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Flashcards in MCAT Chemistry Deck (500):
1

Cathode

A negatively charged electrode which attracts cations. Where reduction takes place.

2

How does a concentration change affect equilibrium?

If one of the substances is added or removed, all the concentrations of substances adjust to a new equilibrium with the same Keq

3

Electronegativity difference

A number found by taking the difference between the electronegativities of two atoms in a bond. Its value determines the type of bond.

4

List some basic facts about solutions

Particle size less than 1 mmicron; Clear (may be colored); Particles don't settle; Can pass through membranes; Particles not visible

5

Endothermic Reaction: ΔH is (positive or negative); Enthalphy is (increased or decreased)

Positive; Increased

6

Atomic Mass Unit

the mass of a proton or neutron is equal to 1 atomic mass unit. Symbol - "amu"; 1 amu=1.66x10^-24

7

Hydrogen bonding

Formed when hydrogen is bonded to oxygen, fluorine, or nitrogen. The hydrogen of one molecule becomes attracted to the electronegative element of the other molecule. These intermolecular attractions cause higher boiling points than predicted

8

Phase equilibrium

For a liquid in a closed container, when the rates of evaporation (liquid to gas) and condensation (gas to liquid) equalize; the concentration of each is stable.

9

Emission Spectrum

A bright line spectrum formed when energy absorbed by an element is emitted at specific wavelengths. Each element has a unique spectrum.

10

Le Chatelier's Principle

If stress is placed on a system at equilibrium, the equilibrium shifts in order the counteract the effects of the stress and regain equilibrium

11

in H2O solutions: 1) How many ˚C is the freezing point depressed for each molal of solute? 2) How many ˚C is the boiling point elevated for each molal of solute?

1) 1.86C for each molal of particles of solute; 2) 0.51C for each molal of particles of solute

12

Electrode Potential

A measure in volts of the tendency of atoms to gain or lose electrons. (Relative to a H2 oxidation reaction which has an assigned value of zero)

13

State the collision theory of reaction rates

There must be collisions between reactants. Reaction rate depends on number of collisions per unit time and the percent which are successful (Have sufficient energy)

14

Use conversion factors to solve: 1) How many kilograms in 2000 g? 2) How many feet in 60"?

1) (2000g)(1kg/1000g) = 2kg; 2) (60")(1'/12") = 5 ft

15

Describe the change in nonmetallic activity across the periodic table

Non-metallic activity increases across the table (L to R) and decreases down the columns

16

Molarity - general equation

The number of moles of solute it a liter of solution; M = n/L

17

Normality

The number of gram-equivalent weights in a liter of solution

18

Anode

A positively charged electrode which attracts anions. Where oxidation takes place.

19

Triple Point

The only temperature and pressure combination at which the 3 phases of a substance (solid, liquid, gas) can co-exist in equilibrium

20

Which sublevels are present in energy levels 1, 2, 3, and 4-7?

1: s; 2: s, p; 3: s, p, d; 4-7: s, p, d, f

21

Alkene

A series of hydrocarbons containing at least one double covalent bond (CnH2n)

22

How does a pressure change affect equilibrium?

Only in reactions where gases are involved. The reaction will shift to oppose pressure change, resulting in fewer moles of gas particles

23

Positive Ion: which elements tend to form them?

Metals tend to form positive ions by losing electrons (Na → Na+ e-)

24

What are the two driving forces that control reactions?

A drive towards increased entropy (disorder). A drive towards decreased enthalpy (lower heat content)

25

Dalton's Law of Partial Pressures

In a gas mixture, the total pressure equals the sum of the partial pressures of each component. Ptotal = P1 + P2 + P3…

26

Oxidation

The loss of electrons (ex. Cumetal →Cu+2 +2e-

27

Bronsted Theory

An acid is a proton donor. A base is a proton acceptor.

28

Metals. List 6 characteristics

Solid (except Hg); malleable & ductile; conduct heat & electricity; shiny reflective & lustrous; lose e- to form cations; good reducing agents

29

STP

Standard temperature (0C or 273 K) and Standard Pressure (1 atm or 760 torr)

30

Scientific notation

A number expressed as Ax10B. "A" is between 1.00 and 9.99 and "B" is an integer.

31

What are the valences of the elements of families IA through VIIIA?

The valences of elements in groups IA - VIIIA are the element's column number. For example, the valence of Na is 1; O is 6

32

Atomic Number: What are the atomic numbers of helium, hydrogen, carbon, oxygen?

The number of protons in the nucleus of an atom of an element. Helium-2; carbon-6; hydrogen-1; oxygen-8

33

Carboxylic Acid

A hydrocarbon containing the carboxyl functional group. (R-C(=O)-O-H)

34

Negative Ion: which elements tend to form them?

Non-metals form negative ions by gaining electrons (Cl +e- →Cl-)

35

List the names of the first four energy levels (or shells)

1 or K; 2 or L; 3 or M; 4 or N

36

Aromatic compounds

Unsaturated ring structures with six carbon atoms. Benzene is the simplest aromatic. (CnH2n-6)

37

Beta particle

An electron ejected from the nucleus when a neutron decays to a proton; Increases atomic number by one; High velocity; Low energy

38

pOH. For what values is a solution acidic, basic, neutral?

pOH = -log[OH-]; The degree of basicity of a solution. 7=acid.

39

Coordinate covalent bond

When both electrons in a covalent bond are supplied by one atom

40

Electrolyte (give 3 examples)

A solute whose aqueous solution contains ions and conducts electricity (acids, bases, salts)

41

What is the order for filling sublevels (aufbau process) from lowest to highest energy

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d

42

For any substance, write a general formula to convert from Number of Particles to Moles

n = P/6.02x1023; n = number of moles; P = number of particles

43

Isotope

Atoms which contain the same numbers of protons but different numbers of neutrons (ex. Hydrogen has 3 isotopes with mass numbers of 1,2,3)

44

List names for alkanes with 1 to 10 carbons in a molecular chain

1 methane; 2 ethane; 3 propane; 4 butane; 5 pentane; 6 hexane; 7 heptane; 8 octane; 9 nonane; 10 decane

45

Endothermic

A chemical reaction which absorbs heat

46

Double replacement

A reaction where two ionic substances "trade" anions; NaCl + AgNO3 ↔ NaNO3 + AgCl

47

Second Law of Thermodynamics

The entropy of the universe increases for any spontaneous process.

48

Atomic Mass

A weighted average mass of the atoms of an element (assuming the carbon-12 isotope is exactly 12) (ex. Atomic mass of C=35.45 is calculated from two isotopes. Cl-35 and Cl-36)

49

Ion (provide examples)

An atom or group of atoms which contains a positive or negative electrical charge (ex. Na+; Cl-; SO4^2-)

50

Explain the rules for dividing the numbers in scientific notation

Divide the first numbers and subtract the exponents. (8x10^6)/(2x10^10) = 4x10^-4

51

Electronegativity

The ability of an atom to attract to itself the e- in a covalent bond. Values range from 0.7 (Cs) to 4.0 (F)

52

Six common characteristics of acids

Form H2O solutions; Conduct electricity; React with active metals; Turn blue litmus red; Neutralize bases; Sour taste

53

Fission

The breakdown of heavy nuclei into lighter nuclei. The source of nuclear power

54

Diatomic; list 7 diatomic elements found in nature

A molecule composed of two atoms. H2, N2, O2, Cl2, F2, Br2, F2, I2

55

Specific Heat. Give value for liquid water in calories and joules.

