Dat - General Chemistry

Front 1

What is the basic building block of matter, representing the smallest unit of an element?

Back 1

An atom

Front 2

An atom is composed of...?

Back 2

An atom is composed of subatomic particles called protons, neutrons, and electrons.

Front 3

Nucleus

Back 3

Protons and neutrons from nucleus. It is the core of the atom.

Front 4

Electrons

Back 4

Carry a charge equal in magnitude but opposite in sign to that of protons.

Front 5

Protons

Back 5

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.

Front 6

Atomic Number (Z)

Back 6

Is equal to the number of protons found in an atom of that element.

All atoms of a given element have same atomic number.

Front 7

Neutrons

Back 7

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.

Front 8

Mass Number (A)

Back 8

Equal to the number of protons and neutrons.

Convention = A/Z (X)

Front 9

Valence Electrons

Back 9

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.

Front 10

Determine the number of protons, neutrons, and electrons in a Nickel-58 atom and in a Nickel-60 2+ cation.

Back 10

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.

Front 11

Avogadro's Number

Back 11

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.

Front 12

Isotopes

Back 12

Same number of protons, but DIFFERENT number of NEUTRONS. (i.e. different mass numbers)

Front 13

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?

Back 13

Answer: 41.15 g/mol
Solution: 0.60(40 amu) + 0.25(44 amu) + 0.15(41 amu) = 41.15 amu

Front 14

Quantum Theory

Back 14

Developed by Max Planck, proposing that energy emitted as electromagnetic radiation from matter comes in discrete bundles called quanta.

Front 15

Energy Value of a Quantum

Back 15

E = hf

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

f = (sometimes designated v) = frequency of radiation.

Front 16

Rydberg constant

Back 16

Experimentally determined constant = 2.18 x 10^-18 J/electron.

Front 17

Electron at Ground State

Back 17

electron is in its lowest energy state.

Front 18

Atomic Emission Spectra

Back 18

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.

Front 19

Calculation of Electromagnetic energy of photons

Back 19

E = hc/wavelength

h = Planck's constant
c = velocity of light (3.00 x 10^8 m/s)
wavelength = wavelength of radiation.

Front 20

Atomic Absorption Spectra

Back 20

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)

Front 21

Orbital

Back 21

Regions of space around the nucleus.

An orbital is a representation of the probability of finding an electron within a given region.

Front 22

Heisenberg uncertainty principle

Back 22

States that it is impossible to determine, with perfect accuracy, the momentum and position of an electron simultaneously.

Front 23

Quantum numbers

Back 23

An electron in an atom can be completely described by four quantum numbers: n, l, m(l), and m(s).

Front 24

Pauli Exclusion Principle

Back 24

No two electrons in a given atom can possess the same set of four quantum numbers.

Front 25

Energy State

Back 25

Is the position and energy of an electron described by its quantum numbers.

Front 26

Principal Quantum Number

Back 26

- 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.

Front 27

Azimuthal Quantum Number

Back 27

- 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.

Front 28

Magnetic Quantum Number

Back 28

- 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

Front 29

Spin Quantum Number

Back 29

- 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.

Front 30

Electron configuration

Back 30

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.

Front 31

Subshell fill chart

Back 31

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)

Front 32

(n + l) rule

Back 32

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.

Front 33

Hund's Rule

Back 33

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.

Front 34

What are the written electron configurations for Nitrogen (N) according to Hund's rule?

Back 34

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.

Front 35

Paramagnetic

Back 35

material that has unpaired electrons.

A magnetic field will align the spins of these electrons and weakly attract the atom.

Front 36

Diamagnetic

Back 36

Materials that have NO unpaired electrons and are slightly repelled by a magnetic field.

Front 37

Valence Electrons

Back 37

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.

Front 38

Periods

Back 38

Rows

- 7 periods, representing the principal quantum numbers n = 1 to n =7, and each is filled sequentially.

Front 39

Groups

Back 39

Columns
- Represent elements that have the same electronic configuration in their valence, or outermost shell, and share similar chemical properties.

Front 40

Roman Number above group

Back 40

Represents the number of valence electrons.

Front 41

Representative elements

Back 41

- A elements, which have either s or p sublevels as outermost orbitals.

Front 42

Nonrepresentative elements

Back 42

B elements, including the transition elements, which have partly filled d sublevels.

Front 43

Atomic Radius

Back 43

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.

Front 44

Ionization Energy (IE)

Back 44

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.

Front 45

First and Second Ionization Energies

Back 45

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.

Front 46

Electron Affinity

Back 46

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.

Front 47

Electronegativity

Back 47

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.

Front 48

Three categories of Periodic Table

Back 48

1. Metals (left side and middle)
2. Nonmetals (right side)
3. Metalloids or semimetals (along diagonal line between the two).

