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MCAT Chemistry Flashcards

1
Q

Cathode

A

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

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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

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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.

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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

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5
Q

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

A

Positive; Increased

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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

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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

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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.

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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.

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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

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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

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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)

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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)

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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

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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

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16
Q

Molarity - general equation

A

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

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17
Q

Normality

A

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

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18
Q

Anode

A

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

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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

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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

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21
Q

Alkene

A

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

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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

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23
Q

Positive Ion: which elements tend to form them?

A

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

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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)

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25
Q

Dalton’s Law of Partial Pressures

A

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

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26
Q

Oxidation

A

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

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27
Q

Bronsted Theory

A

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

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28
Q

Metals. List 6 characteristics

A

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

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29
Q

STP

A

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

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30
Q

Scientific notation

A

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

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31
Q

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

A

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

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32
Q

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

A

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

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33
Q

Carboxylic Acid

A

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

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34
Q

Negative Ion: which elements tend to form them?

A

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

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35
Q

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

A

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

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36
Q

Aromatic compounds

A

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

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37
Q

Beta particle

A

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

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38
Q

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

A

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

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39
Q

Coordinate covalent bond

A

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

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40
Q

Electrolyte (give 3 examples)

A

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

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41
Q

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

A

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

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42
Q

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

A

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

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43
Q

Isotope

A

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)

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44
Q

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

A

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

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45
Q

Endothermic

A

A chemical reaction which absorbs heat

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46
Q

Double replacement

A

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

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47
Q

Second Law of Thermodynamics

A

The entropy of the universe increases for any spontaneous process.

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48
Q

Atomic Mass

A

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)

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49
Q

Ion (provide examples)

A

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

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50
Q

Explain the rules for dividing the numbers in scientific notation

A

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

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51
Q

Electronegativity

A

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)

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52
Q

Six common characteristics of acids

A

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

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53
Q

Fission

A

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

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54
Q

Diatomic; list 7 diatomic elements found in nature

A

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

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55
Q

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

A

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

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56
Q

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

A

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

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57
Q

Heat of Fusion (value for water)

A

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

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58
Q

Non-metals. List 7 characteristics

A

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

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59
Q

Transition element

A

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

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60
Q

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

A

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

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61
Q

Valence electroncs

A

The electrons found in the outermost energy level of an atom

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62
Q

Solubility curves

A

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

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63
Q

Synthesis or combination

A

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

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64
Q

Equilibrium

A

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

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65
Q

Oxidizing agent

A

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

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66
Q

Faraday

A

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

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67
Q

Electron configuration

A

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

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68
Q

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

A

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

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69
Q

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

A

Ksp = [Ag+][Cl-]

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70
Q

Metallic Bonds

A

A sea of electrons surrounding positive metal ions

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71
Q

Physical properties

A

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

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72
Q

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

A

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-)

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73
Q

Specific Rate Constant

A

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

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74
Q

Mole

A

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)

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75
Q

Amino Acid

A

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

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76
Q

Conjugate Base - Write conjugate base of HCl

A

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

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77
Q

Law of Multiple Proportions

A

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

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78
Q

Ionization energy

A

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

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79
Q

Molecule

A

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

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80
Q

Catalyst

A

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.

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81
Q

Equivalent weight

A

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

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82
Q

Write the equation to calculate the atomic mass of an element

A

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

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83
Q

Dynamic equilibrium

A

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

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84
Q

Isomer

A

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

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85
Q

Aldehyde

A

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

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86
Q

List five factors that control reaction rate

A

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

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87
Q

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

A

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

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88
Q

Describe the change in electronegativity across the periodic table

A

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

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89
Q

Solvent (provide example

A

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

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90
Q

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

A

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

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91
Q

List four quantum numbers and their symbols

A

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

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92
Q

Octet Rule

A

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

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93
Q

Single Replacement

A

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

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94
Q

Alcohol

A

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

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95
Q

Mass number

A

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

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96
Q

Nonelectrolyte (give 2 examples)

A

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

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97
Q

Alkyne

A

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

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98
Q

Seven common characteristics of bases

A

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

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99
Q

Vapor Pressure

A

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

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100
Q

Decomposition

A

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

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101
Q

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

A

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

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102
Q

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

A

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

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103
Q

Alkane

A

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

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104
Q

Conjugate Acid - Write conjugate acid of I-

A

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

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105
Q

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

A

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)

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106
Q

Quantum Number

A

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

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107
Q

Describe the change in atomic radius across the periodic table

A

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

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108
Q

List some basic facts about suspensions

A

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

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109
Q

Activation energy

A

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

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110
Q

Gay-Lussac’s Law of Combining Gas Volumes

A

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

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111
Q

Element. List some examples

A

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

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112
Q

Density: Write the general equation and three standard units

A

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

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113
Q

List properties of ionic substances

A

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

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114
Q

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

A

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

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115
Q

Nonpolar covalent bond

A

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.