The amount of heat required to raise one gram of substance 1C. Water: 1cal/g-C or 4.18 J/g-C

56

For any substance, write a general formula to convert from Moles to Number of Particles

(n)(6.02x10^23) = P; n = number of moles; P = number of particles

57

Heat of Fusion (value for water)

The amount of energy required to change a gram of substance from solid to liquid at its melting point (water = 80 cal/g)

58

Non-metals. List 7 characteristics

1/2 are gases; solids are brittle; poor conductors of heat and electricity; dull & non-reflective; gain e- to form anions; good oxidizing agents

59

Transition element

A metal having two incomplete shells of electrons; many have multiple oxidation states; less active than family IA & IIA. Ex., Fe, Ag, Au, Cr, W

60

For any substance, write a general formula to convert from Moles to Grams

g=(n)(MM) where, n = moles MM = molecular mass g = grams

61

Valence electroncs

The electrons found in the outermost energy level of an atom

62

Solubility curves

A curve for a given substance which shows how much dissolves in a given amount of solvent at different temperatures.

63

Synthesis or combination

A reaction where 2 or more elements form a compound; N2 + 3H2 ↔2NH3; 2H2 + O2 ↔ 2H2O

64

Equilibrium

The point in a reversible reaction where the forward and reverse reactions are taking place at the same rate.

65

Oxidizing agent

A substance which causes another substance to be oxidized (oxidizing agent is simultaneously reduced)

66

Faraday

A unit of electric charge which deposits by electrolysis one equivalent weight of an element. Equals 96,500 coulombs

67

Electron configuration

The distribution of electrons into shells and sublevels for an atom of an element. Each element has a unique electron configuration.

68

What is a general rule for solubilities of polar and nonpolar compounds?

Like dissolves like; Ionic and polar solvents dissolve ionic, polar solutes (water dissolves salt). Non polar solvents dissolve nonpolar solutes (acetone dissolves gasoline)

69

Write the solubility product expression for AgCl ↔ Ag+ + Cl-

Ksp = [Ag+][Cl-]

70

Metallic Bonds

A sea of electrons surrounding positive metal ions

71

Physical properties

Properties which can be observed without changing the substance into something different. Color, odor, hardness, density, luster, state, conductivity, solubility, boiling and melting points

72

Compare the ionic radii of 1) Cl, S, P; 2) Na, Mg, Al; What is the reason for their different size?

1) Cl-1 is smallest; P-3 is largest (the P atom gained the most e-); 2) Na+1 is largest Al+3 is smallest (the Al atom lost the most e-)

73

Specific Rate Constant

Symbol is "k" in a rate equation. A constant specific to temperature and reaction which is part of every rate equation

74

Mole

6.02x10^23 items, can be anything. The number of atoms in one mole (atomic mass in grams) of a monoatomic element. The number of formula units in one mole (formula mass in grams) of an ionic compound. The number of molecules in one mole (formula mass in grams) of a molecular substance)

75

Amino Acid

Organic compounds which contain an amine and a carboxyl group. (H2N-CH(-R)-COOH)

76

Conjugate Base - Write conjugate base of HCl

When a Bronsted acid donates a proton, it becomes its conjugate base (conjugate base of HCl is Cl-)

77

Law of Multiple Proportions

The same elements may combine to form more than one compound. The ratios of atoms are in small whole numbers (H2O and H2O2)

78

Ionization energy

The energy change required for the removal of the outermost electron from the gaseous atom to form a +1 ion

79

Molecule

A group of two or more atoms held together by chemical bonds

80

Catalyst

a catalyst is introduced into a reaction to speed it up or slow it down. It is not consumed. An increase or decrease of activation energy results from an alternate reaction path.

81

Equivalent weight

The number of grams of an element which will accept or donate 1 mole of electrons

82

Write the equation to calculate the atomic mass of an element

Atomic mass of X = ((mx1)(%x1)/100%) + ((mx1)(%x1)/100%)) + etc.; mx1, mx2, mxN = atomic masses of each isotope of element; %x1, %x2, %xN = percent composition of each isotope

83

Dynamic equilibrium

In a closed container where opposing changes are taking place at equal rates; the concentration of all components remains constant.

84

Isomer

Compounds with the same molecular formula but different structural formulas (different connectivity)

85

Aldehyde

A hydrocarbon containing the aldehyde functional group (R-C(=O)-H)

86

List five factors that control reaction rate

nature of reactants; exposed surface area; concentrations; temperatures; presence of catalyst

87

Write the nuclear equation: 1) U238 loses an alpha particle; 2) Th234 loses a beta particle.

92U238 →90Th234 +2He4; 90Th234→ 91Pa234 + -1e0

88

Describe the change in electronegativity across the periodic table

Electronegativity tends to increase across the table left to right and decrease down the columns

89

Solvent (provide example

A substance, usually a liquid, into which another substance (solute) is dissolved. Water is the solven in iced tea.

90

List in the order of increasing enthalpy: solid, gas, liquid

Solid (least enthalpy) -> Liquid -> Gas (most enthalpy)

91

List four quantum numbers and their symbols

1) Principal energy level "n"; 2) sublevel "l"; 3) orbital "m1"; 4) spin "ms"

92

Octet Rule

Atoms tend to gain or lose outer shell electrons in order to achieve a noble gas configuration of 8 electrons

93

Single Replacement

A reaction involving the replacement ina compound of an element by another element. Zn + CuCl2 ↔ Cu + ZnCl2

94

Alcohol

hydrocarbons which contain the hydroxyl functional group (OH-) attached to a saturated carbon (R-O-H)

95

Mass number

The sum of protons plus neutrons n the nucleus of an atom

96

Nonelectrolyte (give 2 examples)

A solute whose aqueous solution does not conduct electricity (sugar, benzene, most organic compounds)

97

Alkyne

A series of hydrocarbons containing at least one triple covalent bond (CnH2n-2)

98

Seven common characteristics of bases

Form H2O solutions; Conducts electricity; Turns red litmus blue; Feels slippery; Caustic; Neutralizes acids; Bases + fats form SOAP

99

Vapor Pressure

The pressure the gas phase exerts on its liquid phase in a closed container. This pressure varies with temperature

100

Decomposition

A reaction where a compound breaks down into elements; CO2 ↔ C + O2; 2CaO ↔2Ca + O2

101

Write an equation to calculate energy change when a fixed mass of substance changes temperature

ΔH = mCpΔT, where ΔT = temperature change, Cp = Specific Heat; ΔH = heat absorbed or given off

102

How many grams of NaCl are required to prepare 500 grams of a 5% solution?

%Concentration = (gNaCl/gsolution)(100%); 5% =(x/500g)(100%);x=25 g NaCl

103

Alkane

A series of hydrocarbons with only single covalent bonds (CnH2n+2)

104

Conjugate Acid - Write conjugate acid of I-

When a Bronsted base accepts a proton it becomes its conjugate acid (conjugate acid of I- is HI)

105

Empirical Formula. What are the molecular and empirical formulas of peroxide?

An expression which gives the relative numbers of atoms of the elements in a molecule. Expressed as the lowest possible set of integers (H2O2, HO)

106

Quantum Number

A number used to describe the energy levels available to an electron. Each electron in an atom has a unique set of four.