Front 49

Characteristic of Metals

Back 49

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)

Front 50

Characteristic of Nonmetals

Back 50

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.

Front 51

Characteristic of Metalloids

Back 51

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.

Front 52

Alkali Metals

Back 52

Elements of Group IA

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

Have one loosely bound electron in outermost, giving largest radius.

Front 53

Alkaline Earths

Back 53

Elements of Group IIA

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

Have two electrons in valence.

Front 54

Halogens

Back 54

Elements of Group VIIA

Highly reactive nonmetals with seven valence electrons.

Highly variable in properties.

Front 55

Noble gases

Back 55

Also called inert gases.

Found in Group VIII (Also group O)

Fairly nonreactive because complete octet.

Low bp and all are gases at RT.

Front 56

Transition elements

Back 56

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.

Front 57

Intermolecular forces

Back 57

Forces between molecules. Although weaker than the intramolecular chemical bonds, these are of considerable importance in understanding physical properties of substances.

Front 58

Octet Rule

Back 58

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.

Front 59

Types of chemical bonds

Back 59

Ionic bonds and Covalent bonds

Front 60

Ionic bonding

Back 60

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

Front 61

Covalent bonding

Back 61

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.

Front 62

Polar Covalent Bonds

Back 62

Bond is partially covalent and partially ionic.

Front 63

Cation

Back 63

Atom that loses electrons become a positively charged ion.

Front 64

Anion

Back 64

Atom that gains electrons becomes a negatively charged ion.

Front 65

Ionic Compounds

Back 65

Have high melting and boiling points due to strong electrostatic forces.

Front 66

Bond length

Back 66

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

Front 67

Bond Energy

Back 67

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

Front 68

Nonbonding Electrons

Back 68

Valence electrons not involved in the covalent bond. (or lone electron pairs)

Front 69

Bonding electrons

Back 69

Shared valence electrons of a covalent bond.

Front 70

Lewis Structure

Back 70

Used to represent the bonding and nonbonding electrons in a molecule, facilitating "bookkeeping".

Front 71

Formal Charges

Back 71

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

Front 72

Resonance Structures

Back 72

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.

Front 73

Types of Covalent Bonding

Back 73

1. Polar Covalent Bond
2. Nonpolar Covalent Bond
3. Coordinate Covalent Bond

Front 74

Polar Covalent Bond

Back 74

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.

Front 75

Nonpolar Covalent Bond

Back 75

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.

Front 76

Coordinate Covalent Bond

Back 76

Shared electron pair comes from the lone pair of one of the atoms in the molecule.

Typically found in Lewis acid-base compounds.

Front 77

Lewis Acid

Back 77

Compound that accepts an electron pair to form a covalent bond.

Front 78

Lewis Base

Back 78

Compound that can donate an electron pair to form a covalent bond.

Front 79

Valence Shell Electron-Pair Repulsion Theory (VSEPR)

Back 79

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.

Front 80

Valence Electron Arrangements:

2, 3, 4, 5, 6 (atoms around central atom), what are electron arrangements?

Back 80

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

Front 81

Polar molecule

Back 81

Molecule with a net dipole moment.

Depends on the polarity of the constituent bonds and on the shape of the molecule.

Front 82

S orbitals

Back 82

Spherically symmetrical

Front 83

P orbitals

Back 83

Dumbbell shape

Front 84

Bonding orbital

Back 84

If the signs of the two atomic orbitals are the same.

Front 85

Antibonding orbital

Back 85

If the signs are different.

Front 86

Sigma Bond

Back 86

When orbitals overlap head to head.

Front 87

Pi Bond

Back 87

When the orbitals are parallel.

Front 88

Types of Intermolecular Forces

Back 88

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.

Front 89

Dipole-Dipole Interactions

Back 89

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.

Front 90

Hydrogen Bonding

Back 90

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.

Front 91

Dispersion Forces

Back 91

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.

Front 92

Compound

Back 92

Pure substance that is composed of two or more elements in a fixed proportion.

Front 93

Molecule

Back 93

Combination of two or more atoms held together by covalent bonds.

Front 94

Molecular Weight

Back 94

Sum of the atomic weights (in amu) of the atoms in the molecule.

Front 95

Formula Weight

Back 95

Of an ionic compound is found by adding up atomic weights according to the empirical formula of the substance.

Front 96

What is molecular weight of SOCL2?

Back 96

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

Front 97

Mole

Back 97

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))

Front 98

Molar weight or molar mass

Back 98

Expressed as g/mol

e.g. molar mass of H2CO3 = 62 g/mol

Front 99

How many moles are in 9.52 g of MgCl2?

Back 99

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.

Front 100

Equivalent Weight

Back 100

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.

