier's Principle: Changes in Temperature
Heat can be considered as a product (exothermic rxn) or as a reactant (endothermic rxn).

Depending on reaction, if temp was increased or decreased, it would alter the equilibrium.
Types of Systems
1. Isolated
2. Closed
3. Open
Isolated system
When it cannot exchange energy or matter with the surroundings, as with an insulated bomb reactor.
Closed system
When it CAN exchange ENERGY, but NOT matter with the surroundings, as with a steam radiator
Open System
When it can exchange BOTH matter and energy with the surroundings, as with a boiling pot of water.
Isothermal
Temperature of the system reamins constant.
A Sample of Argon occupies 50L at standard temperature. Assuming constant pressure, what volume will the gas occupy if the temperature is doubled?

a. 100L
b. 25L
c. 20L
d. 200L
e. none of the above
a. 100L
What is the best explanation for why Ca2+ has a larger atomic radius than Ar?

a. the calcium cation has less electrons.
b. the calcium cation is not a noble gas.
c. the calcium cation has more electrons.
d. the calcium cation has less protons.
e. the calcium cation has more protons.
d. the calcium cation has less protons.
Which of the following has the highest second ionization energy?

a. Ca
b. Sr
c. Ba
d. Ra
e. Na
e. Na
How many moles of water are produced in the combustion of 20g of glucose in excess O2?

a. 2/3
b. 1/3
c. 2/6
d. 1/6
e. none of the above
a. 2/3
A zero order reaction has teh form

a. r = 0
b. r = k[A]
c. r = k[A][B]
d. r = [A]^0
e. r = k
e. r = k
A reaction is always spontaneous if....

a. change in gibbs free energy is zero
b. change in gibbs free energy is positive.
c. change in gibbs free energy is negative
d. change in enthalpy is negative
e. change in entropy is negative
a. change in gibbs free energy is zero
Entropy is best defined as

a. degree of chaos
b. degree of conduction
c. amount of heat exchanged in a system at equilibrium
d. degree of order
e. amount of disorder in a system
e. amount of disorder in a system
When a substance absorbs heat, it is always true that:

a. change in entropy is negatve
b. change in enthalpy is negative
c. change in enthalpy is positive
d. change in gibbs free energy is negative
e. change in gibbs free energy is positive
c. change in enthalpy is positive
When an impurity enters a solution, which of the following will occur?

a. the solution's vapor pressure will increase.
b. the solution's boiling point will increase.
c. the solution's freezing point will increase.
d. the solution will have a lower osmotic force
e. all of the above
b. the solution's boiling point will increase.
If a capillary has a higher osmotic force than the interstitial fluid surrounding it, then which of the following will occur?

a. water will flow into the capillary
b . water will flow into the interstitial fluid
c. water will flow out of the capillary
d. it depends on the fluid
e. all of the above
a. water will flow into the capillary
Which of the following is an assumption of the kinetic molecular theory of gases?

a. particles have negligible volume
b. no intermolecular attraction
c. random motion, continuous movement of molecules
d. average kinetic energy of gas particles is proportional to the absolute temperature
e. all of the above.
e. all of the above.
Charle's Law
(V1/T1) = (V2/T2)

V/T (Volume of gas is directly proportional to its absolute temperature)
Boyle's Law
PV = k or P1V1 = P2V2

Volume of gas is proportional to its pressure.
Ideal Gas Law
PV = nRT

Combines the relationships outlined in Boyle's law, Charles' Law, and Avogadro's Principle to yield an expression which can used to predict the behavior of a gas.
Heat
Form of energy which can easily transfer to or from a system, the result of a temperature difference between the system and its surroundings.

Heat can be endothermic (asborbed) or exothermic (released).

Measured in calories (cal), or Joules (J), and more commonly in kcal or kJ.
Calorimetry
Measures heat changes.

Heat (q) absorbed or released in a given process is calculated from the equation:

q = mc(delta)T

m = mass
c = specific heat
delta T = the change in temperature
Standard conditions
25 degrees Celsius and 1 atm.

Normally used to measure enthalpy, entropy, and free energy of a reaction.
Enthalpy
Expression of heat changes at constant pressure.

Delta H = a process is equal to the heat absorbed or evolved by the system at constant pressure.
Delta H
To find enthalpy change, one must substract the enthalpy of the reactants from the enthalpy of the products.

delta H(rxn) = H products - H reactants

Positive H = endothermic
Negative H = exothermic
Standard Heat of Formation
Enthalpy formation of a compound.

Delta H-f = is the enthalpy change that would occur if one mole of compound were formed directly from its elements in their standard states.
Standard Heat of Reaction
Standard Heat of Reaction, Delta H(rxn), is the hypothetical enthalpy change that would occur if the reaction were carried out under standard conditions.

delta H(rxn) = (sum of delta H(formation) of products) - (sum of Delta H (formation) of reactants)
Hess' Law
States that enthalpies of reactions are additive.

When thermochemical equations are added to give the net equation for a reaction, the corresponding heats of reaction are also added to give the net heat of reaction.
Heat of vaporization
Enthalpy change of a liquid to gas, Delta H (vap).
Bond Dissociation Energy
An average of the energy required to break a particular type of bond in one mole of gaseous molecules.

Bond breakage is always ENDOTHERMIC, because it takes energy to pull two atoms apart.