107

Describe the change in atomic radius across the periodic table

The atomic radius tends to decrease from left to right across the table and increase down the columns

108

List some basic facts about suspensions

No brownian motion; Don't pass through filter paper or a membrane; Cloudy but particles settle on standing; Particles visible with microscope or eye

109

Activation energy

The energy necessary for a reaction to begin. Obtained from the kinetic energy released during collision

110

Gay-Lussac's Law of Combining Gas Volumes

When only gases are involved in a reaction, the volumes of reactants and products are in a small, whole number ratio.

111

Element. List some examples

A substance is composed of identical atoms. Gold, silver, oxygen, hydrogen, lead, chlorine, helium, iron, copper, fluorine, arsenic

112

Density: Write the general equation and three standard units

The mass of a unit volume of a substance. Density = mass/volume = g/mL; g/L; kg/L

113

List properties of ionic substances

Solids at 25C; Non-conducting as solids but conducting as aqueous solutions or liquids. Conducting as aqueous solutions or liquids. High melting & boiling points; Brittle; Low volatilities

114

List prefixes for naming hydrocarbons for 1-10 carbons in a molecule

1 meth; 2 eth; 3 prop; 4 but; 5 pent; 6 hex; 7 hept; 8 oct; 9 non; 10 dec

115

Nonpolar covalent bond

when the EN difference is very small (less than 0.5). Two bonded atoms share the valence e-. The resulting molecule has no electrostatic charge.

116

List three basic postulates of the bohr model for the hydrogen atom

1) e- are present only in specific energy states; 2) a quantum of energy is absorbed or emitted to change energy levels; 3) a quantum is the smallest amount of energy that can be gained or lost

117

How can ΔG be used to predict if a reaction is spontaneous?

When ΔG is positive it is not spontaneous; negative, it is spontaneous; equals 0 it is at equilibrium.

118

Half Reaction

One of the two parts, either the reduction or the oxidation, of an oxidation-reduction reaction

119

Write the equilibrium expression for aA + bB ↔ cC + dD

Keq = [C]^c[D]^d / [A]^a[B]^b, where Keq = Equilibrium constant

120

How can you determine if a redox reaction will take place spontaneously

Add the electrode potentials of the two half reactions. If the result is positive, the reaction is spontaneous; if negative, the reaction is not spontaneous.

121

Alpha particle

a helium nucleus; charge =+2; High energy; Low velocity; Ejection reduces atomic number by 2amu and atomic weight by 4amu

122

Fusion

The joining of lighter nuclei to form heavier nuclei. Source of the sun's energy

123

Describe an ionic bond in terms of electronegativity difference

When the EN values differ by 1.7 or more. The atom with higher EN borrows the electrons from the atom with lower EN. The resulting positive and negative ions attract.

124

Thermochemical equation

A balanced chemical equation which includes the enthalpy change. (H2(g) +1/2O2(g) →H2O(l) ; ΔH = -285kJ

125

Conversion Factor

A fraction which expresses an equality between two units of measurement and can be used to convert from one to the other (ex. 1kg/1000g)

126

Free Energy Change (ΔG). Write the free energy equation.

A property which reflects a system's capacity to do useful work. ΔG = ΔH - TΔS; G=free energy; S=entropy; H=enthalpy; T=kelvin

127

Polar covalent bond

When the EN difference is between 0.5 and 1.7, the bonding electrons stay closer to the more electronegative atom. Electrons are shared unequally.

128

Gram-equivalent weight

The amount of substance which reacts with or displaces 1 mole of H+ ions.

129

Compare potential energy and kinetic energy

Potential energy is energy due to position. KE is energy of motion.

130

Which three factors affect the rate of solubility?

Pulverizing; stirring; heating

131

Significant digits

Digits which are measured. All non-zero digits are significant. Zeros are significant unless they are placeholders

132

Resonance structures

Where there is more than one possible bonding structure in a molecule

133

Write conversion forumlas: 1) Celsius to Fahrenheit; 2) Celsius to Kelvin

1) ˚f = (1.8 x ˚celsius) + 32; 2) Kelvin = ˚celsius +273

134

Calculate the volume of 10M NaOH needed to titrate 5L of 2M HCl

MA x VA = MB x VB; 2M x 5L = 10M x VB; 1L = volume of base

135

What is the volume of 1 mole of any gas at STP?

22.4 L

136

Explain how to convert a number less than 1 to scientific notation

Move decimal point to right until only 1 digit remains to left. Indicate number of moves as a negative exponent of 10. (0.00356 = 3.56 x 10-3

137

Write the electron configuration for 1) Lithium; 2) Iron

1) 1s2, 2s1; 2) 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d6

138

10 prefixes used to name covalent compounds

Mono=1; Di=2; Tri=3; Tetra=4; Penta=5; Hexa=6; Hepta=7; Octa=8; Nona=9; Deca=10

139

Addition reaction

In an unsaturated hydrocarbon, two atoms may be added to the structure across a double or triple bond (C2H2 + Br2 → Ch2Br2)

140

Substitution reaction. Which elements commonly substitute in alkanes?

A reaction where one of the hydrogen atoms in a hydrocarbon is replaced by another. Usually a halogen. (CH4 + Br2→Ch3Br + HBr

141

Saturated and Unsaturated. Which are sturated? Alkanes, alkenes, alkynes.

A compound only containing single covalent bonds is saturated. Alkanes are saturated. A compound containing double or triple bonds is unsaturated. Alkenes and alkynes are unsaturated.

142

Charles Law

V1/T1 = V2/T2 (pressure and amount of gas are constant; V=volume; T=Kelvin)

143

Hydrogenation

The process of adding hydrogen to an unsaturated hydrocarbon

144

How do temperature and pressure affect the solubility of a gas?

Solubility usually decreases with increasing temperature. Solubility increases in direct proportion to an increase in pressure.

145

Reactants; Products

1) The starting materials in a chemical reaction; 2) The substances formed in a chemical reaction

146

Column or Family

A verticle group of elements on the periodic table in the same column. They have similar properties and the same number of valence electrons.

147

Cycloalkane

An alkane which has a ring structure instead of a chain

148

Matter

A substance that occupies space has mass.

149

Hybrid orbitals (list three types)

Where 2 or more pure atomic orbitals are mixed to form identical hybrid orbitals (ex. Sp, sp2, sp3)

150

How are the concentrations of [H+] and [OH-] related in a solution?

The product of the concentrations = 1x10-14; [H+][OH-] = 1x10-14

151

The Law of Conservation of Mass

During a chemical reaction, matter is neither created nor destroyed

152

Avogadro's Number

The number of atoms or formula units in "x" grams of an element or molecule where "x" is the atomic or molecular mass. (Always equal to 6.02 x 10^23)

153

Quantum

A packet of energy associated with a specific wavelength of electromagnetic radiation

154

Cation and Anion (provide examples)

cation - positively charged ion (Na+; Fe+2; NH4+; Ag+); anion - negatively charged ion (Cl-; SO4-2; OH-; P-3)

155

Ground State

Electron is at its lowest energy level as close to the nucleus as possible

156

Heat of Vaporization (value for water?)