Front 101

Law of Constant Composition

Back 101

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.

Front 102

Ways to express a formula for a compound

Back 102

Empirical formula and molecular formula.

Front 103

Empirical formula

Back 103

Gives simplest whole number ratio of the elements in the compound.

Front 104

Molecular formula

Back 104

Gives the exact number of atoms of each element in the compound, and is usually a multiple of the empirical formula.

Front 105

Percent composition

Back 105

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.

Front 106

What is the percent composition of chromium in K2Cr2O7?

Back 106

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%

Front 107

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?

Back 107

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

Front 108

Types of Chemical Reactions (Inorganic)

Back 108

1. Combination reactions
2. Decomposition Reactions
3. Single Displacement Reactions
4. Double displacement reactions

Front 109

Combination reactions

Back 109

Reactions in which two or more reactants form one product.

e.g S + O2 --> SO2

Front 110

Decomposition Reactions

Back 110

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

Front 111

Single Displacement Reactions

Back 111

Occurs when an atom (or ion) of one compound is replaced by an atom of another element.

e.g. Zn + CuSO4 --> Cu + ZnSO4

Front 112

Double Displacement Reactions

Back 112

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

Front 113

Net Ionic Equations

Back 113

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)

Front 114

Spectator ions

Back 114

Ions that do not take part in the overall reaction but simply remain in solution throughout.

Front 115

Neutralization Reactions

Back 115

Specific type of double displacements which occur when an acid reacts with a base to produce a solution of SALT and WATER.

Front 116

Stoichiometric coefficients

Back 116

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.

Front 117

Applications of Stoichiometry

Back 117

1. Limiting Reactant
2. Yields

Front 118

Limiting reactant

Back 118

First reactant to be consumed. It limits the amount of product that can be formed in the reaction.

Front 119

Excess reactant

Back 119

Reactant that remains after all the limiting reactant is used up.

Front 120

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

Back 120

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

Front 121

Yields

Back 121

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.

Front 122

Three ways to report Yield

Back 122

1. Theoretical yield
2. Actual Yield
3. Percent Yield

Front 123

Theoretical Yield

Back 123

Amount of product that can be predicted from a balanced equation, assuming all of the limiting reagent has been used.

Seldom obtained.

Front 124

Actual Yield

Back 124

Amount of product that is isolated from the reaction experimentally.

Front 125

Percent Yield

Back 125

used to express the relationship between the actual yield and the theoretical yield.

% yield = (actual yield/theoretical yield) x 100%

Front 126

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?

Back 126

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%

Front 127

Chemical Kinetics

Back 127

Study of the rates of reactions, the effect of reaction conditions on these rates, and mechanisms implied by such observations.

Front 128

Mechanism

Back 128

Actual series of steps through which a chemical reaction occurs.

Front 129

Rate-determining step

Back 129

Slowest step in a proposed mechanism, because overall reaction cannot proceed faster than that step.

Front 130

Rate Law

Back 130

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

Front 131

Reaction Order

Back 131

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.

Front 132

Zero-order reactions

Back 132

Has a CONSTANT rate, which is INDEPENDENT of the reactant's concentrations.

Thus rate law is: rate = k

Units = Msec-1

Front 133

First-order reactions

Back 133

Order = 1
Has a rate proportional to the concentration of one reactant.

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

Units: sec-1

Front 134

Second-order reactions

Back 134

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

Front 135

Higher-order reactions

Back 135

Has an order greater than 2.

Front 136

Mixed-order reactions

Back 136

Has a fractional order; e.g., rate = k[A]^1/3

Front 137

Activation Energy

Back 137

Or the energy barrier.

Is the minimum energy of collision necessary for a reaction to take place.

Front 138

- Delta (H) = ?

Back 138

Exothermic = heat given off

Front 139

+ Delta (H) = ?

Back 139

Endothermic = heat absorbed

Front 140

Factors affecting Reaction Rate

Back 140

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

Front 141

Reactant Concentrations in regards to Reaction Rate.

Back 141

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.

Front 142

Temperature and Reaction Rates

Back 142

For nearly all reactions, the reaction rate will increase as the temperature of a system increases.

Front 143

Medium and Reaction Rates

Back 143

Rate of a reaction may also be affected by the medium in which it takes place.

Certain reaction proceed more rapidly in aqueous solution.

Front 144

Catalysts and Reaction Rates

Back 144

Catalysts are substances that increase reaction rate without themselves being consumed.

They do this by lowering activation energy. (e.g. enzymes)

Front 145

What is the expression for the equilibrium constant for the following reaction?

2 H2 + N2 --> 2 NH3

Back 145

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.

Front 146

Reaction Quotient

Back 146

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.

Front 147

Properties of equilibrium constant

Back 147

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.