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116
Q

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

A

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

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117
Q

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

A

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

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118
Q

Half Reaction

A

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

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119
Q

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

A

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

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120
Q

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

A

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.

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121
Q

Alpha particle

A

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

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122
Q

Fusion

A

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

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123
Q

Describe an ionic bond in terms of electronegativity difference

A

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.

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124
Q

Thermochemical equation

A

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

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125
Q

Conversion Factor

A

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)

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126
Q

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

A

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

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127
Q

Polar covalent bond

A

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.

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128
Q

Gram-equivalent weight

A

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

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129
Q

Compare potential energy and kinetic energy

A

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

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130
Q

Which three factors affect the rate of solubility?

A

Pulverizing; stirring; heating

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131
Q

Significant digits

A

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

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132
Q

Resonance structures

A

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

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133
Q

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

A

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

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134
Q

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

A

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

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135
Q

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

A

22.4 L

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136
Q

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

A

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

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137
Q

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

A

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

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138
Q

10 prefixes used to name covalent compounds

A

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

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139
Q

Addition reaction

A

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

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140
Q

Substitution reaction. Which elements commonly substitute in alkanes?

A

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

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141
Q

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

A

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.

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142
Q

Charles Law

A

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

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143
Q

Hydrogenation

A

The process of adding hydrogen to an unsaturated hydrocarbon

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144
Q

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

A

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

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145
Q

Reactants; Products

A

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

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146
Q

Column or Family

A

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

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147
Q

Cycloalkane

A

An alkane which has a ring structure instead of a chain

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148
Q

Matter

A

A substance that occupies space has mass.

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149
Q

Hybrid orbitals (list three types)

A

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

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150
Q

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

A

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

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151
Q

The Law of Conservation of Mass

A

During a chemical reaction, matter is neither created nor destroyed

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152
Q

Avogadro’s Number

A

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)

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153
Q

Quantum

A

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

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154
Q

Cation and Anion (provide examples)

A

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

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155
Q

Ground State

A

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

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156
Q

Heat of Vaporization (value for water?)

A

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

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157
Q

List the names of the four sublevels and their electron capacities

A

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

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158
Q

Explain the rules for multiplying the numbers in scientific notation

A

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

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159
Q

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

A

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

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160
Q

Energy

A

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)

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161
Q

Solute (provide example)

A

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

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162
Q

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

A

%=(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

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163
Q

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

A

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

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164
Q

Ether

A

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

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165
Q

Reducing agent

A

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

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166
Q

Solubility Product Constant (Ksp)

A

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

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167
Q

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

A

Negative; Decreased

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168
Q

List some basic facts about carbon bonding

A

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

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169
Q

How is the reaction rate related to concentration?

A

Reaction rate is directly proportional to the concentrations of reactants

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170
Q

Dehydrogenation

A

The process of removing hydrogen from a hydrocarbon

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171
Q

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

A

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

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172
Q

Chemical change

A

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

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173
Q

Compound. List some examples

A

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

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174
Q

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

A

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

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175
Q

Enthalpy

A

The heat content of a system

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176
Q

Arhennius Theory

A

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

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177
Q

Enthalpy change (state the equation)

A

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

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178
Q

Write the general equation to calculate ΔH for a chemical reaction

A

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

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179
Q

Transmutation

A

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

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180
Q

Compare primary and secondary alcohols

A

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.

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181
Q

Polymerization

A

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

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182
Q

Describe the change in metallic activity across the periodic table

A

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

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183
Q

Physical change

A

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

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184
Q

Exothermic

A

A chemical reaction which evolves heat

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185
Q

Are concentrations of product and reactant equal at equilibrium?