The amount of energy required per gram to change a liquid to a gas at its boiling point (water = 540 cal/g)

157

List the names of the four sublevels and their electron capacities

sublevel s holds two electrons; p holds 6; d holds 10; f holds 14

158

Explain the rules for multiplying the numbers in scientific notation

Multiply the first numbers and add the exponents. (3x10^5)(2x10^3)= 6x10^8

159

pH. For what values is a solution acidic, basic and neutral?

pH = -log[H+] ; The degree of acidity of a solution. 7 = basic.

160

Energy

The ability to do work; it is released or absorbed during chemical reactions in the form of heat, light, electricity. (calorie, Joule: 1cal = 4.18J)

161

Solute (provide example)

The substance dissolved in another (solvent). Salt is the solute in salt water.

162

Write the equation for Percent Composition. What is the percent composition of Ca in CA(OH)2?

%=(me/FW)(100%); %= percent composition of the element; me=mass of element in one formula unit; FW = formula weight; (40/74)(100%) = 54% Ca in Ca(OH)2

163

For any substance, write a general formula to convert from Grams to Moles

n = g/MM; n = moles; MM = molecular mass; g = grams

164

Ether

A hydrocarbon containing an ether functional group. (R-O-R')

165

Reducing agent

A substance which causes another substance to be reduced (reducing agent is simultaneously oxidized)

166

Solubility Product Constant (Ksp)

An equilibrium exists in a saturated solution between dissolved and undissolved solute. Ksp is the equilibrium constant for this reaction.

167

Exothermic reaction: ΔH is (positive or negative); Enthalphy is (increased or decreased)

Negative; Decreased

168

List some basic facts about carbon bonding

Carbon forms more compounds than any other element; Each atom requires 4 covalent bonds; Carbon can form long chains and rings; Bonds commonly to O, H, N, S, P, and halogens

169

How is the reaction rate related to concentration?

Reaction rate is directly proportional to the concentrations of reactants

170

Dehydrogenation

The process of removing hydrogen from a hydrocarbon

171

Calculate the molality of 10 moles of H2SO4 dissolved in a 4 kg of water

Molality = Moles Solute / kg solvent = 10/4 = 2/5 Molal

172

Chemical change

Where new substances are formed with new chemical and physical properties. Oxygen & hydrogen form water; sodium & chlorine form sale (sodium chloride)

173

Compound. List some examples

substance composed of two or more elements chemically combined. Water - H2O; Salt - NaCl; Sugar - C6H12O6; Ammonia - NH3

174

Forumlas and names of acids formed from: 1) F; 2) Cl; 3) Br; 4) I

1) HF - hydrofluoric acid; 2) HCl - hydrochloric acid; 3) HBr - hydrobromic acid; 4) HI - hydriodic acid

175

Enthalpy

The heat content of a system

176

Arhennius Theory

An acid yields protons in solution (H+ ions); A base yields hydroxide ions in solution (OH- ions)

177

Enthalpy change (state the equation)

The difference in heat content between the products and the reactants (ΔH = ΣHproducts - Σhreactants)

178

Write the general equation to calculate ΔH for a chemical reaction

ΔH = Σ(ΔHF-products) - Σ(ΔHF-reactants); where ΔHf = Heat of formation of reactants or products

179

Transmutation

The conversion of an element to a new element due to a change in number of protons. Ex. Alpha or beta decay

180

Compare primary and secondary alcohols

Primary: The OH- group is attaced to the end carbon of the chain. Secondary: The carbon bearing theOH- goup is directly attached to two other carbons.

181

Polymerization

The combination of two or more unsaturated molecules to form a larger chain molecule. This is how plastics are made.

182

Describe the change in metallic activity across the periodic table

Metallic activity decreases across the table (left to right) and increases down the columns

183

Physical change

Substance changes form or state only. Boiling, melting, freezing, dissolving, grinding, cutting

184

Exothermic

A chemical reaction which evolves heat

185

Are concentrations of product and reactant equal at equilibrium?

No. The are constant but not equal. Their relative concentrations are determined by the value of the equilibrium constant at that temperature.

186

Ideal Gas Law

PV = nRT, (P=pressure in atm; V = volume in L; n = # moles; R = 0.0820 Latm/Mol-K; T = Temp in K)

187

Amine

A hydrocarbon containing an amine functional group. (R-NH2)

188

Oxidation number

A number (positive or negative) representing the charge on an ion or atom involved in a chemical bond

189

Molar Heat of Formation

The molar heat of formation of a compound is equal to ΔH when 1 mold of compound is formed from its elements at 1 atm and 25C

190

Describe the shapes of the orbitals: 1) s; 2) p; 3) d; 4) f

1) s- sphere; 2) p - dumbbell shape with 2 lobes; 3) d - double dumbbell; most have 4 lobes; 4) f - most have 8 lobes

191

Molecular Mass

Found by adding all the atomic masses of an element. Ex. H2O. Molecular mass = 18; H2 = 2(1) O = 16

192

Compare Weight vs. Mass

Mass is the amount of matter. Weight measures gravitational force. Mass never varies. Weight can vary.

193

Three postulates of Dalton's Atomic Theory

1) An element is composed of identical atoms; 2) Atoms of different elements have different properties; 3) Compounds are atoms of 2 or more elements chemically combined

194

Gamma radiation

Usually emitted with beta radiation; Has neither charge nor mass; High energy; Travels at the speed of light

195

In a thermochemical equation, what happens to ΔH when the moles of reactants double?

Enthalpy is directly proportional to mass. Therefore when the moles double, so does ΔH.

196

Molecular Formula. What is the molecular formula for ammonia?

An expression stating the number and kind of each atom present in a molecule of a substance (NH3 has 1 nitrogen atome and 3 hydrogen atoms in each molecule)

197

Homogeneous and Heterogeneous mixtures. Provide examples.

Homogeneous mixtures are uniform in composition (air, metal alloy, salt water). Heterogeneous mixtures are not uniform in composition (dirt, spaghetti sauce)

198

Write oxidation numbers for ions of: Group IA & IIA; Group VIA & VIIA

IA→+1; IIA→+2; VIA→-2; VIIA→-1

199

Ionization constant. Write the expression for the ionization of acid "HA"

For substances in solution that partially ionize. An equilibrium expression may be written with Ki; Ki = [H+][A-] / [HA]

200

How does a change in temperature affect equilibrium

The reaction shifts to a new equilibrium point with a new Keq. If the temperature is raised, the equilibrium is shifted to reaction which absorbs heat.

201

Chemical properties

A chemical property is observed when a substance changes into a new substance. Iron forms rust in air & water; gasoline burns in oxygen

202

How are ΔH for a forward and ΔH for a reverse reaction related?

forward is equal in magnitude but opposite in sign to ΔH reverse.

203

Law of Mass Action

The rate of a chemical reaction is proportional to the product of the concentrations of the reactants

204

Mixture. List some examples.

A combination of substances held together by physical means (dirt, milk, soup, saltwater, granite)

205

Combined Gas Law

P1V1/T1 = P2V2/T2 (amount of gas is contant. P=pressure, V=vol; T=Kelvin)

206

Neutralization Reaction - Write equation for hydrochloric acid and sodium hydroxide

acid + base → salt + water; HCl + NaOH →NaCl +H2O

207

How are reactant & product concentrations related to the magnitude fo Keq?

Keq is large: [reactant] is small and [product] is large; when Keq is small: [reactant] is large and [product] is small

208

Hess' Law

for a reaction is the same regardless of the path travelled from reactants to products.

