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Chem 20 > Chem 20 Final Exam > Flashcards

Chem 20 Final Exam Flashcards

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1
Q

Electro negativity

A

Used to describe the relative ability of an atom to attract a pair of bonding electrons in its valence level
Metals= low electronegativities
Nonmetals= high electronegativities

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2
Q

Valence Electrons

A

Electrons occupying the highest energy level of an atom

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3
Q

Valence energy level

A

The outer most energy level of an atom

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4
Q

Orbital

A

A region in space around an atoms nucleus in which an electron may exist

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5
Q

Valence Orbital

A

Volumes of space that can be occupied by electrons in an atoms highest energy level
First energy level: one orbital with a max of 2 electrons
Energy levels above the first have room enough for 4 orbitals (contain 0,1 or 2 electrons)

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6
Q

Octet Rule

A

A maximum of 8 electrons may occupy a valence energy level

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7
Q

Bonding electron

A

An atom with a valence orbital that is occupied by one single electron can share that electron with another atom

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8
Q

Lone pairs

A

A full valence orbital, occupied by 2 electrons, repels electrons in nearby orbitals= 2 electrons occupying the same orbital

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9
Q

Ionic bond

A

Is the attraction that results from a positive ion (metal) and a negative ion (nonmetal), 2 ions are then attracted to each other, a transfer of electrons occurs in an ionic bond

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10
Q

Covalent bond

A

Is the bond that results from the electrostatic attraction between the electrons of one atom to the nucleus of another and vice versa

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11
Q

Single bond

A

One pair of electrons is shared between 2 atoms

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12
Q

Double bond

A

2 pairs of electrons are shared between 2 atoms

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13
Q

Triple bond

A

3 pairs of electrons are shared between the 2 atoms

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14
Q

Polar molecule

A

A molecule that has an overall charge separation, one end of the molecule is positive and other end of the molecule is negative

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15
Q

Dipole

A

Is a result of unequal sharing of electrons in a molecule, from a difference in electronegativities

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16
Q

Nonpolar molecule

A

Has no net charge separation

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17
Q

Dipole-dipole forces

A

Polar molecules have dipoles, the attraction between dipoles is called dipole dipole force. The positive end of one molecule will be attracted to the negative end of a neighbouring molecule and will extend in all directions

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18
Q

London Dispersion Forces

A

Result from the movement of the electrons in the molecule which generates temporary positive and negative regions in the molecule. Weak attractive forces that result when the electrons of one molecule are attracted to the positive nuclei of a nearby molecule

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19
Q

Hydrogen Bonding

A

Occurs when hydrogen is bonded to a highly electronegative element (fluorine,oxygen, nitrogen). Bond is strongly polar, highly electronegative atom pulls hydrogens electron away from its nucleus, another molecules lone pair of electrons can now approach the nucleus closely on the side away from its covalent bond
*is the strongest of intermolecular bonds

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20
Q

3 structures showing hydrogen bonding

A

HF(g)
Any molecule with an -OH in its structure
Any molecule with an -NH in its structure
Results in high boiling points as lots of energy must be added in order to break these bonds

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21
Q

Ionic crystals

A

Crystal lattice structure

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22
Q

Metallic crystals

A

Metallic bonding

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23
Q

Molecular crystals

A

Molecules arranged in a regular lattice

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24
Q

Network Covalent crystals

A

Network of covalent bonds

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25
Q

Physical properties of gases

A
  1. Gases have no shape or volume of their own, they assume shape/volume of the container they are in
  2. Are highly compressible (pressure increases>volume of gas decreases)
  3. gases diffuse or spread out (will occupy all space that is available to them)
  4. Gases mix readily with other gases
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26
Q

Pressure

A

A force per unit area (Pa), unit for pressure is kilopascal (kPa)
Pressure at sea level= 101kPa
1atm= 101.325kPa

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27
Q

Temperature

A

Is a measurement of the amount of kinetic energy (molecular motion) a sample has

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28
Q

Standard Ambient Temperature and pressure (SATP)