A

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

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186
Q

Ideal Gas Law

A

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

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187
Q

Amine

A

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

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188
Q

Oxidation number

A

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

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189
Q

Molar Heat of Formation

A

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

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190
Q

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

A

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

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191
Q

Molecular Mass

A

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

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192
Q

Compare Weight vs. Mass

A

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

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193
Q

Three postulates of Dalton’s Atomic Theory

A

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

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194
Q

Gamma radiation

A

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

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195
Q

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

A

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

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196
Q

Molecular Formula. What is the molecular formula for ammonia?

A

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)

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197
Q

Homogeneous and Heterogeneous mixtures. Provide examples.

A

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

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198
Q

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

A

IA→+1; IIA→+2; VIA→-2; VIIA→-1

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199
Q

Ionization constant. Write the expression for the ionization of acid “HA”

A

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

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200
Q

How does a change in temperature affect equilibrium

A

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.

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201
Q

Chemical properties

A

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

202
Q

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

A

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

203
Q

Law of Mass Action

A

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

204
Q

Mixture. List some examples.

A

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

205
Q

Combined Gas Law

A

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

206
Q

Neutralization Reaction - Write equation for hydrochloric acid and sodium hydroxide

A

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

207
Q

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

A

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

208
Q

Hess’ Law

A

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

209
Q

Valence

A

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

210
Q

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

A

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

211
Q

Double and Triple covalent bond

A

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

212
Q

Boyles Law

A

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

213
Q

Reversible reaction

A

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

214
Q

Solution (provide example)

A

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

215
Q

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

A

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

216
Q

Period

A

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
Q

Absorption Spectrum

A

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

218
Q

The Law of Constant Composition

A

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
Q

Reduction

A

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

220
Q

Half-life

A

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
Q

Three general rules for determining oxidation numbers

A

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
Q

Ester

A

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

223
Q

Excited State

A

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

224
Q

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

A

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

225
Q

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

A

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

226
Q

List some basic facts about colloids

A

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
Q

Lewis Theory

A

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

228
Q

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

A

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
Q

List properties of molecular substances

A

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

230
Q

Metalloid. List 5 examples.

A

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

231
Q

Normal Boiling Point

A

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

232
Q

Describe hybrid bonding in water, ammonia, methane

A

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

233
Q

Describe the change in Ionization Energy across the periodic table

A

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

234
Q

Ketone

A

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

235
Q

How are pH and pOH of a solution related?

A

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

236
Q

Sigma bonds & Pi bonds

A

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

237
Q

Ionization Constant

A

Kw = 1 x 10-14 at 25C

238
Q

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

A

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

239
Q

Avogadro’s Law re. gases

A

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

240
Q

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

A

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

241
Q

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

A

known volknown vol o

242
Q

ph=

A

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

243
Q

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

A

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

244
Q

a b equiv is equal to?

A

one mole of OH- ions

245
Q

NO2 - is the conjugate b of ?

A

HNO2

246
Q

what combinations of a and b are possible?

A
  1. str a and str b2. str a and weak b3. weak a and str b4. weak a and weak b
247
Q

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

A

t

248
Q

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?

A

its initial concentr 2.0 M- X, the amount dissociated

249
Q

NO2-

A

nitrite

250
Q

give 2 examples of buffers?

A

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
Q

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

A

0 and 1

252
Q

bronst lowry a and ba isb is?

A

acid donate h+bases accepts H+

253
Q

name some strong bases

A

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

254
Q

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?

A

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

255
Q

t or fwater is amphoteric

A

t

256
Q

weak a with a str b the resulting solut is?

A

basic.

257
Q

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

A

amphoteric

258
Q

anion ends in ite it will be __________ acid

A

ous acid

259
Q

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

A

acidic

260
Q

the hydroxides of which metals are usually amphoteric?

A

Al, Zn, Pb, and Cr

261
Q

Kw= [H+][OH-]=

A

41561

262
Q

Henderson- Hasselbach equation is?

A

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

263
Q

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

A

amphoteric

264
Q

H3O+ is the conjugate acid of the base?

A

H2O

265
Q

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

A

H2C03

266
Q

a neutrialization react is when?

A

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

267
Q

this is titration of a __________ base with an acid

A

polyprotic base

268
Q

how do you calcul equival weight?

A

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

269
Q

litmus paper turns blue in?