209

Valence

The number of electrons in the atom's highest numbered shell.

210

How do temperatures and pressure affect the solubility of a solid?

Solubility usually increases with increasing temperature. Pressure has little effect.

211

Double and Triple covalent bond

In a double bond, wo pairs of electrons are shared. In a triple bond, three pairs of electrons are shared.

212

Boyles Law

P1V1 = P2V2 (Temperature and amount of gas are constant, P=pressure, V=vol)

213

Reversible reaction

A system where the following opposite reactions are taking place: reactant becoming product; product becoming reactant

214

Solution (provide example)

A liquid, gas or solid phase containing 2 or more components uniformly dispersed (air, coffee, saltwater)

215

Names of acids: 1)H2SO4; 2) HNO2; 3) H3PO4; 4)HClO

1) sulfuric acid; 2) nitrous acid; 3) phosphoric acid; 4) hypochlorous acid

216

Period

A horizontal row of elements in the periodic table. All have the same number of shells of e-. Across the period, the elements' properties change.

217

Absorption Spectrum

A dark-line spectrum formed when white light is passed through a vaporized element and a few specific wavelengths are absorbed.

218

The Law of Constant Composition

A compound always contains the same elements combined in the same proportions by mass (H2O) is 88% oxygen no matter where it is found)

219

Reduction

The gain of electrons (Zn+2 +2e- → Znmetal)

220

Half-life

The time it takes for half of a radioactive sample to decay. It can range from a fraction of a second to many years.

221

Three general rules for determining oxidation numbers

1) Atoms of uncombined elements equal 0; 2) Hydrogen = +1 (in metallic hydrides =-1); 3) Oxygen = -2 (in peroxides =-1); (bonded with fluorine =+2)

222

Ester

A hydrocarbon containing an ester functional group. (R-O-C(=O)-R')

223

Excited State

An electron absorbs energy and moves to a higher energy level above the ground state

224

Names of the ionic compounds: 1) FeCL3; 2) FeO; 3) Cu(OH); 3) Cu3PO4

1) iron (III) chloride; 2) iron (II) oxide; 3) copper (II) hydroxide; 4) copper (I) phosphate

225

Balance Equations: H2 + N2 ↔ NH3; NaCL + Br2 ↔ NaBr + Cl2

1) 3H2 + N2 ↔2NH3; 2) 2NaCl + Br2 ↔2NaBr + Cl2

226

List some basic facts about colloids

Particles measure 1-100 mmicrons; Particles don't pass through a membrane; Show brownian motion and the Tyndall effect; Particles don't settle; Clear and pass through filter paper

227

Lewis Theory

An acid is an electron pair acceptor; a base is an electron pair donor.

228

Explain how to convert a number greater than 1 to scientific notation

Move decimal point to left until only 1 digit remains to left. Indicate number of moves as a positive exponent of 10. 3301 = 3.301x10^3

229

List properties of molecular substances

Non-conducting as liquids and solids; Volatile liquids & solids; Many are gases at 25C; Low melting and boiling points; Soft and waxy solids

230

Metalloid. List 5 examples.

An element with the properties of both metals and non-metals. Ex. Si, As, Ge, Sb, Te

231

Normal Boiling Point

The temperature at which a liquid phase becomes a gas phase at a pressure of 1 atm.

232

Describe hybrid bonding in water, ammonia, methane

sp3 bonding results in a tetrahedron shape with bond angles of 109.5˚ in methane and slightly less in water and ammonia.

233

Describe the change in Ionization Energy across the periodic table

Ionization energy tends to increase from left to right across the table and decrease down the columns

234

Ketone

A hydrocarbon containing a ketone functional group (R-C(=O)-R')

235

How are pH and pOH of a solution related?

The sum of the two values = 14; (pH +pOH = 14)

236

Sigma bonds & Pi bonds

A sigma bond is present between any 2 orbitals except when 2 p orbitals share electrons; then this is a pi bond

237

Ionization Constant

Kw = 1 x 10-14 at 25C

238

Metric Abbreviations: Mega, Kilo, Hecto, Deka, Deci, Centi, Milli, Nano

Mega=M; Kilo=k; Hecto=h; Deka =da; Deci =d; Centi =c; Milli=m; Nano=n

239

Avogadro's Law re. gases

Under the same conditions of temperature and pressure; equal volumes of gases contain equal numbers of moles

240

Write formula for compounds of: 1) Sodium & Sulfate; 2) Magnesium & Nitrite; 3) Aluminum & Phosphate

1) Na2SO4; 2) Mg(NO2)2; 3) AlPO4

241

in titrationyou react a ______ volume of a solution of unknown concentrat with a _______ volume of a sol of known concentrat

known volknown vol o

242

ph=

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

243

Blood pH is maintained in a relatively small range (slightlyabove 7) by a

bicarbonate buffersystem. This homeostasis can be upset, leading to a condition known as acidosis.

244

a b equiv is equal to?

one mole of OH- ions

245

NO2 - is the conjugate b of ?

HNO2

246

what combinations of a and b are possible?

1. str a and str b2. str a and weak b3. weak a and str b4. weak a and weak b

247

t or fNH3 and CL- are bronst lowry bases but not arrhenius bases

t

248

to calculate the concentration of H+ in a 2.0 M aq sol of acetic acid CH3COOH (Ka=1.8 X 10-5) what is the concentr of acetic acid at equil equal to?

its initial concentr 2.0 M- X, the amount dissociated

249

NO2-

nitrite

250

give 2 examples of buffers?

examples of buffers are: a solution of acetic acid (CH3COOH) and its salt,sodium acetate (CH3COO-Na+); and a solution of ammonia (NH3) and itssalt, ammonium chloride (NH4+Cl-).

251

m-lognsince n is a number betw 1 and 10 its log will be a fractoin betw ?

0 and 1

252

bronst lowry a and ba isb is?

acid donate h+bases accepts H+

253

name some strong bases

NaOH sodium hydroxideKOH potassium hydroxideother soluble hydroxides of gr IA and IIA metals

254

to calculate the concentration of H+ in a 2.0 M aq sol of acetic acid CH3COOH (Ka=1.8 X 10-5) what can you approximate?

that 2.0-x is approx 2.0 bec acetic acid is a weak acid

255

t or fwater is amphoteric

t

256

weak a with a str b the resulting solut is?

basic.

257

t or fspecies that can act as either oxidiz or reducing agents are considered to be?

amphoteric

258

anion ends in ite it will be __________ acid

ous acid

259

str acid with a weak base the resulting solut will be?

acidic

260

the hydroxides of which metals are usually amphoteric?

Al, Zn, Pb, and Cr

261

Kw= [H+][OH-]=

41561

262

Henderson- Hasselbach equation is?

pH = PKa + log [conjugate base]/[weak acid]

263

the hydroxides of Al, Zn, Pb, and Cr are?

amphoteric

264

H3O+ is the conjugate acid of the base?

H2O

265

This the the titration ofNa2CO3 with HCI in which the polyprotic acid is the ultimate prod-uct.

H2C03

266

a neutrialization react is when?

an a and a base reat w each other, forming a salt and usually water

267

this is titration of a __________ base with an acid

polyprotic base

268

how do you calcul equival weight?

divide the gram molecular weight by how many moles of H+ it liberates

269

litmus paper turns blue in?

bases

270

H2SO4 is a. monoproticb. diproticc. triprotic

b. diprotic

271

you estimate the equival point in titrat by what 2 methods?

