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CHEM111 > CHEM111 CHAPTER 8-12 > Flashcards

CHEM111 CHAPTER 8-12 Flashcards

1
Q

Thermodynamics

A

Study of energy and temperature changes

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2
Q

Potential energy

A

Stored energy that matter possesses due to its position, condition and/or composition (cohesive force)

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3
Q

Kinetic energy

A

Energy of motion (disruptive force)

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4
Q

What happens when a substance absorbs heat energy

A

The particles within the substance vibrate and move faster

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5
Q

Equation that represents energy, heat and work

A

ΔE=q+w. Triangle E represents change in energy, q is heat, and w is work

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6
Q

Exothermic change

A

Releases heat energy (-q)

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7
Q

Endothermic change

A

Absorbs energy (+q)

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8
Q

Law of conservation of energy

A

Also called the first law of thermodynamics that states energy cannot be created or destroyed (ΔEsystem)=-(ΔEsurroundings)

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9
Q

Specific heat

A

Relationship between heat, mass, and temperature. It is the amount of heat required to raise the temperature of one gram of material by one degree Celsius. s=q/mΔT

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10
Q

Coffee cup calorimetry

A

When water is placed in an insulated cup covered with a cork lid and a thermometer. The system to be studied is placed inside the coffee cup and the reaction or change is allowed to take place

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11
Q

Bomb calorimetry

A

Measures the heats of reactions for fuels and other high energy chemical changes. T substance is placed in a heavy steal container and outside of the container is a second container filled with water and equipped with a thermometer

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12
Q

Extensive property

A

Energy absorbed or released in a reaction

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13
Q

Fuel value

A

The amount of heat energy that can be released by a combustion reaction of a certain substance (expressed in terms of energy per unit gram)

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14
Q

Reaction enthalpy

A

Amount of heat energy that is absorbed or released in a chemical reaction at constant pressure represented by triangle Hrxn

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15
Q

Physical state depends on (4):

A
  • Chemical identity of the matter (what it is)
  • Temperature of the matter
  • Pressure the matter is subjected to
  • The intermolecular forces that hold it together
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16
Q

Q

A

Representation of total energy released or absorbed by a system

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17
Q

Specific heat capacity

A

When heat capacity is given per gram of substance

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18
Q

Molar heat capacity

A

When heat capacity is given per mole of substance

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19
Q

Sublimination

A

When a material goes from solid to gas without going through the liquid state (endothermic - energy required to break intermolecular bonds)

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20
Q

Deposition

A

When a material goes from gas to solid without going through the liquid state (exothermic - energy released as intermolecular bonds form)

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21
Q

Ionic bonds

A

Attraction between oppositely charged ions

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22
Q

Covalent bonds

A

Attraction between two nonmetals and results when two atoms share electrons

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23
Q

Octet rule

A

States atoms are stabilized by 8 electrons in their valence level

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24
Q

Expanding octet

A

A bonding arrangement in which an atom has 10 or 12 valence electrons. This is ONLY possible with elements in rows 3-7 on periodic table

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25
Q

Single covalent bond

A

Two atoms share one pair of valence electrons

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26
Q

Double covalent bond

A

Two atoms share two pairs of valence electrons

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27
Q

Triple covalent bond

A

Two atoms share three pairs of electrons

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28
Q

Coordinate covalent bond

A

Both electrons of a shared pair originate from the same atom

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29
Q

Strength of bond

A

The more electron pairs being shared, the greater the strength of bond. increase as it goes single -> double -> triple

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30
Q

What element is never the central atom

A

Hydrogen because only two valence electrons. Or Fluorine

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31
Q

Localized electron model

A

Assumes all electrons are in pairs thus it does not handle odd-electron molecules very easily

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32
Q

Steps to drawing Lewis Structures (4)

A
  • Add up all the valence electrons
  • Frame the structure using single bonds (typically the atom that is nearer to the lower left of the periodic table is the central atom)
  • Fill octets of the outer atom first and add lone pairs
  • Fill the octet on the central atom
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33
Q

Formal charges

A

the charge that would reside on the atom if all of the bonding electrons were shared equally.