A

bases

270
Q

H2SO4 is a. monoproticb. diproticc. triprotic

A

b. diprotic

271
Q

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

A
  1. graphical method- use ph meter2. watching for a color ch
272
Q

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

A

C032 istitrated to HC03.

273
Q

if Ka is large then Kb?

A

will be small

274
Q

weak acids and bases only partly?

A

dissociate in aq sol

275
Q

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

A

in a titration, half-way to the equivalent point

276
Q

amphoteric means?

A

can act as an acid or a base

277
Q

nitrate

A

NO3-

278
Q

name 3 polyvalent acids?

A

H2SO4H3PO4H2CO3

279
Q

name of an acid is related to its?

A

parent anion

280
Q

arrhenius def of acid and baseacid is?base is?

A

acid prod H+base prod OH-

281
Q

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

A

m-0

282
Q

why dont indicators change the equival point

A

bec they are in low concetr

283
Q

ClO3-

A

chlorate

284
Q

bromide becomes ________________ acid

A

hydrobromic acid

285
Q

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

A

acidosis

286
Q

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

A

f

287
Q

what is the shortcoming of arrhenius def?

A

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

288
Q

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

A

will react w the water solvent, reforming the weak b

289
Q

litmus paper turns red in ?

A

acid

290
Q

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?

A

write the express for the acid dissoc const

291
Q

the henderson-Hasselbalch equation is used for?

A

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
Q

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?

A

NaVa=NbVb

293
Q

why are buffers useful?

A

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

294
Q

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

A

water

295
Q

name some strong acids

A

HClO4- perchloric acidHNO3 nitric acidH2SO4 sulfuric acidHCl hydrochloric acid

296
Q

CLO-

A

hypochlorite

297
Q

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

A

a proton

298
Q

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

A

relative strenght of the reactants

299
Q

HNO2

A

nitrous acid

300
Q

oxid or reduc agents act amphoteric when they ?

A

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

301
Q

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

A

end pointequivalence point

302
Q

weak bases are usually not?

A

hydroxides

303
Q

if the conjugate a is strong then the conjugate base?

A

will be weak

304
Q

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

A

CH3COOH (aq)—> H+ (aq) + CH3COO- (aq)

305
Q

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=?

A

ERROR!:parse

306
Q

Ka and _______ are inversely related

A

Kb

307
Q

indicators are?

A

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

308
Q

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

A

ph=pKa, bec log1=0

309
Q

lewis def of a and bacidbase?

A

acid is electr pair acceptbase is electr pair donator.

310
Q

HNO3

A

nitric acid

311
Q

a conjugate base is formed when?

A

an a loses a proton

312
Q

HClO2

A

chorous acid

313
Q

Ka measures?

A

the degree to which an a dissoc

314
Q

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

A

it will only change the ph a little

315
Q

what is equival weight of H2SO4?

A

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

316
Q

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

A

that 2.0-x is approx = to 2.0

317
Q

in any aq solut the H2O does what?

A

dissociates slightly

318
Q

t or fthe equival point is always at ph of 7

A

fonly for str a with a str base

319
Q

bronst lowry acid and b occur in pairs called?

A

conjugate a b pairs

320
Q

the partially dissoc conjug base of a polyprotic acid is?

A

usually amphoteric

321
Q

flouride bec acid?

A

hydrofluoric acid

322
Q

what are the react for HCL reacting with NH3?

A

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

323
Q

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?

A

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

324
Q

this is a titration of a ____________ with a strong base

A

weak acid

325
Q

what is the equation for the dissoc of H2SO4?

A

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

326
Q

m-lognthe larger n is the larger ?

A

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

327
Q

bases turn litmus paper?

A

blue

328
Q

t or fneutralization reactions usually go to completion

A

t

329
Q

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

A

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

330
Q

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

A

conjugate a HCO3-conjugate b CO3-

331
Q

how can you estimate a logn x 10-m =

A

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

332
Q

Ka=

A

ERROR!:parse

333
Q

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

A

neutralization

334
Q

the weaker the base the ____________ the Kb

A

smaller

335
Q

ph + pOH=?

A

14

336
Q

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. ?

A

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

337
Q

AlCl3 is a _______________ acid but not a ________________ acid

A

lewis but not a bronst lowry

338
Q

strong acids and bases are those that?