Bond energies can be used to estimate enthalpies of reactions.

Delta H(rxn) = (Delta H bonds broken) - (Delta H bonds formed)

= total energy input - total energy released

formation of bonds = negative
breakage of bonds = positive
Heats of Combustion
Delta h(comb).

Reaction is often fast and spontaneous.
Entropy
(S) is a measure of DISORDER, or randomness, of a system.

Units: energy/temperature, commonly J/K or cal/K

The greater the order in a system, the lower the entropy and vice versa.

Delta S = S final - S initial
Delta S (Standard Entropy Change)
Calculated using the standard entropies of reactants and products.

Delta S(rxn) = (sum of S products) - (sum of S reactants)

Freezing = decrease in entropy
Boiling = increase in entropy
Sublimation is greatest entropy change.

System will spontaneously tend towards an equilibrium state if left alone.
Gibbs Free Energy
(Delta G) = (Delta H) - T(Delta S)

"Goose Hunters Take Shotguns"

Combines the two factors which affect the spontaneity of a reaction - enthalpy and entropy.

Delta G = change in free energy of a system. Represents the maximum amount of energy released by a process, occurring at constant T and P.
Characteristics of Gibbs Free Energy
1. If Delta G = NEGATIVE, the reaction is SPONTANEOUS

2. If Delta G = POSITIVE, the reaction is NOT SPONTANEOUS.

3. If Delta G = ZERO, the system is in a state of equilibrium.

thus Delta G = 0 and (Delta H) = T (Delta S)

(Delta H),(Delta S):

(-,+) = SPONTANEOUS at ALL temperatures
(+,-) = NONSPONTANEOUS at ALL temperatures
(+,+) = SPONTANEOUS only at HIGH temperatures
(-,-) = SPONTANEOUS only at LOW temperatures.
Standard Free Energy
Delta G, is defined as the Delta G of a process occurring at 25 degrees Celsius and 1 atm pressure, and for which the concentrations of any solutions involved are 1M.
Reaction Quotient
Delta G(rxn) can also be derived from the equilibrium constant for the equation:

Delta G = -RT ln Keq

Keq = equilibrium constant,
R = gas constant
T = Temperature (K)
Three Different Forms of Matter
Gas, liquid, Solid
Four variables that define a gaseous samples
pressure (P), volume (V), temperature (T), and number of moles (n).

Pressure usually expressed in terms of atmospheres (atm)0 or millimeters of mercury (mm Hg or torr)

1 atm = 760 mm Hg = 760 torr

Volume in Liters or milliliters.

Temperature in Kelvin.
Standard Temperature and Pressure
273.15 K (0 degrees Celsius) and 1 atm.
Boyle's Law
At constant temperature, the volume is inversely proportional to its pressure:

PV = k or P1V1 = P2V2

k is proportionality constant.
Charles Law
At constant pressure, the volume of gas is directly proportional to its absolute temperature.

V/T = k or (V1/T1) = (V2/T2)
Avogadro's Principle
Proposed that for all gases at constant temperature and pressure, the volume of the gas will be directly proportional to the number of moles of gas present.

n/v = k or n1/V1 = n2/V2
Ideal Gas Law
Combines the relationships outlined in Boyle's, Charles', and Avogadro's Principle to yield an expression which can be used to predict the behavior of a gas.

PV = nRT
R = 8.22 x 10-2 L-atm/(mol-K)
What volume would 12 g of helium occupy at 20 degrees C and a pressure of 380 mm Hg?
Solution:

P = 380 mm Hg x (1/760 mm Hg) = 0.5 atm

T = 20 degrees C + 273.15 = 293.15 K

n = 12 g He x 1 mol He / 4.0 g = 3 mol He

(0.5atm)(V) = (3 mol)(0.0821)(293.15K)

V = 144.4L
Density
Defined as mass per unit volume of a substance and, for gases is usually expressed in unites of g/L.

d = m/v = P(MM)/RT
Calculating change in Volume
V2 = V1(P1/P2)(T2/T1)

V2 is then used to find the density of the gas unders nonstandard conditions.

d = m / V2
Molar Mass
Sometimes identity of gas is unknown, and molar mass must be determined.

Find weight of sample by volume to calculate the density. And MW is found by multiplying the number of grams per liter by 22.4 L per mole.
What is the molar mass of a 2L sample of gas that weights 8 g at a temperature of 15 degrees C and a pressure of 1.5 atm?
d = 8g / 2L at 15 degrees C and 1.5 atm.

V(STP) = (2L)(273/288)(1.5/1) = 2.84 L

8g / 2.84L = 2.82 g/L at STP

(2.82 g/L)(22.4 L/mol) = 63.2 g/mol
Dalton's Law of Partial Pressures
Pressure exerted by each individual gas is called the partial pressure of that gas.

Law states that the total pressure of a gaseous mixture is equal to the sum of the partial pressures of the individual components. The equation is...

P(T) = P(A) + P(B) + P(C) + ....

The partial pressure of a gas is related to its mole fraction and can be determined using the following equations: PA = P(T)X(A)

X(A) = n(A)/n(T) (moles of A/total moles)
A vessel contains 0.75 mol of nitrogen, 0.20 mol of hydrogen, and 0.05 mol of fluorine at a total pressure of 2.5 atm. What is the partial pressure of each gas?
First calculate the mole fraction of each gas.