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34
Q

Resonance structures

A

Set of Lewis structures that show how electrons are distributed around a molecule or ion. Resonance structures are used when a single Lewis structure cannot adequately depict the structure (for 2nd or 3rd covalent bonds, and lone pairs)

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35
Q

Valence shell electron pair repulsion (VSEPR)

A

Model based on the idea that electrons repel each other, and therefore they occupy regions that are as far away from each other as possible

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36
Q

Electronic geometry

A

Arrangement of electrons around a central atom

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37
Q

Molecular geometry

A

Arrangement of atoms within a molecule

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38
Q

Two VSEPR electron pairs

A

linear

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39
Q

Three VSEPR electron pairs

A

Trigonal planar (120 deg) - Bent

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40
Q

Four VSEPR electron pairs

A

Tetrahedral (109.5 deg) - Trigonal pyramidal

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41
Q

Five VSEPR electron pairs

A

Trigonal Bipyramidal

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42
Q

Six VSEPR electron pairs

A

Octahedral

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43
Q

What shape is a water molecule

A

Tetrahedral

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44
Q

Electronegativity

A

Measure of the ability of that atom to attract shared electrons towards itself and is affected by atom size, nuclear charge, and the number of non valence electrons

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45
Q

Pauling scale

A

Numerical scale of electronegativities based on bond-energy calculations for different elements joined by covalent bonds

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46
Q

Pure covalent bond

A

Bond between two identical atoms or atoms of the same electronegativity (equal sharing of electron)

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47
Q

Polar covalent bond

A

Covalent bond between atoms of intermediate difference in electronegativity (unequal sharing of electrons)

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48
Q

Greater difference in electronegativity = ____

A

The greater the polarity of the bond

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49
Q

Bond polarity

A

Measure of the inequality in sharing of electrons in a bond

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50
Q

Dipole

A

An overall polarity in which different sides of a molecule have slight positive and negative charges

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51
Q

Which element is the most electronegative

A

Fluorine. It has the strongest pull for electrons

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52
Q

Which elements are the least electronegative

A

Cesium and francium. It has the weakest hold on electrons

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53
Q

Bond summary (3):

A
  • Covalent: Difference in electronegativity less than 0.5
  • Polar covalent: Difference between 0.5 and 2.0
  • Ionic: Difference greater than 2.0
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54
Q

Types of intermolecular forces (3):

A
  • Hydrogen bonding
  • Dipole-dipole forces (van der waals forces)
  • London dispersion forces (van der waals forces)
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55
Q

Dipole-dipole forces

A

Attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. Can interact with each other electrostatically (positive end of one molecule lines up with negative end)

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56
Q

Hydrogen bonding

A

Strong dipole-dipole forces when H is attached to a highly electronegative elements such as N, O, or F (very polar bond)

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57
Q

London dispersion forces

A

Results from the motion of electrons that creates temporary dipoles in molecules.

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58
Q

Strength of intermolecular forces least to greatest

A

London dispersion –> dipole-dipole –> hydrogen bonding

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59
Q

Key pointers of hydrogen bonding (2):

A
  • Requires donor and acceptors (donor is H bound to an N, O, F, and acceptor is a lone pair on an N, O, F)
  • Intermolecular hydrogen bonding also seen in alcohols (organic molecule containing O-H group) and amines (organic molecule containing N-H group)
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60
Q

Intermolecular forces

A

Usually [weak] forces of attraction or repulsion between individual molecules

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61
Q

Strength of London dispersion also depends on what factors

A

Depends on molecular shape/size. Electrons in elongated molecules are more easily displaced than those in small/symmetrical molecules.

62
Q

N-pentane

A

Large contact area, strong attraction

63
Q

Isopentane

A

Less surface area, less attraction

64
Q

Neopentane

A

Small contact area, weakest attraction

65
Q

Viscosity

A

Measure of a liquids resistance to flow. Liquids with strong intermolecular forces are highly viscous

66
Q

Types of crystalline solids (6):

A
  • Ionic solids
  • Metallic solids
  • Molecular solids
  • Network atomic solids
  • Polymer network
67
Q

Ionic solids

A

Extended network of ions held together by ion-ion interactions (NaCl, MgO)

68
Q

Metallic solids

A

Extended network of atoms held together by metallic bonding (Cu, Fe)

69
Q

Molecular solids

A

Discrete molecules held together by intermolecular forces (HBr, H2O)

70
Q

Network atomic solids

A

Extended network of atoms held together by covalent bonds (C, Si)

71
Q

Polymer network

A

Macromolecule where each constitutional unit is connected to each other’s constitutional unit

72
Q

Ideal gas law

A

Relationship between pressure, volume, and temperature, and the number of moles of an ideal gas expressed as PV = nRT where R is the gas constant

73
Q

Barometer

A

Device used to measure atmospheric pressure. A long tube is filled with a liquid (usually mercury) then turned upside down in a liquid. Gravity pulls the liquid down, creating a vacuum in the top of tube and the pressure from the outside air pushes the liquid up the tube. The higher the air pressure, the higher the liquid is pushed into the tube.

