MCAT Chemistry Flashcards
(500 cards)
1
Q
Cathode
A
A negatively charged electrode which attracts cations. Where reduction takes place.
2
Q
How does a concentration change affect equilibrium?
A
If one of the substances is added or removed, all the concentrations of substances adjust to a new equilibrium with the same Keq
3
Q
Electronegativity difference
A
A number found by taking the difference between the electronegativities of two atoms in a bond. Its value determines the type of bond.
4
Q
List some basic facts about solutions
A
Particle size less than 1 mmicron; Clear (may be colored); Particles don’t settle; Can pass through membranes; Particles not visible
5
Q
Endothermic Reaction: ΔH is (positive or negative); Enthalphy is (increased or decreased)
A
Positive; Increased
6
Q
Atomic Mass Unit
A
the mass of a proton or neutron is equal to 1 atomic mass unit. Symbol - “amu”; 1 amu=1.66x10^-24
7
Q
Hydrogen bonding
A
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
Q
Phase equilibrium
A
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
Q
Emission Spectrum
A
A bright line spectrum formed when energy absorbed by an element is emitted at specific wavelengths. Each element has a unique spectrum.
10
Q
Le Chatelier’s Principle
A
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
Q
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?
A
1) 1.86C for each molal of particles of solute; 2) 0.51C for each molal of particles of solute
12
Q
Electrode Potential
A
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
Q
State the collision theory of reaction rates
A
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
Q
Use conversion factors to solve: 1) How many kilograms in 2000 g? 2) How many feet in 60”?
A
1) (2000g)(1kg/1000g) = 2kg; 2) (60”)(1’/12”) = 5 ft
15
Q
Describe the change in nonmetallic activity across the periodic table
A
Non-metallic activity increases across the table (L to R) and decreases down the columns
16
Q
Molarity - general equation
A
The number of moles of solute it a liter of solution; M = n/L
17
Q
Normality
A
The number of gram-equivalent weights in a liter of solution
18
Q
Anode
A
A positively charged electrode which attracts anions. Where oxidation takes place.
19
Q
Triple Point
A
The only temperature and pressure combination at which the 3 phases of a substance (solid, liquid, gas) can co-exist in equilibrium
20
Q
Which sublevels are present in energy levels 1, 2, 3, and 4-7?
A
1: s; 2: s, p; 3: s, p, d; 4-7: s, p, d, f
21
Q
Alkene
A
A series of hydrocarbons containing at least one double covalent bond (CnH2n)
22
Q
How does a pressure change affect equilibrium?
A
Only in reactions where gases are involved. The reaction will shift to oppose pressure change, resulting in fewer moles of gas particles
23
Q
Positive Ion: which elements tend to form them?
A
Metals tend to form positive ions by losing electrons (Na → Na+ e-)
24
Q
What are the two driving forces that control reactions?
A
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
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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
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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
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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.
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Isoelectronic
Two different elements that share the same electronic configuration (ex K+, Ar)
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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.
446
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.
448
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.
449
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
460
once a react commences a standard state condit?
no longer holds
461
J/K are the units for ?
entropy
462
once a react commences Keq must be replac by?
reaction quotient Q
463
adiabatic proc is when?
no heat exch occurs
464
if ΔG is zerothe system is?
in a state of equil
465
heat is a form of e that transfers bec of?
a temper diff betw syst and its surroundings
466
heat is a?
form of energy that can leave or enter a syst
467
for any given nonspont react or proc, the reverse react or proc will be?
sponaneous
468
ΔSuniv >
0
469
entropyqrev is ?
the heat added to the syst undergoing a rever process.
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bomb calorimeterno ____________ enters or leaves the system
heat
471
entropy is the meas of?
disorder of a system
472
1 cal = ________ J
4.184 J
473
ΔHreact can be estimated by?
#NAME?
474
bond ______________ is always exothermic
formation
475
standard condition for enthalpy etc is?
25 C, 1 atm
476
macroscopic prop include?
temp Tpress Pvolume V
477
a solid has a lower entropy than a ?
gas
478
for a revers procc ΔSuniv is?
zero
479
isolated system can?
it can not exch e or matter with surroundings
480
heat absorbed or given off in a process is =?
q=mc chT
481
const vol calorimetrywhat dev do you use?
bomb calorimeter
482
give example of an open syst?
pot of a boiling water
483
bond e can be used to estim the ?
enthalpies of reactions
484
which 2 processess are common?
isothermal and isobaric. temp and ress are easy to contr
485
H2 (g) ---> 2H (g) ΔH= will be?
positive
486
ΔG=?
ΔG=ΔH-TΔS
487
only ____________ in H can be measured
Δ
488
standard free energy of format of any elem in its most stable form is?
zero
489
no heat exch occurs in a __________ process
adiabatic
490
heat lost by a system to surrounding is?
negative
491
enthalpy is?
heat changes at constant pressure
492
what is the equat for ΔG from equil const?
ΔG=-RTlnKeq
493
thermody and kineticsspontan does not necessaril mean?
instantaneous
494
in the equilbr state, free e is?
at a minimum
495
what are the req for relatively easy measure of ΔH?
that react be fast and spontan
496
joules and calories are units for?
heat
497
ΔG react=(sum of ΔGf of prod)- ??
#ERROR!:parse
498
Gibbs free energy combines what 2 fact?
ΔH and ΔS
499
s is
entropy
500
ex of a closed syst
steam radiator