X(N2) = 0.75 mol/1mol

X(H2) = 0.20 mol/1mol

X (F2) = 0.05 mol/1mol

Then calculate the partial pressure:

P(A) = X(A)P(T)

P(N2) = (2.5 atm)(0.75) = 1.875 atm

P(H2) = (2.5atm)(0.20) = 0.5 atm

P(F2) = (2.5atm)(0.05) = 0.125 atm
Assumptions of the Kinetic Molecular Theory of Gases
1. Gases are made up of particles whose volumes are negligible compared to the container volume.
2. Gas atoms or molecules exhibit no intermolecular attractions or repulsions.
3. Gas particles are in continuous, random motion, undergoing collisions with other particles and the container walls.
4. Collisions between any two gas particles are elastic, meaning that there is no gain or loss of energy.
5. The average kinetic of gas particles is proportional to the absolute temperature of the gas, and is the same for all gases at a given temperature.
Graham's Law of Diffusion and Effusion
Diffusion: Occurs when gas molecules diffuse through a mixture. r1/r2 = (MM2/MM1)^(1/2) = square root of (MM2/MM1)

r1 and MM1 represent the diffusion rate and molar mass of gas 1.

Effusion: flow of gas particles under pressure from one compartment to another through a small opening.

r1/r2 = (MM2/MM1)^(1/2)
Condensed Phases
Liquid or solid phases. Because of their smaller volume relative to gases.
Evaporation
Or Vaporization is a cooling process. Liquid to Gas.
Condensation
Gas to liquid.
Vapor Pressure
Pressure that the gas exerts over the liquid.

Vapor pressure increases as temperature increases.

Temperature at which the vapor pressure of the liquid equals the external pressure is called boiling point.
Boiling point
Temperature at which the vapor pressure of the liquid equals the external pressure.
Fusion
Or melting, solid to liquid.

Temperature at which this occurs is melting point.
Solidification/freezing
Or crystallization is liquid to solid.

Temperature at which this occurs is freezing point.
Sublimation
Solid to gas (e.g dry ice)
Triple Point
Intersection of the three lines is called triple point. At this temperature and pressure, unique for a given susbtance,....

All three phases are in equilibrium.
Critical Point
Tip between liquid and gas phases.

Temperature and pressure above which no distinction between liquid and gas is possible.
Colligative Properties
Physical properties derived solely from the number of particles present, not the nature of those particles.
Freezing-Point Depression
Pure water freezes at 0 degrees Celsius.

Formula for calculating freezing-point depression:

Delta Tf = K(f)m

Delta Tf = freezing point depression

K(f) = proportionality constant

m = molality of the solution (mol solute/kg solvent)
Boiling-Point Elevation
A liquid boils when its vapor pressure equals the atmospheric pressure.

Delta T(b) = K(b)m

Delta T(b) = boiling-point elevation

Kb = proportionality constant

m = molality of the solution.
Osmotic Pressure
Substances tend to diffuse from higher to lower concentrations.

O = MRT

M = molarity
R = ideal gas constant
T = Temperature in Kelvin
Raoult's Law
Vapor-Pressure Lowering

When solute B is added to pure solvent A, the vapor pressure of A above the solvent decreases.

If the vapor pressure of A above pure solvent A is designated P-(A) and the vapor pressure of A above the solution containing B is P(A), the vapor pressure decreases as follows:

Delta P = P-(A) - P(A)
Dissolution
Interaction between solute and solvent molecules.
Percent Composition by Mass
Is the mass of the solute divided by the mass of the solution, multiplied by 100.
Sublimation
Solid to gas (e.g dry ice)
Triple Point
Intersection of the three lines is called triple point. At this temperature and pressure, unique for a given susbtance,....

All three phases are in equilibrium.
Critical Point
Tip between liquid and gas phases.

Temperature and pressure above which no distinction between liquid and gas is possible.
Colligative Properties
Physical properties derived solely from the number of particles present, not the nature of those particles.
Freezing-Point Depression
Pure water freezes at 0 degrees Celsius.

Formula for calculating freezing-point depression:

Delta Tf = K(f)m

Delta Tf = freezing point depression

K(f) = proportionality constant

m = molality of the solution (mol solute/kg solvent)
Boiling-Point Elevation
A liquid boils when its vapor pressure equals the atmospheric pressure.

Delta T(b) = K(b)m

Delta T(b) = boiling-point elevation

Kb = proportionality constant

m = molality of the solution.
Osmotic Pressure
Substances tend to diffuse from higher to lower concentrations.

O = MRT

M = molarity
R = ideal gas constant
T = Temperature in Kelvin
Raoult's Law
Vapor-Pressure Lowering

When solute B is added to pure solvent A, the vapor pressure of A above the solvent decreases.

If the vapor pressure of A above pure solvent A is designated P-(A) and the vapor pressure of A above the solution containing B is P(A), the vapor pressure decreases as follows:

Delta P = P-(A) - P(A)
Dissolution
Interaction between solute and solvent molecules.
Percent Composition by Mass
Is the mass of the solute divided by the mass of the solution, multiplied by 100.
Arrhenius Acid
Species that produces H+ (a proton) in an aqueous solution.
Arrhenius Base
Species that produces OH- (a hydroxide ion) in an aqueous solution.
Bronsted-Lowry Acid
Species that DONATES protons.
Bronsted-Lowry Base
Species that ACCEPTS protons.

e.g. NH3 and Cl-

Not limited to aqueous solutions.
Conjugate acid-base pairs
Bronsted-Lowry acids and bases always occur in pairs.