A

Laboratory conditions
P=101kPa
T=25•C or 298K

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29
Q

Standard Temperature and Pressure (STP)

A

P=101.325kPa

T= 0•C or 273K

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30
Q

Boyles Law

A

As pressure increases, volume decreases provided that all other variables (temp, amount of gas) are kept constant
P1V1=P2V2

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31
Q

Charles Law

A

The higher the temperature of a substance, the greater the molecular motion of the substance. If the temp of a gas increases, the volume of the gas increases. Is valid when the pressure and amount of gas are held constant
V1/T1=V2T2

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32
Q

Absolute Zero and the Kelvin scale

A

Volume vs temp is graphed will result in a straight line, line will cross x-axis at -273•C (gases have no volume at this temp)
-273•C = zero kelvin = 0 K = absolute zero
Convert Celsius to Kelvin add 273, convert Kelvin to Celsius subtract 273
*always use the Kelvin unit

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33
Q

Combined gas Law

A

When boyles, Charles, and Gay-Lussacs laws are combined it states the relationship between pressure, Temperature, and volume of a gas
P1V1/T1=P2V2/T2

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34
Q

Law of combining volumes

A

States that the volume of gaseous reactants and products of chemical reactions are always in simple whole number ratios

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35
Q

Avogadros theory

A

Equal volumes of gas at a constant temperature and pressure contain equal numbers of molecules or atoms
SATP: 24.8L/mol
STP: 22.4L/mol

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36
Q

Ideal gas Law

A

Combining Charles Law, boyles Law, and avogadros theory together
pV=nRT
V=volume(L)
P=pressure(kPa)
T=temp(K)
n=#of moles (mol)
R= universal gas constant (8.314 L•kPa/mol•K)

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37
Q

Solution

A

A homogeneous mixture of 2 or more substances, different substances cannot be mechanically separated or seen (may be solid, liquid or gas)

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38
Q

Solvent

A

The substance present in the largest amount, usually determines the phase of the solution

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39
Q

Solute

A

The substance present in the smallest amount, what is dissolved in the solvent

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40
Q

Aqueous Solution

A

If the solvent is water

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41
Q

Electrolyte

A

A substance that conducts electricity when dissolved in water
Ex: acid, base, soluble ionic solutions

42
Q

Non-electrolyte

A

A substance that does not conduct electricity when dissolved In water
Ex: molecular compounds, insoluble ionic compounds

43
Q

Dissociation

A

Separations of ions from an ionic compound when it dissolved in water, molecular compounds do
Not dissociate I too ions in water

44
Q

Ionization

A

Formation of charged ions from neutral atoms by the action of the solvent, acids ionize inaqueous solution

45
Q

Concentration in Solution

A 
C=n/v
C=mol/L
n=mol
V=L
Concentration is the quantity of solute divided by the quantity of solution
46
Q

Concentration calculations

A

Always include units, make sure units cancel.

Involve quantity of solute, quantity of solution, and concentration of solution

47
Q

Dilution calculations

A

ViCi=VfCf
More concentrated to less concentrated you must add more solvent
The quantity of solute does not change, the mass and # of moles of solute remains constant after dilution

48
Q

Dynamic Equilibrium

A

The rate at which particles dissolve equals the rate at which particles crystallize (2 opposing processes occur at equal rates)

49
Q

Dissolving

A

When a solute dissolved in a solution, the solid particles move into the solvent to form ions

50
Q

Crystallization

A

Some of the dissolved ions crystallize to become solid again

51
Q

Saturated

A

When dynamic Equilibrium is reached and undissolved solute is visible even with stirring (solution will not dissolve any more particles)
Solution contains max amount of solute that can remain dissolved in a given amount of solvent at a specific temp

52
Q

Supersaturated

A

Saturated solution is heated to dissolve extra solute, then cooled to original temp without extra solute crystallizing

53
Q

Unsaturated

A

Solution contains less than the maximum amount of solute than can dissolve at a particular temperature