A

completely dissociate into their component ions in aq solution

339
Q

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

A

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
Q

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

A

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

341
Q

NO3-

A

nitrate

342
Q

acids from oxyanions are called?

A

oxyacids

343
Q

nitric acid

A

HNO3

344
Q

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

A

that H2SO4 is diprotic

345
Q

pOH is=

A

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

346
Q

hypochlorite

A

CLO-

347
Q

HCl is?

A

hydrochloric acid

348
Q

hydrolysis (in terms of a and b) is?

A

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

349
Q

Ka and Kb are _____________ related

A

inversely

350
Q

anion ends in ate, it will be ____________ acid

A

ic acid

351
Q

weaker the acid, the smaller the?

A

Ka

352
Q

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

A

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

353
Q

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

A

C032-

354
Q

MnO4- is what anion?

A

permanganate, even though there are no manganate or manganite ions

355
Q

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

A

t

356
Q

AlCl3 can accept a?

A

electron pair

357
Q

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?

A

write the equilibrium reaction

358
Q

Kw=

A

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

359
Q

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

A

Ka of HCO3-

360
Q

pH above 7 indicates an excess of?

A

OH- ions

361
Q

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

A

arhhenius

362
Q

pure water _____________ is equal to ________________

A

H+ is equal to OH-

363
Q

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?

A

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

364
Q

HClO3

A

chloric acid

365
Q

2M H3PO4 would have be _________ N

A

6

366
Q

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

A

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
Q

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

A

Ka X Kb=Kw=l x10-14

368
Q

HF is?

A

hydrofluoric acid

369
Q

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

A

pH. by carefully choosing a weak acid and its salt

370
Q
  1. 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
A

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
Q

an a equiv is equal to ?

A

one mole of H+

372
Q

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

A

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

373
Q

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

A

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

374
Q

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

A

t

375
Q

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

A

amphoteric

376
Q

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

A

H2C03

377
Q

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

A

equivalence point

378
Q

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

A

its salt

379
Q

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

A

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
Q

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

A

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
Q

HNO3 is?

A

nitric acid

382
Q

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

A

weak acid w a str base

383
Q

ClO4-

A

perchlorate

384
Q

Kw is the?

A

water dissociation constant

385
Q

HBr is ?

A

hydrobromic acid

386
Q

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

A

conjugate base

387
Q

HClO

A

hypochlorous acid

388
Q

the weaker the acid, the _____________ the Ka

A

smaller

389
Q

titration is used to ?

A

determ the molarit of an a or a b

390
Q

perchloric acid

A

HClO4

391
Q

in the early part of the curve, the __________ species

A

acidic species

392
Q

chlorous acid

A

HClO2

393
Q

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

A

t

394
Q

math reminderlog(xy)=?

A

logx + logy

395
Q

equival point is when the?

A

number of acid equival equals the number of base equival

396
Q

bronsted lowryan amphoteric species can?

A

gain or lose a proton

397
Q

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

A

C032- and HC03,first buffer region

398
Q

sulfuric acid is?

A

H2SO4

399
Q

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

A

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

400
Q

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

A

hydrolysis

401
Q

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

A

H2C03 and HC03

402
Q

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

A

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

403
Q

Kb is the

A

base dissociation constant

404
Q

ClO2-

A

chlorite

405
Q

pH below 7 is

A

acidic

406
Q

equation for neutralization

A

HA + BOH —> BA + H2O

407
Q

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?

A

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

408
Q

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

A

NaVa=NbVb

409
Q

the equival point is in the _______ range

A

basic

410
Q

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?

A

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

411
Q

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

A

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

412
Q

chlorate

A

CLO3-

413
Q

one mole of H2SO4 can prod ______ acid equiv

A

2

414
Q

when titrating polyprotic a or b there are several?

A

equivalence points

415
Q

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

A

equivalence point

416
Q

nitrite

A

NO2-

417
Q

In region IV the flatregion is the?

A

second buffer region of the titration curve.

418
Q

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

A

the acid dissoc const Ka

419
Q

hydrofluoric acid is?

A

HF

420
Q

perchlorate

A

ClO4-

421
Q

anion ending in -ide will become ______________ acid

A

hydro ic

422
Q

point at which the indicat ch color is called the?