1. graphical method- use ph meter2. watching for a color ch

272

Region Ill contains the equivalence point, at which all of the ?

C032 istitrated to HC03.

273

if Ka is large then Kb?

will be small

274

weak acids and bases only partly?

dissociate in aq sol

275

henderson-hasselbachpH = PKa + log [conjugate base]/[weak acid]when will the concentr of conj base equals the conc of the weak acid ?

in a titration, half-way to the equivalent point

276

amphoteric means?

can act as an acid or a base

277

nitrate

NO3-

278

name 3 polyvalent acids?

H2SO4H3PO4H2CO3

279

name of an acid is related to its?

parent anion

280

arrhenius def of acid and baseacid is?base is?

acid prod H+base prod OH-

281

m-lognsince n will be a number betw 0 and 1.m-logn will be between m-1 and ?

m-0

282

why dont indicators change the equival point

bec they are in low concetr

283

ClO3-

chlorate

284

bromide becomes ________________ acid

hydrobromic acid

285

in a body you have the bicarbonate buffersystem. This homeostasis can be upset, leading to a condition known as ?.

acidosis

286

t or fa bronsted lowry defin of an a or a b is limited to aq solutions

f

287

what is the shortcoming of arrhenius def?

only desc a and b in aq media, not in non aqueous

288

str a w a weak bthe cation of the salt formed will do what?

will react w the water solvent, reforming the weak b

289

litmus paper turns red in ?

acid

290

to calculate the concentration of H+ in a 2.0 M aq sol of acetic acid CH3COOH (Ka=1.8 X 10-5) what do you do after writing the equil react?

write the express for the acid dissoc const

291

the henderson-Hasselbalch equation is used for?

The Henderson-Hasselbaich equation is used to estimate the pH ota solution in the buffer region where the concentrations of the speciesand its conjugate are present in approximately equal concentrations.

292

What volume of a 3 M solution of NaOH is required to titrate 0.05 Lof a 4 M solution of HCl to the equivalence point?what equat should you use?

NaVa=NbVb

293

why are buffers useful?

Buffer solutions have the usefulproperty of resisting changes in pH when small amounts of acid or baseare added.

294

To find the Ka of the conjugate acid HC03, the reaction with __________must be considered.

water

295

name some strong acids

HClO4- perchloric acidHNO3 nitric acidH2SO4 sulfuric acidHCl hydrochloric acid

296

CLO-

hypochlorite

297

bronst lowry 2 members of a conj pair are related by a transf of?

a proton

298

ph of a weak a and a weak b depends on?

relative strenght of the reactants

299

HNO2

nitrous acid

300

oxid or reduc agents act amphoteric when they ?

accept or donate electr pairs, acting like Lewis a or bases

301

the _______ of the indicator and the __________ of the a and b should be close.

end pointequivalence point

302

weak bases are usually not?

hydroxides

303

if the conjugate a is strong then the conjugate base?

will be weak

304

CH3COOH (Ka=1.8 X 10-5)write the equilibrium react

CH3COOH (aq)---> H+ (aq) + CH3COO- (aq)

305

to calculate the concentration of H+ in a 2.0 M aq sol of acetic acid CH3COOH (Ka=1.8 X 10-5) [H+] is=?

#ERROR!:parse

306

Ka and _______ are inversely related

Kb

307

indicators are?

weak organic a or bases that have differ colors in their undissoc or dissoc states

308

henderson-hasselbachpH = PKa + log [conjugate base]/[weak acid]when the concentr of conj base equals the conc of the weak acid then?

ph=pKa, bec log1=0

309

lewis def of a and bacidbase?

acid is electr pair acceptbase is electr pair donator.

310

HNO3

nitric acid

311

a conjugate base is formed when?

an a loses a proton

312

HClO2

chorous acid

313

Ka measures?

the degree to which an a dissoc

314

in the early part of the curve, the acidic spec dominates so small amounts of base will do what to the ph?

it will only change the ph a little

315

what is equival weight of H2SO4?

98 g/mol is molec weightsince each mole liberates 2 acid equiv, the gram equiv of H2SO4 is 98/2 or 49 g.

316

Ka= [X][X]/[2.0]=1.8 X 10-5what did you approximate?

that 2.0-x is approx = to 2.0

317

in any aq solut the H2O does what?

dissociates slightly

318

t or fthe equival point is always at ph of 7

fonly for str a with a str base

319

bronst lowry acid and b occur in pairs called?

conjugate a b pairs

320

the partially dissoc conjug base of a polyprotic acid is?

usually amphoteric

321

flouride bec acid?

hydrofluoric acid

322

what are the react for HCL reacting with NH3?

HCl (aq) + NH3 (aq) --> NH4+ (aq) + Cl-(aq) Reaction INH4+ (aq) + H20 (aq) --> NH3 (aq) + H3O+(aq) Reaction II

323

Ka= [X][X]/[2.0-X]=1.8 X 10-5if when you solve for X it is close the the original concentr of acetic acid (2.0 M) what do you need to do?

you cant approx that 2.0-x is =2.0instead you have to use the quadratic equation

324

this is a titration of a ____________ with a strong base

weak acid

325

what is the equation for the dissoc of H2SO4?

H2S04(aq) --> H+(aq) + HS04(aq)HS04(aq) --> H+(aq) + S042-(aq)

326

m-lognthe larger n is the larger ?

the fraction logn will be the answer will be closer to m-1

327

bases turn litmus paper?

blue

328

t or fneutralization reactions usually go to completion

t

329

the addit of base most ch the concentr of H+ and OH- where on the graph?

near the equivalence point. this is where the ph the most

330

HCO3-/CO3- which is the conjugate a and b?

conjugate a HCO3-conjugate b CO3-

331

how can you estimate a logn x 10-m =

log(n x 10-m)= -m + logn the negative log ism-logn

332

Ka=

#ERROR!:parse

333

HA + BOH ---> BA + H2Othis is the equatrion for?

neutralization

334

the weaker the base the ____________ the Kb

smaller

335

ph + pOH=?

14

336

1 x 10-8 M HCL solution.Kw=(x+ 1 x lO-8)(x)=1.O x 1O-14 molecules).when do use this kind of equation. ?

when you have to calcul the H+ concentr with a weak acid. you could also use the equation for a weak base

337

AlCl3 is a _______________ acid but not a ________________ acid

lewis but not a bronst lowry

338

strong acids and bases are those that?

completely dissociate into their component ions in aq solution

339

If Ka = 1 .8 x 10-5, then pKa = estimate please

5 - log 1.8. Since 1 .8 is small,its log will be small, and the answer will be closer to 5 thanto 4. (The actual answer is 4.74.)

340

when NaOH dissoc in water, why can you normally ignore the contribution of OH- from H2O?

as long as the amount of OH- or H+ is greater than 10-7 then you can ignore contribution of water.

341

NO3-

nitrate

342

acids from oxyanions are called?

oxyacids

343

nitric acid

HNO3

344

H2S04(aq) --> H+(aq) + HS04(aq)HS04(aq) --> H+(aq) + S042-(aq)what is this reaction showing?

that H2SO4 is diprotic

345

pOH is=

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

346

hypochlorite

CLO-

347

HCl is?

hydrochloric acid

348

hydrolysis (in terms of a and b) is?

reverse react of neutralization, when salt ions react w water to give back the a or b

349

Ka and Kb are _____________ related

inversely

350

anion ends in ate, it will be ____________ acid

ic acid

351

weaker the acid, the smaller the?