Front 148

Le Chatelier's Principle

Back 148

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.

Front 149

Le Chatelier's Principle: Change in Concentration

Back 149

Increasing [] of a species will tend to shift the equilibrium away from the species that is added, in order to reestablish its equilibrium [].

Front 150

Le Chatelier's Principle: Changes in Pressure or Volume

Back 150

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).

Front 151

Le Chatelier's Principle: Changes in Temperature

Back 151

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.

Front 152

Types of Systems

Back 152

1. Isolated
2. Closed
3. Open

Front 153

Isolated system

Back 153

When it cannot exchange energy or matter with the surroundings, as with an insulated bomb reactor.

Front 154

Closed system

Back 154

When it CAN exchange ENERGY, but NOT matter with the surroundings, as with a steam radiator

Front 155

Open System

Back 155

When it can exchange BOTH matter and energy with the surroundings, as with a boiling pot of water.

Front 156

Isothermal

Back 156

Temperature of the system reamins constant.

Front 157

Temperature and Reaction Rates

Back 157

For nearly all reactions, the reaction rate will increase as the temperature of a system increases.

Front 158

Medium and Reaction Rates

Back 158

Rate of a reaction may also be affected by the medium in which it takes place.

Certain reaction proceed more rapidly in aqueous solution.

Front 159

Catalysts and Reaction Rates

Back 159

Catalysts are substances that increase reaction rate without themselves being consumed.

They do this by lowering activation energy. (e.g. enzymes)

Front 160

What is the expression for the equilibrium constant for the following reaction?

2 H2 + N2 --> 2 NH3

Back 160

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.

Front 161

Reaction Quotient

Back 161

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.

Front 162

Properties of equilibrium constant

Back 162

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.

Front 163

Le Chatelier's Principle

Back 163

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.

Front 164

Le Chatelier's Principle: Change in Concentration

Back 164

Increasing [] of a species will tend to shift the equilibrium away from the species that is added, in order to reestablish its equilibrium [].

Front 165

Le Chatelier's Principle: Changes in Pressure or Volume

Back 165

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).

Front 166

Le Chatelier's Principle: Changes in Temperature

Back 166

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.

Front 167

Types of Systems

Back 167

1. Isolated
2. Closed
3. Open

Front 168

Isolated system

Back 168

When it cannot exchange energy or matter with the surroundings, as with an insulated bomb reactor.

Front 169

Closed system

Back 169

When it CAN exchange ENERGY, but NOT matter with the surroundings, as with a steam radiator

Front 170

Open System

Back 170

When it can exchange BOTH matter and energy with the surroundings, as with a boiling pot of water.

Front 171

Isothermal

Back 171

Temperature of the system reamins constant.

Front 172

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

Back 172

a. 100L

Front 173

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.

Back 173

d. the calcium cation has less protons.

Front 174

Which of the following has the highest second ionization energy?

a. Ca
b. Sr
c. Ba
d. Ra
e. Na

Back 174

e. Na

Front 175

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

Back 175

a. 2/3

Front 176

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

Back 176

e. r = k

Front 177

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

Back 177

a. change in gibbs free energy is zero

Front 178

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

Back 178

e. amount of disorder in a system

Front 179

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

Back 179

c. change in enthalpy is positive

Front 180

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

Back 180

b. the solution's boiling point will increase.

Front 181

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

Back 181

a. water will flow into the capillary

Front 182

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.

Back 182

e. all of the above.

Front 183

Charle's Law

Back 183

(V1/T1) = (V2/T2)

V/T (Volume of gas is directly proportional to its absolute temperature)

Front 184

Boyle's Law

Back 184

PV = k or P1V1 = P2V2

Volume of gas is proportional to its pressure.

Front 185

Ideal Gas Law

Back 185

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.

Front 186

Heat

Back 186

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.

Front 187

Calorimetry

Back 187

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

Front 188

Standard conditions

Back 188

25 degrees Celsius and 1 atm.

Normally used to measure enthalpy, entropy, and free energy of a reaction.

Front 189

Enthalpy

Back 189

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.

Front 190

Delta H

Back 190

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

Front 191

Standard Heat of Formation

Back 191

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.

Front 192

Standard Heat of Reaction

Back 192

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)

Front 193

Hess' Law

Back 193

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.

Front 194

Heat of vaporization

Back 194

Enthalpy change of a liquid to gas, Delta H (vap).

Front 195

Bond Dissociation Energy

Back 195

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

Front 196

Heats of Combustion

Back 196

Delta h(comb).

Reaction is often fast and spontaneous.

Front 197

Entropy

Back 197

(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

Front 198

Delta S (Standard Entropy Change)

Back 198

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.

Front 199

Gibbs Free Energy

Back 199

(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.

Front 200

Characteristics of Gibbs Free Energy

Back 200

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.