The two members of a conjugate pair are related by the transfer of a proton.

e.g. H3O+ is conjugate acid of the base H2O
Lewis Acid
Electron-pair ACCEPTOR.

e.g. BCl3 and AlCl3
Lewis Base
Electron-Pair DONOR.
Measurement of pH
pH = -log[H+] = log(1/[H+])
Measurement of pOH
pOH = -log[OH-] = log (1/[OH-])
Water dissociation Constant
Kw = [H+][OH-] = 10^-14

An equilibrium reaction.

Rewrite:

pH + pOH = 14
Fundamental Properties of Logarithms
1. log(xy) = log x + log y

log of a product is equal to the sum of the logs.

2. Log x^n = nLog x, and Log 10^x = x
If [H+] = 0.001, what is the pH?
0.001 = 10^-3, then pH = 3
If Kb = 1.0 x 10^-7, then what is the pKb?
1.0 x 10^-7, then pKb = 7
If Ka = 1.8 x 10^-5, then what is pKa?
pKa = 5 - log 1.8.

Since 1.8 is small, its log will be small, and the answer will be closer to 5 than to 4. (The actual answer is 4.74)
Strong acids and bases
Those that completely dissociate into their component ions in aqueous solution.

e.g. NaOH + excess H2O --> Na+ + OH-

Strong acids: HNO3, H2SO4, HClO4, and HCl

Strong bases: KOH, NaOH, and other soluble hydroxiddes of Group IA and IIA metals.
Weak Acids and Bases
Those that partially dissociate in aqueous solution.

Weak monoprotic acid, HA, in aqueous solution will achieve the following eqn:

HA + H2O --> H3O + A
Acid dissociation Constant
Ka = ([H3O+][A-]) / [HA]

The weaker the acid, the smaller the Ka.
Base dissociation Constant
A weak monovalent base, BOH, undergoes dissociation to give B+ and OH-.

Measure of the degree to which a base dissociates.

Weaker the base, the smaller the Kb.

Kb = ([B+][OH-) / [BOH]
Conjugate acid
Defined as the acid formed when a base gains a proton.
Conjugate base
Formed when an acid loses a proton.

e.g. HCO3-/CO3^2- (conjugate acid/base pair)
Ka x Kb = Kw = 1.0 x 10^-14
Remember this.
Find the concentration of H+ in a 2.0 M aqueous solution of acetic acid, CH3COOH (Ka = 1.8 x 10^-5).
1. Write the equilibrium reaction:

CH3COOH --> H + CH3COO-

2. Write the expression for the acid dissociation constant:

Ka = ([H+][CH3COO-]) / [CH3COOH] = 1.8 x 10^-5

Since acetic acid is a weak acid, the concentration of CH3COOH at equilibrium is equal to its initial concentration, 2.0M, less the amount dissociated, x.

Likewise [H+] = [CH3COO-] = x, since each molecule of CH3COOH dissociates into one H+ ion and one CH3COO- ion. Thus equation is rewritten:

Ka = ([x][x]) / [2.0 - x] = 1.8 x 10^-5

We can approximate that 2.0 - x is approximately 2.0 is a weak acid. Simplifies equation:

Ka = [x][x] / [2.0] = 1.8 x 10^-5

x = 6.0 x 10^-3 M
Neutralization Reaction
Acids and bases may react with each other, forming a salt and (often, but not always) water, in a neutralization reaction.
Hydrolysis
Reverse reaction of a neutralization reaction. Salt ions react with water to give back acid or base.
Titration
Procedure used to determine the molarity of an acid or base.

Accomplished by reacting a known volume of solution of unknown concentration with a known volume of solution of known concentration.
Equivalence point
When number of acid equivalents equals the number of base equivalents added, or vice versa.
Buffer Solution
Consists of a mixture of a weak acid and its salt (which consists of its conjugate base and a cation), or a mixture of a weak base and its salt (which consists of its conjugate acid and an anion).
Henderson-Hasselbach equation
used to estimate the pH of a solution in the buffer region where the concentrations of the species and its conjugate are present in approximately equal concentrations.
Calculating pH for weak acid buffer solution
pH = pKa + log ([conjugate base][weak acid])

When [conjugate base] = [weak acid], pH = pKa
Calculating pOH for weak base buffer solution
pOH = pKb + log ([conjugate acid]/[weak base])

When [conjugate acid] = [weak base], pOH = pKb
- log Ka = ?

- log (5.6 x 10^-11) = ?
- log Ka = pKa

- log (5.6 x 10^-11) = 11 - log 5.6
Electrochemical Reactions
Include spontaneous reactions that produce electrical energy, and nonspontaneous reactions that use electrical energy to produce a chemical change. Both types always involve a transfer of electrons with conservation of charge and mass.
Law of Conservation of Charge
An electrical charge can be neither created NOR destroyed. Thus, an isolated loss or gain of electrons cannot occur.
Oxidation
Loss of electrons
Reduction
Gain of Electrons
Redox reaction
An electron transfer.