A

end point

423
Q

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

A

less than that

424
Q

Azimuthal Quantum Number

A

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
Q

Atomic Emission Spectra

A

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

426
Q

Electron Configuration

A

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
Q

Atomic Number

A

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

428
Q

Photon

A

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
Q

Atomic Absorption Spectrum

A

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

430
Q

Mass Number

A

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

431
Q

Pfund Series

A

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

432
Q

Heisenberg Uncertainty Principle

A

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
Q

Principal Quantum Number

A

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
Q

Balmer Series

A

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
Q

Paschen Series

A

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

436
Q

Bohr Model

A

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
Q

Lyman Series

A

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
Q

Quantum Mechanics

A

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

439
Q

Atomic Mass Unit

A

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

440
Q

Electron

A

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

441
Q

Isoelectronic

A

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

442
Q

Quantum Numbers

A

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

443
Q

Valence Electrons

A

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).

444
Q

Pauli Exclusion Principle

A

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

445
Q

Atomic Weight

A

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

446
Q

Atom

A

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

447
Q

Paramagnetic

A

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

448
Q

Magnetic Quantum Number

A

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
Q

Aufbau Principle

A

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

450
Q

Isotopes

A

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

451
Q

Diamagnetic

A

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

452
Q

Nucleus

A

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

453
Q

Proton

A

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

454
Q

Spin Quantum Number

A

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.

455
Q

Orbital

A

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.

456
Q

Hund’s Rule

A

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.

457
Q

Neutron

A

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

458
Q

a sytm reaches its maxim entropy at?

A

equilibrium

459
Q

what are units of entropy?

A

j/K

460
Q

once a react commences a standard state condit?

A

no longer holds

461
Q

J/K are the units for ?

A

entropy

462
Q

once a react commences Keq must be replac by?

A

reaction quotient Q

463
Q

adiabatic proc is when?

A

no heat exch occurs

464
Q

if ΔG is zerothe system is?

A

in a state of equil

465
Q

heat is a form of e that transfers bec of?

A

a temper diff betw syst and its surroundings

466
Q

heat is a?

A

form of energy that can leave or enter a syst

467
Q

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

A

sponaneous

468
Q

ΔSuniv >

A

0

469
Q

entropyqrev is ?

A

the heat added to the syst undergoing a rever process.

470
Q

bomb calorimeterno ____________ enters or leaves the system

A

heat

471
Q

entropy is the meas of?

A

disorder of a system

472
Q

1 cal = ________ J

A

4.184 J

473
Q

ΔHreact can be estimated by?

A

NAME?

474
Q

bond ______________ is always exothermic

A

formation

475
Q

standard condition for enthalpy etc is?

A

25 C, 1 atm

476
Q

macroscopic prop include?

A

temp Tpress Pvolume V

477
Q

a solid has a lower entropy than a ?

A

gas

478
Q

for a revers procc ΔSuniv is?

A

zero

479
Q

isolated system can?

A

it can not exch e or matter with surroundings

480
Q

heat absorbed or given off in a process is =?

A

q=mc chT

481
Q

const vol calorimetrywhat dev do you use?

A

bomb calorimeter

482
Q

give example of an open syst?

A

pot of a boiling water

483
Q

bond e can be used to estim the ?

A

enthalpies of reactions

484
Q

which 2 processess are common?

A

isothermal and isobaric. temp and ress are easy to contr

485
Q

H2 (g) —> 2H (g) ΔH= will be?

A

positive

486
Q

ΔG=?

A

ΔG=ΔH-TΔS

487
Q

only ____________ in H can be measured

A

Δ

488
Q

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

A

zero

489
Q

no heat exch occurs in a __________ process

A

adiabatic

490
Q

heat lost by a system to surrounding is?

A

negative

491
Q

enthalpy is?

A

heat changes at constant pressure

492
Q

what is the equat for ΔG from equil const?

A

ΔG=-RTlnKeq

493
Q

thermody and kineticsspontan does not necessaril mean?

A

instantaneous

494
Q

in the equilbr state, free e is?

A

at a minimum

495
Q

what are the req for relatively easy measure of ΔH?

A

that react be fast and spontan

496
Q

joules and calories are units for?

A

heat

497
Q

ΔG react=(sum of ΔGf of prod)- ??

A

ERROR!:parse

498
Q

Gibbs free energy combines what 2 fact?

A

ΔH and ΔS

499
Q

s is

A

entropy

500
Q

ex of a closed syst

A

steam radiator

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