Ka

352

when cant you ignore the contrib of water to acid and base calculations?

when it is a very weak acid or base whose H+ or OH- concentr is close to 10-7

353

In region I, little acid has been added and the predominant species is

C032-

354

MnO4- is what anion?

permanganate, even though there are no manganate or manganite ions

355

t or fspec that can act as red or oxidiz agents are consid to be amphoteric

t

356

AlCl3 can accept a?

electron pair

357

to calculate the concentration of H+ in a 2.0 M aq sol of acetic acid CH3COOH (Ka=1.8 X 10-5) what should you do first?

write the equilibrium reaction

358

Kw=

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

359

HC03- (aq) + H20 (I) ---> H3O+ (aq) + C03 2- (aq)this reaction is used to calculate the ?

Ka of HCO3-

360

pH above 7 indicates an excess of?

OH- ions

361

the ______________ defin of an a or a b is limited to aq solut

arhhenius

362

pure water _____________ is equal to ________________

H+ is equal to OH-

363

to calculate the concentration of H+ in a 2.0 M aq sol of acetic acid CH3COOH (Ka=1.8 X 10-5) write the acid dissoc const?

Ka=[H+][CH3COO-] ------------ = 1.8 X 10-5 [CH3COOH]

364

HClO3

chloric acid

365

2M H3PO4 would have be _________ N

6

366

At equilibrium, a certain acid, HA, in solution yields 0.94 M [HA] and0.060 M [A-].Calculate Ka.

HA ---> H+ + A- The molar ratio of A- to H+ is 1:1, so [H+] must also be 0.060 M at equilibrium. It follows, then, that: Ka [A-][H+]/[HA] = (0.060)(0.060)/(0.94) = 3.8 X 10-3

367

what can we calculate for an acid and its conjugate base?

Ka X Kb=Kw=l x10-14

368

HF is?

hydrofluoric acid

369

you can use the henderson-hasselbach equation to make a buffer at any?

pH. by carefully choosing a weak acid and its salt

370

13. Which of the following combinations would produce a buffer solutionof pH = 4? (Ka HN02=4.5 X 10-4)A. 0.30 M HNO2, 0.22 M NaNO28. 0.22 M HNO2, 0.30 M NaNO2C. 0.11 M HNO2, 0.50 M NaNO2D. 0.50 M HNO2, 0.11 M NaN

pH = PKa + log [A-/[HA]4 = 3.35 + log [A-]/[HA]0.65 = log [A-]/[HA][A-]/[HA] = 4.5 Only Choice C fulfills this criterion as 0.50/0.11 = 4.5.

371

an a equiv is equal to ?

one mole of H+

372

weak a with a str b why is the result solut basic?

bec of hydrolysis of the salt to reform the aci and format of OH- from hydrolyzed water molecules.

373

1 x 10-8 M HCL solutionKw=(x+ 1 xlO-8)(x)=1.O x 1O-14 molecules).what is X?

where x=[H+]=[OH-] (bothfromthe dissociation of water

374

t or fthe titration of any monoprotic a with a str b will give a similar curve to this

t

375

a substance that can act as an a or a b is?

amphoteric

376

In region V. theequivalence point for the entire titration is reached, as all of the HCO3-,is converted to?.

H2C03

377

Region Ill contains the _______________ , at which all of the C032 istitrated to HC03. As the curve illustrates, a rapid change in pH occurs atthe ____________

equivalence point

378

a buffer consists of a weak acid or a weak base and?

its salt

379

Consider a buffer solution of acetic acid and sodium acetate:CH3COOH H+ + CH3COO-what happens when NaOH is added to the buffer?

When a small amount of NaOH is added to the buffer, the OH- ionsfrom the NaOH react with the H+ ions present in the solution; subse-quently, more acetic acid dissociates (equilibrium shifts to the right),restoring the [H+]. Thus, an increase in [OH-] does not appreciablychange pH.

380

Consider a buffer solution of acetic acid and sodium acetate:what happens when a small amount of HCl is added?

H+ ions from the HCl react with the acetate ions to form acetic acid. Thus[H+] is kept relatively constant and the pH of the solution is relativelyunchanged.

381

HNO3 is?

nitric acid

382

this is an ex of?HClO + NaOH ---> NaClO + H2O

weak acid w a str base

383

ClO4-

perchlorate

384

Kw is the?

water dissociation constant

385

HBr is ?

hydrobromic acid

386

Ka X Kb=Kw=l x10-14this is for an acid and its?

conjugate base

387

HClO

hypochlorous acid

388

the weaker the acid, the _____________ the Ka

smaller

389

titration is used to ?

determ the molarit of an a or a b

390

perchloric acid

HClO4

391

in the early part of the curve, the __________ species

acidic species

392

chlorous acid

HClO2

393

t or fKa=[H3O][A-]/[HA]

t

394

math reminderlog(xy)=?

logx + logy

395

equival point is when the?

number of acid equival equals the number of base equival

396

bronsted lowryan amphoteric species can?

gain or lose a proton

397

In region Il, more acid has been added and the predominantspecies are ______________, in relatively equal concentrations. The flatpart of the curve is the ______________, c

C032- and HC03,first buffer region

398

sulfuric acid is?

H2SO4

399

accord to this graph, has the best indicator been chosen?

no, bec the end point has a pretty different ph than the equival point of the titration

400

salt ions reacting with water to give back the acid or base is ?

hydrolysis

401

In region IV, the acid has neutralized approximately half of the HC03,and now _____________are in roughly equal concentrations.

H2C03 and HC03

402

prod of a react betw equal concentr of a str acid and str base are?

salt and water, a and b neutr each oth so ph is 7

403

Kb is the

base dissociation constant

404

ClO2-

chlorite

405

pH below 7 is

acidic

406

equation for neutralization

HA + BOH ---> BA + H2O

407

the acid HClO has a Ka = 3.2 X 10-8,and the base NH3 has a Kb = 1.8 x 1O-5. what will the resulting solut be? acidic, basic, or neutral?

an aqueous solution of HCIOand NH3 is basic since Ka for HCIO is less than Kb for NH3

408

what formula do yo use to calculate the volume added to reach the endpoint?

NaVa=NbVb

409

the equival point is in the _______ range

basic

410

to calculate the concentration of H+ in a 2.0 M aq sol of acetic acid CH3COOH (Ka=1.8 X 10-5) How can you rewrite the Ka expression?

Ka= [X][X]/[2.0-X]=1.8 X 10-5

411

prod of a react betw str a and weak b are?

SALT, but usually no water is formed bec weak b are usually not hydroxides

412

chlorate

CLO3-

413

one mole of H2SO4 can prod ______ acid equiv

2

414

when titrating polyprotic a or b there are several?

equivalence points

415

In region V. the___________________for the entire titration is reached, as all of the HCO3-,is converted to H2C03. Again, a rapid change in pH is observed near theequ

equivalence point

416

nitrite

NO2-

417

In region IV the flatregion is the?

second buffer region of the titration curve.

418

the degree to which an an a dissoc is measur by?

the acid dissoc const Ka

419

hydrofluoric acid is?