"OIL RIG"

Oxidation is LOST and Reduction is GAIN
Oxidizing agent
Causes another atom in a redox reaction to undergo oxidation, and is itself reduced.
Reducing agent
Causes the other atom to be reduced, and is itself oxidized.
Assigning Oxidation Numbers
1. The oxidation number of free elements is ZERO. (e.g. N2, P4, S8.)

2. Oxidation number for a monatomic ion is EQUAL to the charge of the ion. (e.g. Na+, Cu2+, Fe3+, Cl-, and N3- are +1, +2, +3, -1, and -3, respectively)

3. The oxidation number of each Group IA element in a compound is +1. The oxidation number of each Group IIA element in a compound is +2.

4. The oxidation number of each Group VIIA element in a compound is -1, except when combined with an element of higher electronegativity. (e.g. HCl, the oxidation number of Cl- is -1; in HOCl, however, the oxidation number of Cl is +1).

5. The oxidation number of hydrogen is -1 in compounds with less electronegative elements than hydrogen (Groups IA and IIA). (e.g. NaH and CaH2. The more common oxidation number of hydrogen is +1)

6. In most compounds, the oxidation number of oxygen is -2. This is not the case, however, in OF2. Here, because F is more electronegative, the oxidation number of O is +2. Also, peroxides such as BaO2, the oxidation number of O is -1.

7. The sum of the oxidation numbers of all atoms present in a neutral compound is zero. The sum of the oxidation numbers of the atoms present in a polyatomic ion is equal to the charge of the ion. Thus, for SO4^2-, the sum of the oxidation numbers must be -2.
Oxidation Numbers
Are assigned to atoms in order to keep track of the redistribution of electrons during a chemical reaction.
Assign oxidation numbers to the atoms in the following reaction in order to determine the oxidized and reduced species and the oxidizing and reducing agents.

SnCl4 + PbCl4 --> SnCl4 + PbCl2
Solution:

All of these species are neutral, so the oxidation numbers of each compound must add up to ZERO.

In SnCl2, since there are two chlorines present, and chlorine has an oxidation number of -1, Sn must have an oxidation number of +2.

Similarly, the oxidation number of Sn in SnCl4 is +4.

Pb is +4 in PbCl4 and +2 in PbCl2.

Notice the oxidation number of Sn goes from +2 to +4; it loses electrons and thus is oxidized, making it the reducing agent.

Since oxidation number of Pb has decreased from +4 to +2, it has gained electrons and been reduced. Pb is oxidizing agent.

The sum of charges on both sides of the reaction is equal to zero, so charge has been conserved.
Half-reaction method
Most common method for balancing redox reactions. Also known as ion-electron method.

Equation is separated into two half-reactions - the oxidation part and the reduction part.

Each half reaction is balanced separately, and they are then added to give a balanced reaction.
Balance this redox reaction:

MnO4^- + I^- --> I2 + Mn^2+
Step 1. Separate the two half-reactions:

I^- --> I2

MnO4^- --> Mn^2+

Step 2. Balance the atoms of each half reaction. First, balance all atoms except H and O. next, in an acidic solutiono, add H2O to balance the O atoms and then add H+ to balance the H+ atoms. (In a basic solution, use OH- and H2O to balance the O and H).

To balance the iodine atoms, place a coefficient of two before the I^- ion.

2 I^- --> I2

For the permanganate half-reaction, Mn is already balanced. Next, balance the oxygens by adding 4H2O to the right side.

Finally, add H+ to the left side to balance the 4 H2Os. These two half-reactions are now balanced.

MnO4^- + 8 H+ --> Mn^2 + 4 H2O

Step 3. Balance the charges for each half reaction. The reduction half must consume the same number of electrons as are supplied by the oxidation half. For the oxidation reaction, add 2 electrons to the right side of the reaction:

2 I^- --> I2 + 2e-

For the reduction reaction, a charge of +2 must exist on both sides. Add 5 electrons to the left side of the reaction to accomplish this:

5e- + 8 H+ + MnO4^- --> Mn^2+ + 4H2O

Next, both half reactions must have the same number of electrons so that they will cancel. Multiply the oxidation half by 5 and the reduction half by 2.

5(2I^- --> I2 + 2e-)
2 (5e- + 8 H+ + MnO4^- --> Mn^2+ + 4 H2O)

Step 4: Add half-reactions:

10 I^- --> 5 I2 + 10e-
16 H+ + 2 MnO4^- + 10e- --> 2 Mn^2+ + 8 H2O

The final equation is:

10 I^- + 10e- + 16 H+ + 2 MnO4^- --> 5 I2 + 2 Mn^2+ + 10 e- + 8 H2O

To get overall reaction, cancel out electrons and any H2Os, H+, and OH- that appear on both sides:

10 I- + 16 H+ + 2 MnO4^- --> 5 I2 + 2 Mn^2+ + 8 H2O

Step 5: Confirm mass and charge are balanced. There is +4 net charge on each side.
Electrochemical Cells
Contained systems in which a redox reaction occurs.

Two types: galvanic and electrolytic cells.

Both types contain electrodes at which oxidation and reduction occurs.

For all electrochemical cells, the electode at which oxidation occurs is called the anode, and reduction occurs at the cathode.

"AN OX and a RED CAT"
Galvanic Cells
Spontaneous reactions occur.

Redox reaction occurring in a galvanic cell has a negative delta G and is therefore a SPONTANEOUS reaction.

Commonly used as batteries.
Electrolytic Cells
Nonspontaneous reactions occur.