Front 201

Standard Free Energy

Back 201

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.

Front 202

Reaction Quotient

Back 202

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)

Front 203

Three Different Forms of Matter

Back 203

Gas, liquid, Solid

Front 204

Four variables that define a gaseous samples

Back 204

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.

Front 205

Standard Temperature and Pressure

Back 205

273.15 K (0 degrees Celsius) and 1 atm.

Front 206

Boyle's Law

Back 206

At constant temperature, the volume is inversely proportional to its pressure:

PV = k or P1V1 = P2V2

k is proportionality constant.

Front 207

Charles Law

Back 207

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

V/T = k or (V1/T1) = (V2/T2)

Front 208

Avogadro's Principle

Back 208

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

Front 209

Ideal Gas Law

Back 209

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)

Front 210

What volume would 12 g of helium occupy at 20 degrees C and a pressure of 380 mm Hg?

Back 210

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

Front 211

Density

Back 211

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

Front 212

Calculating change in Volume

Back 212

V2 = V1(P1/P2)(T2/T1)

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

d = m / V2

Front 213

Molar Mass

Back 213

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.

Front 214

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?

Back 214

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

Front 215

Dalton's Law of Partial Pressures

Back 215

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)

Front 216

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?

Back 216

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

Front 217

Assumptions of the Kinetic Molecular Theory of Gases

Back 217

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.

Front 218

Graham's Law of Diffusion and Effusion

Back 218

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)

Front 219

Condensed Phases

Back 219

Liquid or solid phases. Because of their smaller volume relative to gases.

Front 220

Evaporation

Back 220

Or Vaporization is a cooling process. Liquid to Gas.

Front 221

Condensation

Back 221

Gas to liquid.

Front 222

Vapor Pressure

Back 222

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.

Front 223

Boiling point

Back 223

Temperature at which the vapor pressure of the liquid equals the external pressure.

Front 224

Fusion

Back 224

Or melting, solid to liquid.

Temperature at which this occurs is melting point.

Front 225

Solidification/freezing

Back 225

Or crystallization is liquid to solid.

Temperature at which this occurs is freezing point.

Front 226

Sublimation

Back 226

Solid to gas (e.g dry ice)

Front 227

Triple Point

Back 227

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.

Front 228

Critical Point

Back 228

Tip between liquid and gas phases.

Temperature and pressure above which no distinction between liquid and gas is possible.

Front 229

Colligative Properties

Back 229

Physical properties derived solely from the number of particles present, not the nature of those particles.

Front 230

Freezing-Point Depression

Back 230

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)

Front 231

Boiling-Point Elevation

Back 231

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.

Front 232

Osmotic Pressure

Back 232

Substances tend to diffuse from higher to lower concentrations.

O = MRT

M = molarity
R = ideal gas constant
T = Temperature in Kelvin

Front 233

Raoult's Law

Back 233

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)

Front 234

Dissolution

Back 234

Interaction between solute and solvent molecules.

Front 235

Percent Composition by Mass

Back 235

Is the mass of the solute divided by the mass of the solution, multiplied by 100.

Front 236

Sublimation

Back 236

Solid to gas (e.g dry ice)

Front 237

Triple Point

Back 237

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.

Front 238

Critical Point

Back 238

Tip between liquid and gas phases.

Temperature and pressure above which no distinction between liquid and gas is possible.

Front 239

Colligative Properties

Back 239

Physical properties derived solely from the number of particles present, not the nature of those particles.

Front 240

Freezing-Point Depression

Back 240

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)

Front 241

Boiling-Point Elevation

Back 241

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.

Front 242

Osmotic Pressure

Back 242

Substances tend to diffuse from higher to lower concentrations.

O = MRT

M = molarity
R = ideal gas constant
T = Temperature in Kelvin

Front 243

Raoult's Law

Back 243

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)

Front 244

Dissolution

Back 244

Interaction between solute and solvent molecules.

Front 245

Percent Composition by Mass

Back 245

Is the mass of the solute divided by the mass of the solution, multiplied by 100.

Front 246

Arrhenius Acid

Back 246

Species that produces H+ (a proton) in an aqueous solution.

Front 247

Arrhenius Base

Back 247

Species that produces OH- (a hydroxide ion) in an aqueous solution.

Front 248

Bronsted-Lowry Acid

Back 248

Species that DONATES protons.

Front 249

Bronsted-Lowry Base

Back 249

Species that ACCEPTS protons.

e.g. NH3 and Cl-

Not limited to aqueous solutions.

Front 250

Conjugate acid-base pairs

Back 250

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

Front 251

Lewis Acid

Back 251

Electron-pair ACCEPTOR.

e.g. BCl3 and AlCl3

Front 252

Lewis Base

Back 252

Electron-Pair DONOR.