54
Q

Factors that affect Solubility

A

Temperature (increase in temp will increase solubility of solids, but will decrease solubility of gases)
Pressure (increase in pressure will increase solubility of gases, has no effect on solid or liquids)
Nature of solute or solvent: solutes will dissolve better in some solvents than others, it depends on bonds between solute and solvent

55
Q

Acids empirically defined

A

Taste sour, conduct electricity, react with metals to produce hydrogen gas, turn blue litmus red, pH between 0 and 7

56
Q

Bases empirically defined

A

Taste bitter, conduct electricity, feel soapy/slippery, turn red litmus blue, pH between 7 and 14

57
Q

Arrhenius Theory Acids

A

Acids produce hydrogen ions when they are dissolved in water, acids are substances that ionize in aqueous solution to form hydrogen ions, hydrogen ions are responsible for the acidic properties of these solutions

58
Q

Arrhenius theory bases

A

Produce hydroxide ions when they are dissolved in water, are substances that dissociate to form hydroxide ions in aqueous solution, hydroxide ions are responsible for basic properties

59
Q

Hydronium Ions

A

Hydrogen ion does not exist on its own but bonds to polar water molecules (H30+(aq)), is responsible for acidic properties

60
Q

Modified Arrhenius Theory

A

Acidic solutions form when substances react with water to form hydronium ions, bases are substances that react with water to produce hydroxide ions

61
Q

Acid base indicator

A

Substances that change color when the acidity of the solution changes, exist in 2 different forms (colours)

62
Q

pH paper

A

Universal Indicator that will give the exact pH of substance (indicates acidic/basic and pH), use info from various indicators to determine pH of unknown sample or use pH meter

63
Q

pH scale

A

Is a logarithmic scale that gives an indication of the nature of solution, measures relative acidity of a solution
Lower the pH the more acidic, decrease of 1 pH unit of a substance is 10x more acidic
**number of significant digits in the concentration equals the number of digits behind the decimal

64
Q

pH calculations

A

pH=-log(H30+)

(H30+)=10^-pH

65
Q

pOH calculations

A

pOH= -log(OH-)

(OH-)=10^-pOH

66
Q

Relationship between pH and hydronium ion concentrations

A

There is an inverse relationship, the higher the hydronium ion concentration the stronger the acid and the lower the pH.
The higher the hydroxide ion concentration the stronger the bases and the lower pOH (low pOH=high pH)
14-pH=pOH
14-pOH=pH

67
Q

Neutralization reaction

A

Double replacement reaction where an acid reacts with a base to produce water
MODIFIED: the reaction between hydronium and hydroxide ions to produce water

68
Q

Strong acids

A

High conductivity, high rate of reaction with active metals and carbonates, low pH, high concentration of hydronium ions, reacts completely with water to form hydronium ions

69
Q

Weak acids

A

Low conductivity, lower rate of reaction with active metals and carbonates, high pH, low concentration of hydronium ions, reacts incompletely with water to form few hydronium ions

70
Q

Strong bases

A

High conductivity, high pH, high concentration of hydroxide ions, dissociate completely in solution to produce hydroxide ions

71
Q

Weak Bases

A

Lower conductivity, pH closer to 7, lower concentration of hydroxide ions, react partially with water to produce few hydroxide ions, are either ionic or molecular compounds in their pure state

72
Q

Polyprotic acids

A

Contain more than one acidic hydrogen ion in their formula and react more than once with water to produce hydronium ions (weak acids)

73
Q

Monoprotic acids

A

Have only one acidic hydrogen ion in their formula and react only once with water to produce hydronium ions

74
Q

Monoprotic bases

A

React only once with water to produce hydroxide ions

75
Q

Polyprotic bases

A

React more than once with water, but all the reactions with water are much less then 50%, generally are weak bases decreases with each successive step

76
Q

Net ionic equations

A

Write the chemical equation to show that entities of the dissolved ionic compounds are actually present as separate dissociated aqueous ions, leave out spectator ions (ions that are present but do not take part in the reaction)