74
Q

Gauge pressure

A

Difference between a compressed gas pressure and atmospheric pressure

75
Q

Atmosphere (atm)

A

Unit of gas pressure (1atm=760mmHg)

76
Q

The gas laws

A

Generalizations that summarize in mathematical terms, experimental observations about the relationships among the amount (moles), pressure, temperature and the volume of gas

77
Q

Boyles law

A

As pressure increases, volume decreases. As pressure decreases, volume increases (temperature and number of moles remain unchanged). PV= k

78
Q

Charles law

A

As temperature increases, volume increases. As temperature decreases, volume decreases (pressure and number of moles remain unchanged) V/T = b

79
Q

Gay-Lussacs law

A

As temperature increases, pressure increases. As temperature decreases, pressure decreases (volume and number of moles remain unchanged) P/T =c

80
Q

Combined gas law equation

A

P1V1/T1 = P2V2/T2

81
Q

Avogadro’s law

A

States that, if pressure and temperature are constant, the volume of a gas is proportional to the number of moles of gas present. V/n =a

82
Q

Partial pressure

A

Pressure caused by one gas in a mixture. Adding up all the partial pressures gives the total pressure

83
Q

Universal gas constant (R)

A

R= 0.08206 L atm/K mol

84
Q

Effusion

A

Process of a gas escaping from a container; lighter gases effuse more quickly than heavier gases

85
Q

Diffusion

A

Term used to describe the mixing of gases; Particles randomly move, and as they do, they slowly spread from areas of higher concentration to lower concentration

86
Q

How does the pressure change if the temperature of the gas increases?

A

At higher temperatures, the particles move faster

87
Q

What happens to the pressure of a gas if the volume increases?

A

If the volume of a container increases, the particles have to travel farther before reaching the walls of the container. As a result, the particles collide with the container less frequently

88
Q

What happens to the pressure of a gas if the number of particles in a container is increased?:

A

More gas particles mean the particles will strike the sides of the container more often, creating more pressure (think of pumping a tire)

89
Q

Solution

A

A homogenous mixture of 2 or more substances (can be gases, liquids, solids)

90
Q

Solute

A

Substance that is dissolved in a solution; smaller component

91
Q

Solvent

A

A major component of a solution

92
Q

Solvent

A

The component of the mixture present in the greatest amount

93
Q

Aqueous solution

A

When water is the solvent

94
Q

Dilute

A

A solution that has very little solute

95
Q

Concentrated

A

A solution that has lots of solute

96
Q

Saturated

A

A solution containing the maximum amount of dissolved solute

97
Q

Percent by mass equation

A

mass of solute/mass of solution x 100%. To find measure of concentration

98
Q

Percent by volume equation

A

Volume of solute/volume of solution x 100%. To find measure of concentration

99
Q

Molarity (M)

A

Moles of solute/litre of solution; measure of concentration

100
Q

Molality (m)

A

Moles of solute/kg of solvent

101
Q

Percent composition

A

Percent of the solute present in the solution

102
Q

Parts per million (ppm)

A

A measure of concentration used for dilute solutions (1g of solute per 1,000,000 grams of solution); mass of solute/mass of solution x 10^6

103
Q

Parts per billion (ppb):

A

A measure of concentration used for dilute solutions (1g of solute per 1,000,000,000 grams of solution); mass of solute/mass of solution x 10^9

104
Q

What do colligative properties depend on

A

Depend only on the number of molecules or ions of the solute that are dissolved. Examples of colligative properties are boiling points, freezing points, etc

105
Q

Osmosis

A

Tendency of water to move towards regions of greater concentrations

106
Q

Freezing point depression

A

A colligative property of water; presence of solute in an aqueous solution lowers the freezing point below that of pure water

107
Q

Boiling point elevation

A

A colligative property of water; the presence of solute in an aqueous solution raises the boiling point above that of pure water

108
Q

Semipermeable membrane

A

Lipid bilayer where water can pass back and forth through membrane, but most molecules and ions cannot

109
Q

Precipitation reactions

A

When two solutions are combined and a solid product is formed

110
Q

Metal displacement reaction

A

A reaction between two metals in which one metal is oxidized to its ionic form while the other metal ion is reduced to its elemental form

111
Q

Gravimetric analysis:

A

A technique that uses the mass of a precipitate to determine the concentration of a reactant

112
Q

Acids (5):

A
  • ACIDS ARE PROTON (H+) DONORS AND PRODUCES H+ OR H3O+ IN WATER
  • Sour taste
  • Turns litmus red
    -Some metals dissolve in acid, giving off H2 gas
  • Examples: Vinegar, citric acid
113
Q

Base (6):

A
  • BASES ARE PROTON (H+) ACCEPTORS AND PRODUCES OH- IN WATER
  • Bitter taste
    -Turns litmus blue
  • Fatty substances dissolve in base
  • Slippery to touch
  • Examples: Soaps, cleaners containing ammonia
114
Q

What is the number one produced chemical worldwide?