Has a positive Delta G and is therefore NONSPONTANEOUS.
Rules for Constructing a Cell Diagram:
1. Reactants and products are always listed from left to right in the form:

anode - anode solution -- cathode solution - cathode

2. Single vertical line indicates phase boundary.

3. A double vertical line indicates the presence of a salt bridge or some other type of barrier.
Anode of Electrolyic Cells
Positive

Since attached to the positive pole of the battery and so attracts anion from the solution.
Anode of Galvanic Cell
NEGATIVE

Because the spontaneous oxidation reaction that takes place at the galvanic cell's anode.
Cathode of Electrolytic Cells
NEGATIVE
Cathode of Galvanic Cell
Positive

But in both types oxidation occurs at anode and reduction occurs at cathode.
Standard Electromotive Force (EMF or E(cell))
Difference in potential between two half-cells. EMF is determined by adding the standard reduction potential of reduced and standard oxidation of the oxidized.

EMF = E(red) + E(ox)

Standard EMF of galvanic cells is POSITIVE and the standard EMF of electrolytic is NEGATIVE.
Relation between EMF and Delta G
Delta G = -nFE(cell)

n = number of moles of electrons exchanged
F = Faraday's constant
E cell = EMF of the cell.

Keep in mind that if Faraday's constant is expressed in coulombs (J/V), then delta G must be expressed in J, not kJ.
Nernst Equation
E(cell) = E(cell) - (RT/nF)(lnQ)

Q = reaction quotient for a given reaction.

e.g. reaction quotient for:

a A + b B --> c C + d D

Q = ([C]^c[D]^d) / ([A]^a[B]^b)
Relation between EMF and the Equilibrium Constant (Keq)
Delta G = -RT ln Keq

Combining equations:

Delta G = -nFE = -RT ln Keq

OR

nFE = RT ln Keq
Mass defect
E = mc^2
Nucleons
Protons and neutrons
Isotope
For a nucleus of a given element with a given number of protons, the various nuclei with different numbers of neutrons.
Fusion
Combine smaller nuclei into larger nucleus.
Fission
Split larger nucleus into smaller nuclei
Alpha Decay
A = -4 and Z = -2

Emission of an alpha particle, which is a 4 He nucleus that consists of 2 protons and 2 neutrons.

alpha = 4/2 He
Beta Decay
A = no change, Z = -1 (positron) or +1

Emission of a beta-particle, which is an electron given the symbol e-.

B- decay =
A stays the same, Z + 1

B+ decay = positron
A stays the same, Z - 1
Gamma Decay
NO CHANGE

Emission of gamma-particles, which are HIGH energy photons. Carry NO charge and simply lower the energy of the emitting nucleus without changing the mass number or the atomic number.

Daughters A and Z is same as parents.
Calculating Half-life
(t 1/2) =

(1/2)^n

n = half-lives.
If the half-life of a certain isotope is 4 years, what fraction of a sample of that isotope will remain after 12 years?
Solution:

If 4 years is one half-life, then 12 years is three half-lives.

Thus the fraction remaining after 3 half-lives is (1/2)^3 or (1/8)
Alpha Decay involves release of/loss of....
Loss of a CATIONIC helium
A lewis acid, by definition
Accepts an electron pair
A Bronsted-Lowry base, by definition....
Donates H+
In beta- decay, the mass number (A)
Remains unchanged
In beta- decay, the atomic number (Z)...
increased by one
In Beta+ (positron) decay, the atomic number (Z)...
Decreases by ONE
In gamma decay, the atomic number and mass number
Remain unchanged
In electron capture, the atomic number (Z)...
Decreases by one
Electron Capture
Decreases atomic number by 1 and mass number remains the same.

Rare process that is perhaps best thought of as an inverse of B- decay.

Captures an electron and forms a neutron.
Exponential Decay
Delta n / Delta t = -wavelength(n)

wavelength = decay constant
n = number that reamin
t = time
The first law of thermodynamics is best summed in the two words:

a. energy constant
b. energy changing
c. energy increasing
d. energy decreasing
e. disorder constant
a. energy constant
The second law of thermodynamics is best summed in the 6 words:

a. spontaneous processes experience an increase entropy.
b. spontaneous processes experience a decreased entropy.
c. disorder is constant and quiescent always.
d. disorder occurs during cooling processes always.
e. disorder is proportional to increased entropy.
a. Spontaneous processes experience an increase entropy.
What is the closest value to Faraday's constant?

a. 9500 C
b. 96,500 C
c. 96,499 C
d. 96,501 C
e. 96,450 C
C. 96,499 C

Faraday constant = 96,485.34 C/mole
Charge of a single electron
1.602 x 10^-19 Coulombs
Units of Coulomb
J/V

(Joules/Volts)
Units of Current
C/S

(Coulombs/Second)
If you were to add more of the same solvent to a container of the same solvent, the vapor pressure of the solvent above the liquid would....
Remain unchanged.
In a galvanic Cells, the anode is labeled...
Negative
In an electrolytic cells, the cathode is labeled...
Negative
Electrolytic cells differ from galvanic cells in that they....
Electrolytic cells are NONSPONTANEOUS and contain the anode and cathode in the same solution.