Front 253

Measurement of pH

Back 253

pH = -log[H+] = log(1/[H+])

Front 254

Measurement of pOH

Back 254

pOH = -log[OH-] = log (1/[OH-])

Front 255

Water dissociation Constant

Back 255

Kw = [H+][OH-] = 10^-14

An equilibrium reaction.

Rewrite:

pH + pOH = 14

Front 256

Fundamental Properties of Logarithms

Back 256

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

Front 257

If [H+] = 0.001, what is the pH?

Back 257

0.001 = 10^-3, then pH = 3

Front 258

If Kb = 1.0 x 10^-7, then what is the pKb?

Back 258

1.0 x 10^-7, then pKb = 7

Front 259

If Ka = 1.8 x 10^-5, then what is pKa?

Back 259

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)

Front 260

Strong acids and bases

Back 260

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.

Front 261

Weak Acids and Bases

Back 261

Those that partially dissociate in aqueous solution.

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

HA + H2O --> H3O + A

Front 262

Acid dissociation Constant

Back 262

Ka = ([H3O+][A-]) / [HA]

The weaker the acid, the smaller the Ka.

Front 263

Base dissociation Constant

Back 263

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]

Front 264

Conjugate acid

Back 264

Defined as the acid formed when a base gains a proton.

Front 265

Conjugate base

Back 265

Formed when an acid loses a proton.

e.g. HCO3-/CO3^2- (conjugate acid/base pair)

Front 266

Ka x Kb = Kw = 1.0 x 10^-14

Back 266

Remember this.

Front 267

Find the concentration of H+ in a 2.0 M aqueous solution of acetic acid, CH3COOH (Ka = 1.8 x 10^-5).

Back 267

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

Front 268

Neutralization Reaction

Back 268

Acids and bases may react with each other, forming a salt and (often, but not always) water, in a neutralization reaction.

Front 269

Hydrolysis

Back 269

Reverse reaction of a neutralization reaction. Salt ions react with water to give back acid or base.

Front 270

Titration

Back 270

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.

Front 271

Equivalence point

Back 271

When number of acid equivalents equals the number of base equivalents added, or vice versa.

Front 272

Buffer Solution

Back 272

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).

Front 273

Henderson-Hasselbach equation

Back 273

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.

Front 274

Calculating pH for weak acid buffer solution

Back 274

pH = pKa + log ([conjugate base][weak acid])

When [conjugate base] = [weak acid], pH = pKa

Front 275

Calculating pOH for weak base buffer solution

Back 275

pOH = pKb + log ([conjugate acid]/[weak base])

When [conjugate acid] = [weak base], pOH = pKb

Front 276

- log Ka = ?

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

Back 276

- log Ka = pKa

- log (5.6 x 10^-11) = 11 - log 5.6

Front 277

Electrochemical Reactions

Back 277

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.

Front 278

Law of Conservation of Charge

Back 278

An electrical charge can be neither created NOR destroyed. Thus, an isolated loss or gain of electrons cannot occur.

Front 279

Oxidation

Back 279

Loss of electrons

Front 280

Reduction

Back 280

Gain of Electrons

Front 281

Redox reaction

Back 281

An electron transfer.

"OIL RIG"

Oxidation is LOST and Reduction is GAIN

Front 282

Oxidizing agent

Back 282

Causes another atom in a redox reaction to undergo oxidation, and is itself reduced.

Front 283

Reducing agent

Back 283

Causes the other atom to be reduced, and is itself oxidized.

Front 284

Assigning Oxidation Numbers

Back 284

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.

Front 285

Oxidation Numbers

Back 285

Are assigned to atoms in order to keep track of the redistribution of electrons during a chemical reaction.

Front 286

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

Back 286

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.

Front 287

Half-reaction method

Back 287

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.

Front 288

Balance this redox reaction:

MnO4^- + I^- --> I2 + Mn^2+

Back 288

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.

Front 289

Electrochemical Cells

Back 289

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"

Front 290

Galvanic Cells

Back 290

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.

Front 291

Electrolytic Cells

Back 291

Nonspontaneous reactions occur.

Has a positive Delta G and is therefore NONSPONTANEOUS.

Front 292

Rules for Constructing a Cell Diagram:

Back 292

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.

Front 293

Anode of Electrolyic Cells

Back 293

Positive

Since attached to the positive pole of the battery and so attracts anion from the solution.

Front 294

Anode of Galvanic Cell

Back 294

NEGATIVE

Because the spontaneous oxidation reaction that takes place at the galvanic cell's anode.

Front 295

Cathode of Electrolytic Cells

Back 295

NEGATIVE

Front 296

Cathode of Galvanic Cell

Back 296

Positive

But in both types oxidation occurs at anode and reduction occurs at cathode.