77
Q

Writing net ionic equations

A

Cancel identities that appear on both reactant and product sides, when cancelling out spectator ions they must be identical in every way- # of moles, form(atom, ion, molecule), and state of matter

78
Q

Stoichiometric calculations

A

Calculations that makes use of a balanced equation

79
Q

Gravimetric stoichiometry

A

The procedure for calculating the masses of reactants and products in a chemical reaction, chemical amount calculations are made from mass measurements

80
Q

Gas stoichiometry

A

Extended to gases using gas volume, pressure, temperature, molar volume and ideal gas Law
Use molar volumes to convert between moles/volume if at STP or SATP(STP=22.4L/mol, SATP=24.8L/mol)
If gas is not at SATP or STP conditions then you must use ideal gas law (PV=nRT) to convert between moles and volume

81
Q

Solution stoichiometry

A

The application of stoichiometric calculation principles to substances in solution, provided with the volume or amount concentration of one substance and asked to find the volume or amount concentration of another substance

82
Q

Steps to solve stoichiometric calculations

A
  1. Write balanced chemical equation
  2. Determine the 2 substances as the known and the unknown
  3. Calculate the # of moles of given substance
  4. Use mole ratio to determine the unknown # of moles
  5. Convert # of moles of unknown to a mass, volume, or concentration
83
Q

Percent yield

A

Evaluate the overall experiment by calculating percent yield, the ratio of actual or experimental quantity of product obtained to the theoretical yield obtained from stoichiometry calculation
Percent yield=actual yield/predicted yield x100

84
Q

Common sources of uncertainty

A

All measurements, purity of the grade of chemical used, washing a precipitate, qualitative measurements

85
Q

Qualitative analysis

A

Identification of a specific substance present

86
Q

Quantitative analysis

A

Determination of the quantity of a substance present

87
Q

Colorimetry

A

Analysis by color, light emitted, absorbed or transmitted by chemical used

88
Q

Gravimetric analysis

A

Uses stoichiometric calculations from the mass of a reagent

89
Q

Titration analysis

A

Uses stoichiometric calculations from the volume of a reagent

90
Q

Flame test

A

Some ions impart specific colours when subjected to a flame

91
Q

Solution color

A

Some jones create specific colours when in solution

92
Q

Precipitates

A

Using solubility table can determine what ions are in solution by whether or not a precipitate is formed

93
Q

Precipitation reaction

A

A reaction where a slightly soluble product (precipitate) is formed

94
Q

Limiting reagent

A

Reactant whose entities are completely consumed in a reaction, the reaction stops when and because all of this reactant is used up and none remains

95
Q

Excess reagent

A

Other reactant must be present in a greater quantity than is required or you would run out of it first, some of it will remain when limiting reagent has completely reacted (to ensure it goes to completion add 10% to amount of excess reagent)

96
Q

Determining limiting reagent

A
  1. Write balanced equation
  2. Calculate # of moles of each reactant
  3. Set up mole ratio of 2 reagents
  4. Find which of the 2 reagents will be used up in the reaction
  5. Use molar ratio of limiting reagent and the unknown to solve
  6. convert #of moles of unknown into mass/volume/concentration
97
Q

Titration

A

Used to determine amount concentration of substances in solution, process of carefully measuring and controlling the addition of a solution (titrant) from a burrette into a measured volume of another solution (sample) in an Erlenmeyer flask until the reaction is judged to be complete

98
Q

Burette

A

Precisely marked cylinder with a stopcock at one end, allows you to control and measure the volume of the reacting solution

99
Q

Equivalence point

A

Point at which the exact theoretical amount of titrant has been added to completely react with the sample, look for a sudden change in an observable property (colour,pH,conductivity)

100
Q

Endpoint

A

Point during a titration when this sudden change is observed, titration is stoped and the volume of titrant is determined

101
Q

Observable endpoint

A

Use indicators that will change colour if in an acid or a base, will show momentary colour change when titrant contacts sample, closer to endpoint color change lingers, to be accurate make sure observed endpoint occurs as closer as possible to reactions equivalence point

Chem 20 (8 decks)

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