GALVANIC cells are SPONTANEOUS and half-reactions take place in SEPARATE containers.
The formula for the Gibbs Free Energy Change for a Redox reaction is....
Delta G = -nFE
The Values of the Spin Quantum number are directly dependent on....

a. the principle quantum number
b. the angular momentum quantum number
c. the magnetic quantum number
d. none of the above
e. all of the above
d. none of the above

spin is always -1/2 and +1/2, opposite to each other.
(n + l) rule
Indication of which subshell is lower in energy and will fill first.

e.g. 3d, n = 3 and l = 2, so (n + l) = 5
An electron with 26 protons, and is in its ground state, will have how many electrons in its 3d orbital?
6 electrons
What is the maximum number of electrons that can be present in an atom with a principal quantum number of 5?
50
Which of the following is most likely to increase the rate of a reaction?

a. decrease in T
b. increase in T
c. decrease in pressure
d. increase in pressure
e. two of the above
b. increase in T
Decreasing the activation energy of the forward reaction will do what to the reverse reaction?

a. decrease the rate
b. increase the rate
c. have no effect
d. effect the equilibrium
e. make the reaction reach equilibrium slower
b. increase the rate.
If Keq << 1, then which of the following statements is true?

a. there are way more reactants than products.
b. the products most likely outnumber the reactants.
c. there are way more products than reactants.
d. the reactants do not outnumber the products.
e. none of the above.
a. there are way more reactants than products.

Keq = products/reactants.
Which of the following will effect Keq?

a. decrease in T
b. increase in T
c. decrease in P
d. A & B
e. A & C
d. A and B
For a system at equilibrium, which of the following statements is false?

a. the system is at maximum entropy
b. the forward reaction will continue.
c. the reverse reaction will continue.
d. the change in entropy of the system is zero
e. none of the above
e. none of the above.
What is the Normality of a 2.0 M solution of phosphoric acid, H3PO4, for an acid-base titration?
N = 6

N = molarity x n

n = number of hydrogens exchanged.

N = (2.0moles/L) x (3 equivalents H3SO4/L) = 6 N
How many moles of water are produced in the combustrion of 20 g of glucose in excess O2?
Answer 2/3

Solution:
First, write out balanced equation:

C6H12O6 + 6 O2 --> 6 CO2 + 6 H2O

Now solve for moles of H2O:

20g glucose x (1mol/180g glucose) x (6 mol H2O/1 mol glucose) = 120/180 = 2/3
A zero order reaction has the form...
r = k
A reaction is always spontaneous if....

a. change in gibbs free energy is zero
b. change in gibbs free energy is positive
c. change in gibbs free energy is negative
d. change in enthalpy is negative
e. change in entropy is negative
c. change in gibbs free energy is NEGATIVE
Maximum entropy is seen in a system when which of the following is occurring?

a. change in gibbs free energy is zero
b. change in gibbs free energy is positive
c. change in gibbs free energy is negative
d. change in enthalpy is negative
e. change in entropy is negative
a. change in gibbs free energy is ZERO.
Entropy is best defined as:
Amount of disorder in a system
When a substance absorbs heat, it is always true that

a. change in entropy is negative
b. change in enthalpy is negative
c. change in enthalpy is positive
d. change in gibbs free energy is negative
e. change in gibbs free energy is positive
c. change in enthalpy is positive.
When an impurity enters a solution, which of the following will occur?

a. the solution's vapor pressure will increase.
b. the solution's boiling poing will increase
c. the solution's freezing point will increase
d. the solution will have a lower osmotic force
e. all of the above.
b. the solution's boiling point will increase
Titration of a weak acid with a strong base will result in a slightly basic solution due to which of the following events:

a. a mistake occurred and too much base was added.
b. buffering
c. the base used was too strong
d. the acid was too weak
e. the acid was too strong
b. buffering
When H2CO3 is allowed to sit in solution, which of the following will you observe:

a. O2 formation
b. CO2 formation
c. H2O formation
d. A & B
e. B & C
e. B & C

H2CO3 --> H2O and CO2
Dalton's law of partial pressure states that the total pressure of all gases in a system:

a. is less than the sum of the partial pressures.
b. is more than the sum of the partial pressures.
c. is equal to the sum of the partial pressures of the gases.
d. is not related to the partial pressures of the gases.
e. none of the above
c. is equal to the SUM of the partial pressures of the gases.

P(T) = P(A) + P(B) +...
Define specific gravity?
The ratio of density of a substance in any units to the density of water in the same units.
Define anhydride
A molecule formed by the removal of hydrogen and oxygen from Water to form another molecule.
Which molecule is soluble?

a. AgCl
b. PbCl2
c. HgBr2
d. HgCl2
e. none of the above
e. none of the above
Which of the following molecules is soluble?

a. all molecules containing CO2
b. all molecules containing hydroxides
c. all molecules containing ammonium ions
d. all molecules containing hydrogen
e. none of the above
c. all molecules containing ammonium ions.
Sulfate is not soluble in water when combined with which of the following?

a. Ca
b. Ba
c. Sr
d. B and C only
e. all of the above
e. all of the above.
Fusion
Melting

Solid to liquid
Crystallization
Freezing

Liquid to Solid
Vaporization
Evaporation

Liquid to gas
Condensation
Gas to liquid.
Triple point
All three phase are in equilibrium at this T and P. And is where all three phases are present at the same time.
Deposition
Gas to Solid
Electron Motive Force (EMF)
Difference in potential between two half cells. Adding reduction potential of the reduced to the oxidation.