Front 297

Standard Electromotive Force (EMF or E(cell))

Back 297

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.

Front 298

Relation between EMF and Delta G

Back 298

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.

Front 299

Nernst Equation

Back 299

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)

Front 300

Relation between EMF and the Equilibrium Constant (Keq)

Back 300

Delta G = -RT ln Keq

Combining equations:

Delta G = -nFE = -RT ln Keq

OR

nFE = RT ln Keq

Front 301

Mass defect

Back 301

E = mc^2

Front 302

Nucleons

Back 302

Protons and neutrons

Front 303

Isotope

Back 303

For a nucleus of a given element with a given number of protons, the various nuclei with different numbers of neutrons.

Front 304

Fusion

Back 304

Combine smaller nuclei into larger nucleus.

Front 305

Fission

Back 305

Split larger nucleus into smaller nuclei

Front 306

Alpha Decay

Back 306

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

Front 307

Beta Decay

Back 307

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

Front 308

Gamma Decay

Back 308

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.

Front 309

Calculating Half-life

Back 309

(t 1/2) =

(1/2)^n

n = half-lives.

Front 310

If the half-life of a certain isotope is 4 years, what fraction of a sample of that isotope will remain after 12 years?

Back 310

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)

Front 311

Alpha Decay involves release of/loss of....

Back 311

Loss of a CATIONIC helium

Front 312

A lewis acid, by definition

Back 312

Accepts an electron pair

Front 313

A Bronsted-Lowry base, by definition....

Back 313

Donates H+

Front 314

In beta- decay, the mass number (A)

Back 314

Remains unchanged

Front 315

In beta- decay, the atomic number (Z)...

Back 315

increased by one

Front 316

In Beta+ (positron) decay, the atomic number (Z)...

Back 316

Decreases by ONE

Front 317

In gamma decay, the atomic number and mass number

Back 317

Remain unchanged

Front 318

In electron capture, the atomic number (Z)...

Back 318

Decreases by one

Front 319

Electron Capture

Back 319

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.

Front 320

Exponential Decay

Back 320

Delta n / Delta t = -wavelength(n)

wavelength = decay constant
n = number that reamin
t = time

Front 321

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

Back 321

a. energy constant

Front 322

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.

Back 322

a. Spontaneous processes experience an increase entropy.

Front 323

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

Back 323

C. 96,499 C

Faraday constant = 96,485.34 C/mole

Front 324

Charge of a single electron

Back 324

1.602 x 10^-19 Coulombs

Front 325

Units of Coulomb

Back 325

J/V

(Joules/Volts)

Front 326

Units of Current

Back 326

C/S

(Coulombs/Second)

Front 327

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....

Back 327

Remain unchanged.

Front 328

In a galvanic Cells, the anode is labeled...

Back 328

Negative

Front 329

In an electrolytic cells, the cathode is labeled...

Back 329

Negative

Front 330

Electrolytic cells differ from galvanic cells in that they....

Back 330

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.

Front 331

The formula for the Gibbs Free Energy Change for a Redox reaction is....

Back 331

Delta G = -nFE

Front 332

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

Back 332

d. none of the above

spin is always -1/2 and +1/2, opposite to each other.

Front 333

(n + l) rule

Back 333

Indication of which subshell is lower in energy and will fill first.

e.g. 3d, n = 3 and l = 2, so (n + l) = 5

Front 334

An electron with 26 protons, and is in its ground state, will have how many electrons in its 3d orbital?

Back 334

6 electrons

Front 335

What is the maximum number of electrons that can be present in an atom with a principal quantum number of 5?

Back 335

50

Front 336

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

Back 336

b. increase in T

Front 337

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

Back 337

b. increase the rate.

Front 338

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.

Back 338

a. there are way more reactants than products.

Keq = products/reactants.

Front 339

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

Back 339

d. A and B

Front 340

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

Back 340

e. none of the above.

Front 341

What is the Normality of a 2.0 M solution of phosphoric acid, H3PO4, for an acid-base titration?

Back 341

N = 6

N = molarity x n

n = number of hydrogens exchanged.

N = (2.0moles/L) x (3 equivalents H3SO4/L) = 6 N

Front 342

How many moles of water are produced in the combustrion of 20 g of glucose in excess O2?

Back 342

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

Front 343

A zero order reaction has the form...

Back 343

r = k

Front 344

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

Back 344

c. change in gibbs free energy is NEGATIVE

Front 345

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

Back 345

a. change in gibbs free energy is ZERO.

Front 346

Entropy is best defined as:

Back 346

Amount of disorder in a system

Front 347

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

Back 347

c. change in enthalpy is positive.

Front 348

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.