A

Sulfuric Acid

115
Q

Arrhenius Acid

A

A hydrogen containing compound that dissolves in water releasing H+ ions and are polar covalent molecules when not in solutions but form ions in presence of water

116
Q

Arrhenius Base

A

A hydroxide containing compound that dissolves in water to give OH- ions and are ionic compounds that dissociate in water

117
Q

Arrhenius theory

A

Theory of ionization. States that acids are hydrogen containing compounds which give H+ ions or protons in dissociation in water and bases are the hydroxide compounds which give OH- ions in dissociation in water

118
Q

Bronsted-Lowry Theory (more common)

A

States that a substance can function as an acid only in the presence of a base; a substance can only function as a base in the presence of an acid

119
Q

Ionization

A

When molecular compounds dissolve in water to produce ions

120
Q

Dissociation

A

When ionic compounds dissolve to produce ions and occurs when water molecules “pull apart” the ionic crystal

121
Q

Amphoteric substance

A

Substances that can act as an acid or base such as water

122
Q

Hydronium ion

A

H3O+, the ion formed in acidic aqueous solutions

123
Q

Protonation

A

Gain of an H+

124
Q

Deprotonation

A

Loss of H+

125
Q

Monoprotic acids

A

Acids that can release only one proton into solution

126
Q

Polyprotic acids

A

Acids that donate two or more protons

127
Q

Diprotic acids

A

Acids that can release 2 protons (Sulfuric acid, carbonic acid)

128
Q

Tripotic acid

A

Acids that can release 3 protons (Phosphoric acid)

129
Q

Strong acids

A

Acids that completely ionize in water

130
Q

Weak acids

A

Acids that only partially ionize in water (eg. HF)

131
Q

Equilibrium

A

Two opposite processes taking place at equal rates

132
Q

Conjugate acid

A

Acid on the right hand side of chemical equation in an acid-base equilibrium; acid formed when a base reacts with H+

133
Q

Conjugate base

A

Base on the right hand side of a chemical equation in an acid-base equilibrium; base formed when an acid releases an H+

134
Q

Neutralization reaction

A

Reaction in which an acid and a base combine; acids and hydroxide bases continue to produce water and an ionic compound = salt

135
Q

Acid rain

A

The effect observed when nonmetal oxides combine with moisture in the atmosphere to produce acidic rainfall

136
Q

Common strong acids (7):

A
  • Nitric acid
  • Sulfuric acid
  • Perchloric acid
  • Chloric acid
  • Hydrochloric acid
  • Hydrobromic acid
  • Hydroiodic acid
137
Q

Strong electrolyte

A

Soluble salts, strong acids and bases

138
Q

Weak electrolyte

A

Weak acids and bases

139
Q

Non-electrolyte

A

Molecular compounds

140
Q

Acid-base titration

A

Determination of an unknown concentration of an acid or base by exactly neutralizing it with an acid or base of known concentration

141
Q

Litmus

A

Turns blue in the presence of a base and red in the presence of acid

142
Q

Phenolphthalein

A

Turns bright pink in a basic solution but colourless in an acidic solution

143
Q

pH Paper

A

Paper containing a blend of indicators that can be used to estimate pH based on color; produces colors ranging from deep red to deep blue

144
Q

pH scale

A

Scale that indicates the relative concentration of acid or base in an aqueous solution; pH is defined as the negative log of the hydronium concentration.

145
Q

pH scale indication

A

7 is neutral, lower pH is acidic, higher pH is basic

146
Q

Buffer

A

A solution that can resist pH change upon addition of an acidic or basic components and able to neutralize small amounts of acid or base thus maintaining the pH of the solution relatively stable

147
Q

How are buffers able to function

A

Buffers are able to resist pH change because the conjugate acid and conjugate base are both present in appreciable amounts of equilibrium

148
Q

Buffers contain either….

A
  • A mixture of a weak acid and its conjugate base
  • A mixture of a weak base and its conjugate acid
149
Q

pH of fluid inside cells

A

6.8

150
Q

pH of bloodstream

A

7.4

CHEM111 (6 decks)

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