EMF = E(red) + E(ox)

standard EMF for galvanic cell is POSITIVE

standard for electrolytic is NEGATIVE

Higher E (reduction potential) = reduced

Lower E = oxidized
If the atomic weight of carbon-12 is exactly 12 amu, what is the mass of a single carbon-12 atom?
2.0 x 10^-23

12 g/mol x (1 mol/6.022x10^-23) = 2.0 x 10 ^-23
What is the percent yield if 78 g of C6H6 reacts and 82 g of C6H5NO2 is formed according to the reaction below?

C6H6 + HNO3 --> C6H5NO2 + H2O
Answer: 67%

Percent Yield = (Actual Yield/Theoretical Yield) x 100%

Actual Yield = 82 g
Theoretical Yield =
78g C6H6 x (1mol/78g C6H6) x (1mol C6H5NO2/1mol C6H6) x (123 g/mol C6H5NO2) = 123 g

% yield = 82/123 = 0.667 x 100 = 67%
Hydrogen Fluoride, HF, is a gas at room temperature, while water, H2O, is a liquid. The best explanation for this observed difference in physical properties is that:

a. the difference in MW indirectly accounts for the difference in boiling points because of van der Waals forces.
b. the O-H bond dipoles of water are greater than the F-H bonds dipoles of HF and account for the greater dipole dipole interactions btween water molecules.
c. hydrogen bonding between water molecules is significantly greater than that between HF molecules.
d. Dispersion forces are significant btween HF molecules, but not between water molecules.
e. water, as a universal solvent, is more likely than HF to be contaminated with impurities.
C. Hydrogen bonds are a specific type of dipole-dipole interaction. Water is liquid at RT because each molecule can form two hydrgen bonds with neighboring molecules. HF can form only one intermolecular hydrogen bond.
Ka < 1, then pKa = ?
pKa > 1

smaller Ka = weaker the acid
therefore, the bigger the pKa the weaker the acid.
Freezing-Point Depression
Delta T(f) = K(f)m

K(f) = proportionality constant characteristic of a particular solvent
m = molality of the solution (mol solute/kg solvent)

If compound is ionic, the formula for the freezing point depression has to be multiplied by the number of particles formed upon dissolving.
When carbon-14 undergoes beta decay, the daughter element is:

a. carbon-12
b. carbon-13
c. nitrogen-14
d. oxygen-15
e. silicon-28
c. nitrogen-14

In beta decay, a neutron decays into a proton and an electron. The generic for beta decay is A = 0, Z = +1. Since Carbon-14 has atomic number of 6. Add 1 to Z and it will be 7, which is nitrogen. Mass number does not change.
The balanced equation below is for a spontaneous oxidation-reduction reaction.

8 Al + 3 NO3 + 5 OH + 18 H2O --> 8 Al(OH)4 + 3 NH3

Which of the following is the BEST oxidizing agent?

a. Al
b. NO3-
c. NH3
d. Al(OH)4
e. OH-
b. NO3-

The best oxidizing agent is the species getting reduced. The nitrogen goes from +5 in NO3- to -3 in NH3, getting reduced and serving as an oxidizing agent.

A is wrong because Al goes from 0 to +3.
C is wrong because they represent products of rxn.
E is wrong because neither O2 nor Hydrogen in OH are changed in rxn.
Isoelectric species
Atoms and ions that have the same number of electrons.
Characteristics of Covalent Bonds
1. Electrons are shared between the elements.
2. Elements bonded together are usually two nonmetals.
NH3(aq) + HCl(aq) --> ?
Acid base reaction.

Product is NH4Cl
BaCO3(s) + HCl(aq) --> ?
Gas forming reaction.

Product is H2O + CO2 + BaCl2

Sum of coefficents is 6

Balanced net reaction:
BaCO3 + 2 H --> H2O + CO2 + Ba2+
NaF(aq) + Ba(NO3)2(aq) --> ?
Net Ionic Eqation for reaction:

2F- + Ba^2- --> BaF2
Ionic Compounds
Made up of cations and anions.
-ous
represents lesser charge
-ic
represents greater charge
Fe2+
Ferrous
Fe3+
Ferric
Cu2+
Cuprous
Cu3+
Cupric
-ide
Monatomic elements are named by dropping ending of the name of element and adding -ide
H-
Hydride
F-
Fluoride
O^2-
Oxide
S^2-
Sulfide
N^3-
Nitride
P^3-
Phosphide
Oxyanions
Polyatomic anions that contain OXYGEN.
-ite
When an element forms two oxyanions, then name of the one with LESS oxygen ends in -ite.
-ate
When an element forms two oxyanions, then name of the one with MORE oxygen ends in -ate.
NO2-
Nitrite
NO3-
NItrate
SO3^2-
Sulfite
SO4^2-
Sulfate
Hypo-
When the series of oxyanions contain FOUR oxyanions, then prefixes used.

LESS Oxygen = Hypo-
Per-
When the series of oxyanions contain FOUR oxyanions, then prefixes used.

MORE oxygen = Per-
ClO-
Hypochorite
ClO2-
Chlorite
ClO3--
Chlorate
ClO4
Perchlorate
Percent composition by mass
[Mass of solute/mass of the solution (solute + solvent)] x 100
Hess' Law
States that enthalpies of reactions are additive.