Back 348

b. the solution's boiling point will increase

Front 349

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

Back 349

b. buffering

Front 350

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

Back 350

e. B & C

H2CO3 --> H2O and CO2

Front 351

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

Back 351

c. is equal to the SUM of the partial pressures of the gases.

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

Front 352

Define specific gravity?

Back 352

The ratio of density of a substance in any units to the density of water in the same units.

Front 353

Define anhydride

Back 353

A molecule formed by the removal of hydrogen and oxygen from Water to form another molecule.

Front 354

Which molecule is soluble?

a. AgCl
b. PbCl2
c. HgBr2
d. HgCl2
e. none of the above

Back 354

e. none of the above

Front 355

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

Back 355

c. all molecules containing ammonium ions.

Front 356

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

Back 356

e. all of the above.

Front 357

Fusion

Back 357

Melting

Solid to liquid

Front 358

Crystallization

Back 358

Freezing

Liquid to Solid

Front 359

Vaporization

Back 359

Evaporation

Liquid to gas

Front 360

Condensation

Back 360

Gas to liquid.

Front 361

Triple point

Back 361

All three phase are in equilibrium at this T and P. And is where all three phases are present at the same time.

Front 362

Deposition

Back 362

Gas to Solid

Front 363

Electron Motive Force (EMF)

Back 363

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

Front 364

If the atomic weight of carbon-12 is exactly 12 amu, what is the mass of a single carbon-12 atom?

Back 364

2.0 x 10^-23

12 g/mol x (1 mol/6.022x10^-23) = 2.0 x 10 ^-23

Front 365

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

Back 365

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%

Front 366

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.

Back 366

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.

Front 367

Ka < 1, then pKa = ?

Back 367

pKa > 1

smaller Ka = weaker the acid
therefore, the bigger the pKa the weaker the acid.

Front 368

Freezing-Point Depression

Back 368

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.

Front 369

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

Back 369

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.

Front 370

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-

Back 370

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.

Front 371

Isoelectric species

Back 371

Atoms and ions that have the same number of electrons.

Front 372

Characteristics of Covalent Bonds

Back 372

1. Electrons are shared between the elements.
2. Elements bonded together are usually two nonmetals.

Front 373

NH3(aq) + HCl(aq) --> ?

Back 373

Acid base reaction.

Product is NH4Cl

Front 374

BaCO3(s) + HCl(aq) --> ?

Back 374

Gas forming reaction.

Product is H2O + CO2 + BaCl2

Sum of coefficents is 6

Balanced net reaction:
BaCO3 + 2 H --> H2O + CO2 + Ba2+

Front 375

NaF(aq) + Ba(NO3)2(aq) --> ?

Back 375

Net Ionic Eqation for reaction:

2F- + Ba^2- --> BaF2

Front 376

Ionic Compounds

Back 376

Made up of cations and anions.

Front 377

-ous

Back 377

represents lesser charge

Front 378

-ic

Back 378

represents greater charge

Front 379

Fe2+

Back 379

Ferrous

Front 380

Fe3+

Back 380

Ferric

Front 381

Cu2+

Back 381

Cuprous

Front 382

Cu3+

Back 382

Cupric

Front 383

-ide

Back 383

Monatomic elements are named by dropping ending of the name of element and adding -ide

Front 384

H-

Back 384

Hydride

Front 385

F-

Back 385

Fluoride

Front 386

O^2-

Back 386

Oxide

Front 387

S^2-

Back 387

Sulfide

Front 388

N^3-

Back 388

Nitride

Front 389

P^3-

Back 389

Phosphide

Front 390

Oxyanions

Back 390

Polyatomic anions that contain OXYGEN.

Front 391

-ite

Back 391

When an element forms two oxyanions, then name of the one with LESS oxygen ends in -ite.

Front 392

-ate

Back 392

When an element forms two oxyanions, then name of the one with MORE oxygen ends in -ate.

Front 393

NO2-

Back 393

Nitrite

Front 394

NO3-

Back 394

NItrate

Front 395

SO3^2-

Back 395

Sulfite

Front 396

SO4^2-

Back 396

Sulfate

Front 397

Hypo-

Back 397

When the series of oxyanions contain FOUR oxyanions, then prefixes used.

LESS Oxygen = Hypo-

Front 398

Per-

Back 398

When the series of oxyanions contain FOUR oxyanions, then prefixes used.

MORE oxygen = Per-

Front 399

ClO-

Back 399

Hypochorite

Front 400

ClO2-

Back 400

Chlorite

Front 401

ClO3--

Back 401

Chlorate

Front 402

ClO4

Back 402

Perchlorate

Front 403

Percent composition by mass

Back 403

[Mass of solute/mass of the solution (solute + solvent)] x 100

Front 404

Hess' Law

Back 404

States that enthalpies of reactions are additive.