HF

420

perchlorate

ClO4-

421

anion ending in -ide will become ______________ acid

hydro ic

422

point at which the indicat ch color is called the?

end point

423

Note that as a weak acid orbase, the effect on pH willalways be _______________ of astrong acid or base of the sameconcentration.

less than that

424

Azimuthal Quantum Number

Second quantum number, designated by the letter "l." This "angular momentum" refers to the subshells within each principle quantum energy level. It can take on the value of an integer in the 0 to (n-1) range.

425

Atomic Emission Spectra

Discontinuous line spectra of light produced when excited atoms return to their ground state and emit photons of a certain frequency.

426

Electron Configuration

Patterned order by which electrons fill subshells and energy levels in an atom. First number designates principal quantum number (n); s, p, d, f, or g specify the subshell (l); and the superscript indicates the number of electrons in that subshell.

427

Atomic Number

Number of protons in an element, often denoted by the letter Z.

428

Photon

Unit of energy in the form of light equal to hf, where h is Planck's constant and f is the frequency of radiation.

429

Atomic Absorption Spectrum

Spectrum of certain absorbed wavelengths of light corresponding to an atom's spectrum of emitted frequencies of light.

430

Mass Number

Sum of the protons and neutrons in an element, often denoted by the letter A.

431

Pfund Series

Set of spectral lines resulting when a hydrogen atom undergoes a transition from energy levels n>5 to n=5.

432

Heisenberg Uncertainty Principle

Quantum mechanical idea that we cannot measure the exact momentum and position of an orbiting electron simultaneously. The more accurate we measure an electron's momentum, the less we know about it's exact position.

433

Principal Quantum Number

First quantum number, designated as the letter "n." It takes on any positive integer value and describes an electron's energy level. An electron with a higher n value is at a higher energy state.

434

Balmer Series

A set of spectral lines that appear in the visible light region when a hydrogen atom undergoes a transition from energy levels n>2 to n=2.

435

Paschen Series

Set of spectral lines resulting when a hydrogen atom undergoes a transition from energy levels n≥4 to n=3.

436

Bohr Model

Atomic model postulating that electrons are located in discrete circular orbits about the nucleus. Electrostatic force between the positive nucleus and negative electron acts as the centripetal force keeping the electron in orbit, like a satellite around the Earth.

437

Lyman Series

Set of spectral lines appearing in the UV region when a hydrogen atom undergoes a transition from energy levels n>1 to n=1.

438

Quantum Mechanics

Study of physics at the atomic level where energy is quantized in discrete, rather than continuous, levels.

439

Atomic Mass Unit

Unit of mass equal to 1/12 gram of a carbon-12 atom, roughly equal to the mass of one proton.

440

Electron

Subatomic particle orbiting the nucleus with a charge of -1. It has a negligible mass and is often denoted as e^-1.

441

Isoelectronic

Two different elements that share the same electronic configuration (ex K+, Ar)

442

Quantum Numbers

Set of four numbers used to describe an electron's energy state (position and energy).

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Valence Electrons

Electrons occupying the outermost electron shell of an atom, participating in chemical bonds. Atoms with the same number of valence electrons tend to have similar properties (families in the Periodic Table).

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Pauli Exclusion Principle

No two electrons in an atom can have the same set of four quantum number values.

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Atomic Weight

Average mass of all naturally-occurring isotopes of a given element, measured in AMU.

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Atom

Basic building block of all matter in the universe. Made up of 3 main components: protons, neutrons and electrons.

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Paramagnetic

Atom or substance containing unpaired electrons and is consequently attracted by a magnet.

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Magnetic Quantum Number

Third quantum number, designated as ml. Describes a particular orbital within a subshell where an electron is very likely to be found. Possible values are integers in the -1 to 1 range, including 0.

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Aufbau Principle

Electrons fill an atom in order of increasing energy level, starting at the bottom.

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Isotopes

Atoms that share the same atomic number (Z) but have a different number of neutrons.

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Diamagnetic

An atom or substance containing no unpaired electrons and is consequently repelled by a magnet.

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Nucleus

Dense, positively charged center of an atom, containing protons and neutrons.

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Proton

Subatomic particle with a charge of +1 and 1.0073 amu.

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Spin Quantum Number

Fourth quantum number. Designated as ms. It specifies an electron's intrinsic spin or angular momentum value in an orbital. There can be no more than two electrons per orbital, the value of ms can only be +1/2 or -1/2. According to Hund's Rule, the electrons are likely to be in parallel spin, hence there's no 0.

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Orbital

Three-dimensional region about the nucleus where a rapidly orbiting electron is likely to be found. Each level has a unique assignment for n, l, and ml quantum numbers.

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Hund's Rule

Electrons will first fill equal-energy orbitals of a subshell unpaired and with parallel spins before being coupled with other electrons of opposite spins in the same orbital. This maximizes the number of half-filled orbitals allowing for the most stable distribution of electrons in a subshell.

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Neutron

Subatomic particle with zero electric charge, slightly heavier than a proton.

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a sytm reaches its maxim entropy at?

equilibrium

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what are units of entropy?

j/K

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once a react commences a standard state condit?

no longer holds

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J/K are the units for ?

entropy

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once a react commences Keq must be replac by?

reaction quotient Q

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adiabatic proc is when?

no heat exch occurs

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if ΔG is zerothe system is?

in a state of equil

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heat is a form of e that transfers bec of?

a temper diff betw syst and its surroundings

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heat is a?

form of energy that can leave or enter a syst

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for any given nonspont react or proc, the reverse react or proc will be?

sponaneous

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ΔSuniv >

0

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entropyqrev is ?

the heat added to the syst undergoing a rever process.

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bomb calorimeterno ____________ enters or leaves the system

heat

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entropy is the meas of?

disorder of a system

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1 cal = ________ J

4.184 J

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ΔHreact can be estimated by?

#NAME?

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bond ______________ is always exothermic

formation

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standard condition for enthalpy etc is?

25 C, 1 atm

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macroscopic prop include?

temp Tpress Pvolume V

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a solid has a lower entropy than a ?

gas

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for a revers procc ΔSuniv is?

zero

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isolated system can?

it can not exch e or matter with surroundings

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heat absorbed or given off in a process is =?

q=mc chT

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const vol calorimetrywhat dev do you use?

bomb calorimeter

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give example of an open syst?

pot of a boiling water

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bond e can be used to estim the ?

enthalpies of reactions

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which 2 processess are common?

isothermal and isobaric. temp and ress are easy to contr

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H2 (g) ---> 2H (g) ΔH= will be?

positive

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ΔG=?

ΔG=ΔH-TΔS

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only ____________ in H can be measured

Δ

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standard free energy of format of any elem in its most stable form is?

zero

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no heat exch occurs in a __________ process

adiabatic

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heat lost by a system to surrounding is?

negative

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enthalpy is?

heat changes at constant pressure

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what is the equat for ΔG from equil const?

ΔG=-RTlnKeq

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thermody and kineticsspontan does not necessaril mean?

instantaneous

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in the equilbr state, free e is?

at a minimum

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what are the req for relatively easy measure of ΔH?

that react be fast and spontan

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joules and calories are units for?

heat

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ΔG react=(sum of ΔGf of prod)- ??

#ERROR!:parse

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Gibbs free energy combines what 2 fact?

ΔH and ΔS

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s is

entropy

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ex of a closed syst

steam radiator